Third World Congress on Oxidation Catalysis

Studies in Surface Science and Catalysis 110 3rd WORLD CONGRESS ON OXIDATION CATALYSIS

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S t u d i e s in S u r f a c e S c i e n c e a n d C a t a l y s i s Advisory Editors: B. Delmon and J.T. Yates

Vol. 110

3rd WORLD CONGRESS ON OXIDATION CATALYSIS Proceedings of the 3rd World Congress on Oxidation Catalysis, San Diego, CA, U.S.A., 21-26 September 1997

Editors R.K. Grasselli

University of Delaware, Newark, DE, U.S.A.

S.T. O y a m a

Virginia Polytechnic Institute, Blacksburg, VA, U.S.A.

A.M. Gaffney

ARCO Chemical Company, Newton Square, PA, U.S.A. J.E. Lyons

Sun Corporation, Marcus Hook, PA, U.S.A.

1997 ELSEVIER Amsterdam - - Lausanne m New York--- Oxford m Shannon m Singapore m Tokyo

ELSEVIER SCIENCE B.V. Sara Burgerhartstraat 25 P.O. Box 211, 1000 AE Amsterdam, The Netherlands

ISBN 0-444-82772-2 91997 Elsevier Science B.V. All rights reserved. No part of this publication may be reproduced, stored in a retrieval system or transmitted in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, without the prior written permission of the publisher, Elsevier Science B.V., Copyright & Permissions Department, P.O. Box 521, 1000 AM Amsterdam, The Netherlands. Special regulations for readers in the U.S.A.- This publication has been registered with the Copyright Clearance Center Inc. (CCC), 222 Rosewood Drive, Danvers, MA 01923. Information can be obtained from the CCC about conditions under which photocopies of parts of this publication may be made in the U.S.A. All other copyright questions, including photocopying outside of the U.S.A., should be referred to the copyright owner, Elsevier Science B.V., unless otherwise specified. No responsibility is assumed by the publisher for any injury and/or damage to persons or property as a matter of products liability, negligence or otherwise, or from any use or operation of any methods, products, instructions or ideas contained in the material herein. This book is printed on acid-free paper. Printed in The Netherlands

3rd World Congress on Oxidation Catalysis - Atom Efficient Catalytic Oxidations for Global Technologies

PREFACE

The 3rd World Congress on Oxidation Catalysis has its roots in the European Workshop on Selective Oxidation held in Louvain, Belgium in 1985. Out of this workshop grew the 1st World Congress held in Rimini, Italy in 1989, the 2nd in Benalmadena, Spain in 1992, and the 3rd being held now in 1997 in San Diego, California, USA. Out of the small core of dedicated and enthusiastic scientists assembled in Louvain in 1985, grew now a broad base of scientists and technologists from academia, industry and government laboratories who are fervently pursuing the subject of oxidation catalysis and are eager and willing to exchange their findings at the current meeting. The overall theme of the 3rd World Congress is "Atom Efficient Catalytic Oxidations for Global Technologies". We chose this theme to stimulate the participants to report their findings with an emphasis on conserving valuable material in their catalytic transformations, as well as conserving energy, and that in an environmentally responsible manner. Progress towards this stated goal is substantial as evidenced by the tremendous response of our community in their participation of quality publications compiled in these Proceedings of the Congress. The subjects presented span a wide range of oxidation reactions and catalysts. These include the currently important area of lower alkane oxidation to the corresponding olefins, unsaturated aldehydes, acids and nitriles. In this manner, the abundant and less expensive alkanes replace the less abundant and more expensive olefins as starting materials for industrially important intermediates and chemicals. In the oxidative activation of methane the emphasis is shifting towards the use of extremely short contact times and newer more rugged catalysts. In the area of olefin oxidations, of particular note are the high efficiency epoxidation of propylene, and new detailed mechanistic insights into the oxidation of ct,~-unsaturated aldehydes to the corresponding unsaturated acids. Substantial progress is reported in the area of the selective oxidation and ammoxidation of substituted aromatics and heteroaromatics. These include higher yields of desired products, new and more durable catalysts, as well as a reduction of undesirable byproducts. The use of oxidation catalysis to produce fine chemicals is experiencing an explosive growth. A plethora of novel approaches are presented which include shape selective epoxidations. Oxidation in confined structures is coming out of its infancy and the use of TS-'I as catalyst is becoming a standard. New approaches are being presented invoking the shape selective character of the nano-environment of peroxytungstates anchored on selected supports. The important areas of combustion, engineering and environmental aspects of oxidation catalysis, as well as their theoretical, computational and modeling approaches round off the program. Not to be overlooked is perhaps the most ambitious, the subject of structure selectivity/activity correlations, an area always worthy of further attention and in depth study. A noble effort thereof has been put forward and is reported here. We are coming ever closer to the ultimate goal of a rational approach to catalyst design and synthesis. There is still ample room for further efforts towards this goal, but the foundations are being formed for a bright and rewarding future of rationally predictive oxidation catalysis. The five featured lectures and seven plenary lectures constitute the general background and overview of the subject matter at hand. The 104 contributed papers and 13 poster manuscripts, summarized in this compendium, probe new avenues to achieve catalytically efficient oxidation reactions for the future needs of mankind in a global

environment. A large number of countries responded to this challenge by their representatives giving oral presentations or posters, and in particular by supplying the written scientific documents contained in this volume. Our sincere thanks go out to all of the contributors. The countries participating in the Congress and contributing to the Proceedings reported here made the 3rd World Congress on Oxidation Catalysis a truly international event, they are: Argentina, Azerbaijan, Belgium, Brazil, Bulgaria, Canada, China, Czech Republic, Finland, France, Germany, Greece, India, Iran, Ireland, Israel, Italy, Japan, Korea, Latvia, Netherlands, New Zealand, Poland, Portugal, Russia, Saudi Arabia, Singapore, Slovenia, South Africa, Spain, Sweden, Taiwan, Thailand, Ukraine, United Kingdom, and United States. We conclude on the basis of the foregoing, that the future of oxidation catalysis is secure and has never been brighter than at this juncture. We are confident, and the Proceedings support this view, that many new and improved selective oxidation processes and catalysts will be discovered and commercialized over the next decade, and that our knowledge towards a rational design of selective oxidation catalysts is within our grasp. With this optimism we look confidently towards the future and to a successful 4th World Congress on Oxidation Catalysis in the next millennium, in the year 2001. Thank you all for partaking in our Congress and for working in an exciting and promising area of catalysis. May success come your way in abundance in the coming years.

Robert K. Grasselli S. Ted Oyama Anne M. Gaffney James E. Lyons

vii

TABLE OF CONTENTS

Preface R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons Featured Lectures F-1

Molecular Mechanism of Heterogeneous Oxidation-Organic and Sofid State Chemists' Views J. Haber

F-2

The Multifunctional Properties of Heterogeneous Catalysts, Active and Selective in the Oxidation of Light Paraffins F. Cavani and F. Trifirb

19

Selective Oxidation of Hydrocarbons Catalyzed by Heteropoly Compounds M. Misono, N. Mizuno, K. Inumaru, G. Koyano and X. H. Lu

35

The Future of Industrial Oxidation Catalysis Spurred by Fundamental Advances B. Delmon

43

F-3

F-4

Plenary Lectures P-1

P-2

P-3

P-4

Molecular Approach to Active Sites on Metallic Oxides for Partial Oxidation Reactions J.C. V6drine

61

In Situ Electrochemically Controlled Promotion of Complete and Partial Oxidation Catalysts C.G. Vayenas and S.I. Bebelis

77

Reductive and Oxidative Activation of Oxygen for Selective Oxygenation of Hydrocarbons K. Otsuka

93

The Selective Oxidation of Methanol: A Comparison of the Mode of Action of Metal and Oxide Catalysts D. Herein, H. Werner, Th. SchedeI-Niedrig, Th. Neisius, A. Nagy, S. Bernd and R. Schl5gl 103

viii

P-5

P-6

P-7

Gold as a Low-Temperature Oxidation Catalyst: Factors Controlling Activity and Selectivity M. Haruta

123

The Selective Epoxidation of Non-Allylic Olefins over Supported Silver Catalysts J.R. Monnier

135

Redox Molecular Sieves as Heterogeneous Catalysts for Liquid Phase Oxidations R.A. Sheldon

151

PartA Structure Selectivity/Activity Correlation A-1

Synergistic Effects in Mulficomponent Catalysts for Selective Oxidation P. Courtine and E. Bordes

177

A-2

Synergetic Effects Promoted by in Operandi Surface Reconstructions of Oxides E.M. Gaigneaux, J. Naud, P. Ruiz and B. Delmon

185

Further Study on the Synergetic Effects between MoO3 And Sn02 E.M. Gaigneaux, S.R.G. Carrazan, L. Ghenne, A. Moulard, U. Roland, P. Ruiz and B. Delmon

197

The Nature of the Active/Selective Phase in VPO Catalysts and the Kinetics of n-Butane Oxidation D. Dowell and J.T. Gleaves

199

Understanding the Microstructural Transformation Mechanism which Takes Place During the Activation of Vanadium Phosphorus Oxide Catalysts G.J. Hutchings, A. Burrows, S. Sajip, C.J. Kiely, K.E. Bere, J.C. Volta, A. Tuel and M. Abon

209

Structural and Catalytic Aspects of Some Nasicon - Based Mixed Metal Phosphates P.A. Agaskar, R.K. Grasselli, D.J. Buttrey and B. White

219

Selective Reactivity of Oxygen Adatoms on Mo(112) for Methanol Oxidation K. Fukui, K. Motoda and Y. Iwasawa

227

A-3

A-4

A-5

A-6

A-7

A-8

A-9

A-10

Mechanistic Studies of Alkane Partial Oxidation Reactions on Nickel Oxide by Modern Surface Science Techniques N.R. Gleason and F. Zaera

235

Structure and Catalysis of LixNi2_x02 Oxide Systems for Oxidative Coupling of Methane T. Miyazaki, T. Doi, T. Miyamae and I. Matsuura

245

Reaction Induced Spreading of Metal Oxides: in situ Raman Spectroscopic Studies During Oxidation Reactions Y. Cai, C-B. Wang and I.E. Wachs

255

A-11

Temperature Programmed Desorpfion of Ethylene and Acetaldehyde on Uranium Oxides. Evidence of Furan Formation from Ethylene 265 H. Madhavaram and H. Idriss

A-12

Active Sites of Vanadium-Molybdenum-Containing Catalyst for Allyl Alcohol Oxidation: ESR Study in situ O.V. Krylov, N.T. Tai and B.V. Rozentuller

275

Lower Alkane Oxidation

A-13

Oxidative Dehydrogenation of Ethane over Vanadium and Niobium Oxides Supported Catalysts P. Ciambelli, L. Lisi, G. Ruoppolo, G. Russo and J.C. Volta

285

A-14

Partial Oxidation of Ethane over Monolayers of Vanadium Oxide. Effect of the Support and Surface Coverage. M.A. BaSares, X. Gao, J.L.G. Fierro and I.E. Wachs 295

A-15

The Ethane Oxidative Chlorination Process and Efficient Catalyst for It M.R. Flid, 1.1. Kurlyandskaya, Yu.A. Treger and T.D. Guzhnovskaya

305

A-16

Oxidative Conversion of LPG to Olefins with Mixed Oxide Catalysts: Surface Chemistry and Reactions Network M.V. Landau, M.L. Kaliya, A. Gutman, L.O. Kogan, M. Herskowitz and P.F. van den Oosterkamp

315

Free Radicals as Intermediates in Oxidative Transformations of Lower Alkanes M.Yu. Sinev, L.Ya. Margolis, V.Yu. Bychkov and V.N. Korchak

327

A-17

A-18

Alternative Methods to Prepare and Modify Vanadium-Phosphorus Catalysts for Selective Oxidation of Hydrocarbons V.A. Zazhigalov, J. Haber, J. Stoch, A.I. Kharlamov, I.V. Bacherikova and L.V. Bogutskaya

337

Active Species and Working Mechanism of Silica Supported MoO3 and V205 Catalysts in the Selective Oxidation of Light Alkanes A. Parmaliana, F. Arena, F. Frusteri, G. Martra, S. Coluccia and V. Sokolovskii

347

Mechanistic Aspects of Propane Oxidation over Ni-Co-Molybdate Catalysts D.L. Stern and R.K. Grasselli

357

Oxidative Dehydrogenation of Propane by Non-Stoichiometric Nickel Molybdates D. Levin and J.Y. Ying

367

Selective Oxidation of Propane into Oxygenated Compounds over Promoted Nickel-Molybdenum Catalysts J. Barrault, C. Batiot, L. Magaud and M. Ganne

375

Oxidative Dehydrogenation of Propane on CeNixOy (0 ~ x <_ 1) Mixed Oxides Hydrogen Acceptors L. Jalowiecki-Duhamel, A. Ponchel and Y. Barbaux

383

The Role of Adsorption in the Oxidation of a,[J-unsaturated Aldehydes on Mo-V-Oxide Based Catalysts B. Stein, C. Weimer and J. Gaube

393

A-25

A New Catalyst for Propane Ammoxidation: The Sn/V/Sb Mixed Oxide S. Albonetti, G. Blanchard, P.Burattin, S. Masetti and F. Trifir6

403

A-26

Formation of Active Phases in the Sb-V-, AI-Sb-V-, and AI-Sb-V-W-Oxide Systems for Propane Ammoxidation J. Nilsson, A.R. Landa-C&novas, S. Hansen and A. Andersson 413

A-27

Influence of Antimony Content in the Iron Antimony Oxide Catalyst and Reaction Conditions on the (Amm)Oxidation of Propene and Propane E. van Steen, G. Kuwert, A. Naidoo and M. Williams

423

Catalytic Selective Oxidation of C2-C4Alkanes over Reduced Heteropolymolybdates W. Li and W. Ueda

433

A-19

A-20

A-21

A-22

A-23

A-24

A-28

A-29

A-30

The Role of Metal Oxides as Promoters of V205A/-AI20~ Catalysts in the Oxidative Dehydrogenation of Propane J.M. L6pez Nieto, R. Coenraads, A. Dejoz and M.I. Vazquez

443

Alkane Oxidation over Bulk and Silica-Supported VO(H2PO4)2-Derived Catalysts G.K. Bethke, D. Wang, J.M.C. Bueno, M.C. Kung and H.H. Kung

453

A-31

The Nature of the Active Site of the (VO)2P207 Catalyst: An Investigation of the Chemical Composition and Dynamics of the Catalyst Surface B. Kubias, F. Richter, H. Papp, A. Krepel and A. Kretschmer 461

A-32

Partial Oxidation of C5 Hydrocarbons to Phthalic and Maleic Anhydrides over Suboxides of Vanadia: Use of Dicyclopentadiene as a Probe Molecule U.S. Ozkan, G. Karakas, B.T. Schilf and S.S. Ang 471

A-33

Role of Homogeneous Reactions in the Control of the Selectivity to Maleic and Phthalic Anhydrides in the Oxidation of n-Pentane Z. Sobalik, P. Ruiz and B. Delmon

481

A-34

Catalytic Oxidation of Alkanes at Millisecond Contact Times L.D. Schmidt and C.T. Goralski, Jr.

491

A-35

Catalytic Oxidative Dehydrogenation of Isobutane in a Pd Membrane Reactor T.M. Raybold and M.C. Huff

501

Fine Chemicals and Pharmaceuticals

A-36

A-37

A-38

Chemoselective Catalytic Oxidation of Polyols with Dioxygen on Gold Supported Catalysts L. Prati and M. Rossi

509

Promoting Effects of Bismuth in Carbon-Supported Bimetallic Pd-Bi Catalysts for the Selective Oxidation of Glucose to Gluconic Acid M. Wenkin, C. Renard, P. Ruiz, B. Delmon and M. Devillers

517

Oxidative Dehydrogenation of Glycofic Acid to Glyoxylic Acid over Fe-P-O Catalyst M. Ai and K. Ohdan

527

xii

A-39

Shape Selective Epoxidation of Crotyl Alcohol with H202 in the Presence of TS-1 G.J. Hutchings, P.G. Firth, D.F. Lee, P. McMorn, D. Bethell, P.C. Bulman Page, F. King and F. Hancock 535

A-40

Epoxidation of Tertiary Allylic Alcohols and Subsequent Isomerization of Tertiary Epoxy-Alcohols: A Comparison of some Catalytic Systems for Demanding Ketonization Processes J.-M. Br6geault, C. Lepetit, F. Ziani-Derdar, O. Mohammedi, L. Salles and A. Deloffre 545

A-41

Metal-Catalyzed Oxidations with Alkyl Hydroperoxides: A Comparison between tert-Butyl Hydroperoxide and Pinane Hydroperoxide H.E.B. Lempers and R.A. Sheldon

557

On the Way to Redox-Molecular Sieves as Multifuncfional Solid Catalysts for the One-Step Conversion of Olefins to Aldehydes or Ketones M. van Klaveren and R.A. Sheldon

567

Liquid-Phase Oxidation of Cyclohexane to Adipic Acid Catalysed by Cobalt Containing ~-zeofites I. Belkhir, A. Germain, F. Fajula and E. Fache

577

A-42

A-43

A-44

A-45 A-46

Nitrogen Oxides Catalyzed Selective Oxidation by Oxygen in the Liquid Phase A.B. Levina, S.S. Chornaja, I.A. Grigorjeva, O.N. Sergejeva and S.R. Trusov 585 Oxidative Coupling of Isobutene in a Two Step Process

H. Hiltner and G. Emig

593

Sofid Solutions for Cleaning up Chemical Processes using Hydrogen Peroxide S.L. Wilson and C.W. Jones

603

Engineering and Environmental Applications A-47 A-48

Catalytic Wet Air Oxidation of Wastewaters

J.C. B6ziat, M. Besson, P. Gallezot, S. Juif and S. Dur~cu

Catalytic Partial Oxidation of Methanol: H2 Production for Fuel Cells

L. Alejo, R. Lago, M.A. Pefia and J.L.G. Fierro

615 623

xiii

A-49

Catalytic Liquid-Phase Phenol Oxidation over Metal Oxides and Molecular Sieves. Reaction Kinetics and Mechanism A. Pintar, G. Ber~i(~, J. Batista and J. Levec 633

A-50

Ammonia Oxidation over CuO/Ti02 Catalyst: Selectivity and Mechanistic Study G. Bagnasco, G. Peluso, G. Russo, M. Turco, G. Busca and G. Ramis 643

A-51

Metalloporphyrin-Catalysed Oxidation of Azonaphthol Dyes: The Mechanism of Oxidative Bleaching by Oxoiron (IV) Porphyrins in Aqueous Solution G. Hodges, J.R. Lindsay Smith and J. Oakes

A-52

653

VOC's Abatement: Photocatalytic Oxidation of Toluene in Vapour Phase on Anatase Ti02 Catalyst V. Augugliaro, S. Coluccia, V. Loddo, L. Marchese, G. Martra, L. Palmisano, M. Pantaleone and M. Schiavello 663 PartB Methane Activation

B-1

Oxidation Processes on Stoichiometric and Nonstoichiometric Hydroxyapatites H. Hayashi, H. Kanai, Y. Matsumura, S.Sugiyama and J.B. Moffat

673

B-2

Oxidative Coupling of Methane in Sofid Oxide Fuels Cells Guo Xiu-Mei, Kus Hidajat and Chi-Bun Ching

683

B-3

Partial Oxidation of Methane to Synthesis Gas in a Fast Flow Membrane Reactor M. Alibrando and E.E. Wolf 693

B-4

Sustainable Ni/Ba TiO3 Catalysts for Partial Oxidation of Methane to Synthesis Gas R. Shiozaki, A.G. Andersen, T. Hayakawa, S. Hamakawa, K. Suzuki, M. Shimizu and K. Takehira

B-5

701

Synthesis of Early Transition Metal Carbides and their Application for the Reforming of Methane to Synthesis Gas A.P.E. York, J.B. Claridge, C. Marquez-Alvarez, A.J. Brungs, S.C. Tsang and M.L.H. Green 711

xiv

B-6

Partial Oxidation of Methane to Synthesis Gas using LnCoO3 Perovskites as Catalyst Precursors 721 R. Lago, G. Bini, M.A. PeSa and J.L.G. Fierro

B-7

Performance of Catalytic Properties of Reagent Catalyst in the Processes such as Methane Oxidative Coupling and Hydrogen Production by Methane Conversion M.I. Levinbuk, N.Y. Usachev, M.L. Pavlov, A.U. Loginov, L.V. Surkova, E.M. Savin, V.K. Smirnov and I.V. Ivkova 731

Combustion B-8

The Effect of the PbO Loading in the Oxidative Coupling of Methane over PbO/Si02 Catalysts H.J. Lugo, N. Teran, L. Villasmil, G. Castillo and D.M. Finol 737

B-9

Catalytic Combustion of Ethane over High Surface Area Lnl_xKxMnOa (Ln=La, Nd) Perovskites: The Effect of Potassium Substitution Y.Ng Lee, F. Sapi~a, E. Martinez, J.V. Folgado and V. Cortes Corber&n 747

B-10

Effect of Redox Treatment on Methane Oxidation over Binary Catalyst Yu.P. Tulenin, M.Yu. Sinev, V.V. Savkin and V.N. Korchak

757

B-11

Catalytic Combustion of Methane: Activation and Characterization of Pd/AI203 M.G. Carneiro da Rocha and Roger Frety

767

B-12

Activity of Manganese Dioxides towards VOC Total Oxidation in Relation with their Crystallographic Characteristics C. Lahousse, A. Bernier, E. Gaigneaux, P. Ruiz, P. Grange, B. Delmon 777

Catalyst Preparation B-13

Understanding the Surface Chemistry for Supported Vanadium Oxide Systems Modified with Phosphorus Oxide at Hydrocarbons Oxidation V.A. Zazhigalov, L.V. Bogutskaya, L.V. Lyashenko and I.V. Bacherikova 787

B-14

Effects of Cesium Doping on the Kinetics and Mechanism of the n-Butane Oxidative Dehydrogenation over Nickel Molybate Catalysts L.M. Madeira and M.F. Portela 797

B-15

A Comparison of Iron Molybdate Catalysts for Methanol Oxidation Prepared by Coprecipitation and New Sol-Gel Method A.P. Vieira Soares, M. Farinha Portela and A. Kiennemann

807

XV

B-16

B-17

B-18

Oxidation Catalysts Prepared by Mechanically and Thermally Induced Spreading of Sb203 and V205 on Ti02 U.A. Schubert, J. Spengler, R.K. Grasselli, B. Pillep, P. Behrens and H. KnOzinger

817

The Effect of Preparation Parameters on the BET Surface Area of Zr02 Powder YuanYang Wang, YanZhen Fan, YuHan Sun and SongYing Chen

829

Preparation of VOHPO4-O.5H20 and (VO)2P207 and their Catalytic Performance for Maleic Anhydride Synthesis T. Miyake and T. Doi

835

Alternate Oxidants

B-19 B-20

B-21

B-22

B-23

Hydroxylation of Benzene on ZSM5 Type Catalysts M. H~fele, A.Reitzmann, E. Klemm and G. Emig

847

Direct Hydroxylation of Benzene to Phenol by Nitrous Oxide

A.K. Uriarte, M.A. Rodkin, M.J. Gross, A.S. Kharitonov and G.I. Panov

857

Rapid Catalytic Oxygenation of Hydrocarbons with Perhalogenated Ruthenium Porphyrin Complexes J.T. Groves, K.V. Shalyaev, M. Bonchio and T. Carofiglio

865

Ethanol Oxidation Using Ozone over Supported Manganese Oxide Catalysts: An in situ Laser Raman Study Wei Li and S.T. Oyama

873

Generation of Singlet Oxygen from the Catalytic System H20/Ca(OH)2 and Appfications to the Selective Oxidation of Unsaturated Compounds J.M. Aubry and V. Nardello

883

Oxidation of Olefins and Aromatics

B-25

B-26

Toluene Gas Phase Oxidation to Benzaldehyde and Phenol over V-containing Micro- and Mesoporous Materials G. Centi, F. Fazzini and L. Canesson and A. Tuel

893

A Novel Selective Oxidation Catalyst: Ultrafine Complex Molybdenum Based Oxide Particles Y. Fan, W. Kuang, W. Zhang and Yi Chen

903

xvi

Liquid Phase Oxidation of Alkylaromatic Hydrocarbons over Titanium Silicalites G.N. Vayssilov, Z. Popova, S. Bratinova and A. Tuel

909

Coupled Vanadyl Centres in Vanadium Phosphorus Oxide Catalysts: Essential Structural Units for Effective Catalytic Performance in the Ammoxidation of Methylaromatics A. Br0ckner, A. Martin, B. LOcke and F.K. Hannour

919

B-29

Ammoxidation of Xylenes- Kinetics and Selectivity K. Beschmann, S. Fuchs and T. Hahn

929

B-30

Vanadium-Titanium Oxide System in B-Picoline Oxidation E.M. Al'kaeva, T.V. Andrushkevich, G.A. Zenkovets, G.N. Kryukova, S.V. Tsybulya and E.B. Burgina

939

Selective Alkene Epoxidation by Molecular Oxygen in the Presence of Aldehyde and Different Type Catalysts Containing Cobalt O.A. Kholdeeva, I.V. Khavrutskii, V.N. Romannikov, A.V. Tkachev and K.I. Zamaraev

947

Epoxidation of Olefins over Thermally Stable Polyimide-Supported Mo(VI) Complexes J.H. Ahn, J.C. Kim, S.K. Ihm and D.C. Sherrington

957

Selective Partial Oxidation of Propylene to Propylene Oxide on Au/Ti-MCM Catalysts in the Presence of Hydrogen and Oxygen Y.A. Kalvachev, T. Hayashi, S. Tsubota and M. Haruta

965

B-27

B-28

B-31

B-32

B-33

B-34

Immobilization of Triazacyclononane-type Metal Complexes on Inorganic Supports via Covalent Linking: Spectroscopy and Catalytic Activity in Olefin Oxidation Y.V. Subba Rao, D.E. DeVos, B. Wouters, P.J. Grobet and P.A. Jacobs 973

B-35

Simultaneous Determination of Reaction Kinetics and Oxygen Activity during Selective Oxidation of an Aldehyde over an Oxidic Multicomponent Catalyst M. Estenfelder and H.-G. Lintz

981

On the Mechanism of the Selective Oxy-Dehydrogenation of n-Butenes to 1,3-Butadiene on Magnesium Ferrite: an FT-IR Study E. Finocchio, G. Busca, G. Ramis and V. Lorenzelli

989

B-36

xvii

Oxidation in Confined Structures B-37

Cyclohexene Oxidation Catalyzed by Titanium Modified Hexagonal Y Type Zeofites

K.J. Balkus, Jr., A.K. Khanmamedova and J. Shi B-38 B-39

Oxidations Catalyzed by Zeofite Ti-UTD-1

K.J. Balkus, Jr. and A.K. Khanmamedova

Zeofite Titanium Beta: A Selective Catalyst in the Meerwein-PonndorfVerley-Oppenauer Reactions J.C. van der Waal, P.J. Kunkeler, K. Tan and H. van Bekkum

B-40

Selective Oxidation of Cyclohexane over Rare Earth Exchanged Zeofite Y E.L. Pires, M. Wallau and U. Schuchardt

B-41

Rationally Designed Oxidation Catalysts: Functionalized Metalloporphyrins Encapsulated in Transition Metal-Doped Mesoporous Silica Lei Zhang, Tao Sun and J.Y. Ying

B-42 B-43

Catalytic Oxidations with Biomimetic Vanadium Systems I.W.C.E. Arends, M. Pellizon Birelli and R.A. Sheldon

1015

1025

1029 1031

1043

Catalytic Oxidations by in situ Generated Peroxotungsten Complexes Immobilized on Layered Double Hydroxides (LDH): Relation Between Catalytic Properties and Peroxotungstate Micro-Environment B.F. Sels, D.E. DeVos and P.A. Jacobs

B-45

1007

Highly Selective Photochemical and Dark Oxidation of Hydrocarbons by 02 in Zeofites H. Frei

B-44

999

Homobimetallic Heptadentate Coordinated Iron Complexes in Montmorillonites as Methane MonoOxygenase Mimics P.P. Knops-Gerrits, S. Dick, A. Weiss, M. Genet, P.Rouxhet, X.Y. Li

and P.A. Jacobs

1051

1061

Theoretical, Computational and Modeling Studies B-46

Modeling the Transient CO Oxidation over Platinum

T.A. Nijhuis, M. Makkee, A.D. van Langeveld and J.A. Moulijn

1071

xviii

B-47

Dioxygen Activation with Sterically Hindered Tris(pyrazolyl)borate Cobalt Complexes K.H. Theopold, O.M. Reinaud, D. Doren and R. Konecny

1081

Designing Industrial Redox Catalysts for Selective Autoxidations of Hydrocarbons - A New Paradigm D. Masilamani

1089

B-49

Selectivity of Active Sites on Oxide Catalysts C. Batiot, F.E. Cassidy, A.M. Doyle and B.K. Hodnett

1097

B-50

A Novel Computer-Aided Technique for the Development of Catalysts for Propane Ammoxidation to Acrylonitrile X.-Q. Wu, Q.-X. Zhang, Q.-L. Dai, Z.-Y. Hou and D.-W. Lu

1107

Catalysts by Rational Design: Prediction and Confirmation of the Properties of the Co/Ce/Br Liquid-Phase Autoxidation Catalyst Based on the Kinetic Similarity to the Co/Mn/Br Catalyst R.K. Gipe and W. Partenheimer

1117

B-48

B-51

PartC Additional Oxidation Studies

C-1

The Kinetics of the Partial Oxidation of Methane to Formaldehyde over a Silica-Supported Vanadia Catalyst A.W. Sexton and B.K. Hodnett 1129

C-2

Catalytic Destruction of Volatile Organic Compounds on Platinum/Zeolite A. O'Malley and B.K. Hodnett 1137

C-3

High Temperature Propane Oxidation to Reducing Gas over Promoted Ni/MgO Catalysts. Role of Impregnation Condition and Promoter on Properties of Catalysts M.V. Stankovi6 and N.N. Jovanovi6 1145

C-4

Structural Sensitivity of the Oxidation Reactions Catalyzed by Dispersed Transition Metal Oxides: Role of Defect Structure V.A. Sadykov, S.F. Tikhov, S.V. Tsybulya, G.N. Kryukova, S.A. Veniaminov,V.N. Kolmiichuk, N.N. Bulgakov, L.A. Isupova, E.A. Paukshtis, V.I. Zaikovskii, G.N. Kustova and L.B.Burgina 1155

C-5

Oxidation of Cyclohexane using Polymer Bound Ru(lll) Complexes as Catalysts J. John, M.K. Dalai and R.N. Ram

1165

xix

C-6

Photoimmobilized Catalysts for Low-Temperature Oxidation of Olefins L.V. Lyashenko, V.M. Belousov, E.V. Kashuba

1175

C-7

Selective Oxidation Catalysis over Heteropoly Acid Supported on Polymer In Kyu Song, Jong Koog Lee, Gyo Ik Park and Wha Young Lee

1183

A Study of V205-K2S04-Si02 Catalysts for Catalytic Vapor-Phase Oxidation of Toluene to Benzaldehyde A.O. Rocha Jr., A.L. Chagas, L.S.V.S. Sufi6, M.F.S. Lopes and J.A.F.R. Pereira

1193

C-8

C-9

OyH2 Oxidation of Hydrocarbons on the Catalysts Prepared from Pd(ll) Complexes with Heteropolytungstates N.I. Kuznetsova, L.I. Kuznetsova, L.G. Detusheva, V.A. Likholobov, 1203 M.A. Fedotov, S.V. Koscheev and E.B. Burgina

C-10

Oxidation of n-Pentane to Phthalic or Maleic Anhydride: The Role of the VPO Catalyst Structural Disorder Z. Sobalik, S.Gonzalez, P. Ruiz and B. Delmon

1213

C-11

Electrochemical Oxidation of Propene using a Membrane Reactor with Sofid Electrolyte S. Hamakawa, T. Hayakawa, K. Suzuki, M. Shimizu and K. Takehira 1223

C-12

Vanadium Pentoxide Catalytic Membrane Reactor for Partial Oxidation of 1-Butene Sangjin Moon, Tayoon Kim, Seungdoo Park, Jihoon Jung and Sukin Hong 1231

C-13

Peroxidase Oxidation of Phenol by Catalase Immobilized on Carbon Materials E. Horozova, N. Dimcheva and Z. Jordanova

1239

Author Index

1245

This Page Intentionally Left Blank

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

Molecular mechanism of heterogeneous oxidation - organic and solid state chemists' views Jerzy Haber Institute of Catalysis and Surface Chemistry, Polish Academy of Sciences, Krakow, Poland Based on the discussion of the state-of-the-art in heterogeneous oxidation of hydrocarbons this paper reviews the most important but yet unanswered questions. It is emphasized that the oxide surface is in dynamic interaction with the gas phase. The role of electrophilic oxygen is illustrated and the scarcity of information on its formation at the oxide surface with point and extended defects is stressed. The molecular mechanisms of C-H bond activation are discussed and a conclusion is substantiated that both parts of the cleaved C-H bond bind to surface oxide ions with injection of electrons into the conductivity band of the oxide. The importance of the transfer of electrons between the oxide and the hydrocarbon+oxygen gas mixture as a redox system is illustrated. The lack of data on the role of coordination of surface cations, of the presence of defects and of the degree of surface hydroxylation in catalytic oxidation is underlined. The transition metal oxide monolayer catalysts are discussed and the question is raised concerning the mechanism, by which the support influences the behaviour of the monolayer. Also the mechanism of the addition of nucleophilic oxygen to the hydrocarbon molecule remains unanswered, although a hypothesis has been advanced that is favoured when crystallographic shear mechanism operates. A number of questions may be raised as to the mechanism of the operation of an oxide monolayer catalyst are isolated metal-oxygen polyhedra playing the role of active sites, or is the presence of oligomeric species required? Where are these sites located and how are their properties influenced by the type of crystal plane and the presence of defects? General conclusions concerning the mechanism of heterogeneous catalytic oxidation are discussed. 1.

INTRODUCTION

Oxygen is one of the most interesting elements playing a fundamental role in catalysis, because on the one hand it is a component of the most widely used type of catalysts - oxides, and on the other hand it is the reactant in one of the most important types of catalytic reactions - oxidation. The attack of oxygen on the hydrocarbon molecular is the easiest route to functionalize this molecule, and selective oxidation processes, in which hydrocarbon molecules are oxygenated to form alcohols, aldehydes or acids are the basis of the modern petrochemical

industry. They may be divided into vapour or liquid phase reactions, which are catalyzed by solid oxide catalysts and are carried out as heterogeneous catalytic processes and reactions in the liquid phase, catalyzed by transition metal complexes or by enzymes, which are commonly realized as homogeneous catalytic processes. In all these processes the essence of the catalytic act is the same and consists in the interaction of reacting hydrocarbon molecules with a group of atoms playing the role of an active site of the catalyst. In the case of heterogeneous catalytic processes carried out in the presence of solid oxide catalysts the atoms of active sites are either constituents of the oxide lattice or are supported on its surface forming part of the solid. When discussing the mechanism of heterogeneous catalytic oxidation of hydrocarbons, the organic chemists (1) usually consider only the fate of the reacting molecule, trying to unravel the reaction network in terms of consecutive oxidative additions and reductive eliminations and molecular rearrangements. The analysis is based on the consideration of possible transition states formed as a result of rearrangement of electron pairs due to exchange of electrons with an active site of the catalyst, which serves as an acceptor or a donor of electrons. The observed influence of different substituents in the organic molecule on the reaction rate and selectivity helps in elucidation of the transition state. When radical reactions are involved, one electron exchange with the catalyst serves as chain initiation. Little attention is paid to the nature and structure of the active site and the role of its environment. The solid state chemist approaches the problem in a different way(2). His main interest focus on the phase composition of the solid, type of crystal planes exposed, presence of additives and impurities, oxidation states of the cations and their changes in the course of the reaction, type of defects in the oxide lattice, etc. Correlation is sought between these parameters and the activity and selectivity of the oxide system in the given reaction, but little attention is usually paid to the type of interactions between the hydrocarbon molecule and the surface and to the possible transition states. When these two approaches are integrated, several general conclusions may be formulated, but also a number of important yet unanswered questions emerge. 2. DYNAMIC

INTERACTIONS

AT OXIDE SURFACES

On the formation of the oxide surface dangling bonds and the dangling bonds charge density appears in an energetically unfavourable conformation. Therefore, a driving force is generated at the surface to redistribute this charge density so as to create an insulating surface (pair up the dangling bond electrons and open a gap between the occupied and unoccupied surface states). This can be done structurally by rearranging the surface atoms, electronically through strong electron correlation effects or by forming new bonds at the surface of the oxide between surface atoms and the adsorbate molecules (3). Experimental evidence (4-6) and results of quantum chemical calculations (7-9) indicate that the oxide surface is in dynamic interaction with reactants of the catalytic reaction. The interaction of an adsorbing molecule with the surface not only produces the changes of the structure

of the adsorbate, but induces also reconstruction of the oxide surface. The surface adapts itself to the requirements of the reaction, generating a more facile pathway for a concerted rearrangement of electrons and nuclei. Ample experimental evidence (10,11) shows that the same is true for metal surfaces. Reconstruction of the surface of heterogeneous catalysts under the influence of a reaction mixture demonstrates the validity of Newton's third law in heterogeneous catalysis (12). Thus, the heterogeneous catalytic system should be considered not as a twophase, but rather as a three phase system, composed of the gas phase reactants, the solid catalyst and the surface region built of adsorbed molecules interacting with the surface layer of the solid (13). The reacting molecules and the atoms of the solid surface constitute one quantum chemical system. The state of this surface region may be modified either by changing the parameters of the gas phase or by altering the parameters of the solid. Therefore, to obtain information about the mechanism of a heterogeneous catalytic reaction, "in situ" studies of all parameters are essential. In every oxidation reaction two reactants take part: the oxygen molecule and the hydrocarbon molecule to be oxidized. Molecular oxygen may be activated in different ways (14-16): by excitation to the singlet state or by transfer of electrons from the catalyst to the oxygen molecule to form molecular or atomic ion radicals 02or O-. All these forms are strongly electrophilic reactants. They may abstract hydrogen from the hydrocarbon molecule with the formation of alkyl radicals, which may start a chain reaction. Under mild conditions, e.g. in the case of the reaction in the liquid state the chain reaction leads to the formation of alcohols or carbonyl compounds. Under harsh conditions of heterogeneous reactions radical chain leads usually to total oxidation so that in oxidative coupling of methane the problem consists in finding reaction conditions, in which two methyl radicals combine before they can be attacked by oxygen (17). In reactions with olefins or aromatics the electrophilic oxygen species attack the regions of highest electron density - the ~bond system (18). Peroxo- and super oxo-complexes are formed, which decompose by C-C bond cleavage to give oxygenated fragments or undergo combustion. These reactions may be classified as electrophilic oxidations (19). The second route of heterogeneous oxidation starts with the activation of the hydrocarbon molecule by abstraction of hydrogen from a given carbon atom, which becomes prone to nucleophilic addition of the oxide ion O2-. It should be emphasized that the latter has no oxidizing properties, but is a nucleophilic reactant. The consecutive steps of hydrogen abstraction and oxygen addition may be then repeated to obtain selectively more oxygenated molecules. These reactions are classified as nucleophilic oxidation (19). The role of oxidizing agents in these steps of the reaction sequence is played by cations of the catalyst lattice. After the nucleophilic addition of the lattice. After the nucleophilic addition of the lattice oxygen ion the oxygenated product is desorbed, leaving a vacancy at the catalyst surface. 3. INTERACTION OF OXIDES WITH GAS PHASE OXYGEN Little is known about the population of the surfaces of transition metal oxides

with electrophilic oxygen species. Transition metal oxides are non-stoichiometric compounds, their composition depending on the equilibrium between the lattice and its constituents in the gas phase. Changes in oxygen pressure cause changes in stoichiometry of the oxide, which may be accommodated by the crystal lattice in two ways: either by generation of point defects or by alteration of the mode of linkage between the oxide coordination polyhedra, resulting in the formation of extended defects (crystallographic shear). When non-stoichiometry is introduced by the presence of point defects, a series of equilibria is established at the surface on the pathway of lattice oxygen from the bulk to the gas phase or in the reversed process (Figure 1).

02

-

0

2-

~

02

-

~-

02

0 2- M(n-1)+'~ 'Mn+ 0 2- Mn+ 0 2- Mn+ 0 2- [~ ~0

Mn + 0 2- Mn +

2-

0 2-..

Mn+ Mn+ 0 2- M n+ 0 2- M n+ 0 2- M n+

0 2- Mn+ 0 2- M(n-1)+[[[]

Mn

0 2- Mn+ 0 2-

Mn+ 0 2-

Figure 1. Oxygen equilibria at the surface of an oxide When the temperature increases, the equilibrium shifts in the direction of higher dissociation pressure of the oxide and the surface becomes more and more populated with electrophilic oxygen species. When used as catalysts in oxidation of hydrocarbons, such oxides may show high selectivity to partial oxidation products at low temperatures, wherein the surface coverage with transient electrophilic oxygen forms is low. Under such conditions the conversion is very low. On raising the temperature the selectivity to partial oxidation products rapidly decreases, whereas the conversion in total oxidation increases, becoming the predominant reaction

pathway (4). It should be borne in mind that in the case of those oxides, in which the transition metal cation is not in its highest oxidation state, chemisorption of oxygen takes place at low temperatures, electrons being transfered from the oxide to adsorbed oxygen molecules with formation of higher valent cations and electrophilic oxygen species. This may be followed by surface reconstruction and formation of a monolayer of higher valent oxide at the surface of the lower valent one (4). Practically no experimental data exist on the mechanism of the dissociation of transition metal oxides. Little information can be found on the temperature dependence of the equilibrium oxygen pressure, and oxygen adsorption isobars are in most cases unknown. One should also bear in mind that doping of the oxide by altervalent ions may strongly influence its equilibrium oxygen pressure (20), that could permit a control of the activity of surface electrophilic species and hence the increase of selectivity. Different behaviour is shown by oxides, in which the change of stoichiometry is accommodated by the formation of shear structures. Because there is no vacancy formation on extraction of nucleophilic oxygen ion from the oxide surface due to its addition to the hydrocarbon molecule and desorption of the oxygenated product, but instead a shear plane is formed or an existing one grows (21,22), only nucleophilic oxygen species remain exposed at the surface. Very scarce data indicate that these oxides, in which the metal cation is in its highest oxidation state, do not adsorb oxygen. On raising the temperature the activity increases with growing mobility of oxygen, but the selectivity to partial oxidation products remains very high because of the absence of electrophilic oxygen. 4. ACTIVATION OF THE C-H BOND

The process of oxidation of a hydrocarbon molecule must begin with the activation of the C-H bond. Realization of this first elementary step is particularly challenging for physical chemists and chemical engineers because it must be achieved in the presence of many constraints. Namely, the C-H bonds in the initial reactant are usually stronger than those in the intermediate products, which makes these intermediates prone to rapid further oxidation and renders the C-H bond activation rate determining. Very little information on the molecular mechanism of the C-H bond activation at oxide surfaces can be found in the literature, and practically nothing is known about the activation of this bond in alkanes. Coordination of aikanes to transition metals has been observed both in liquid and in the gas phase (23), but the complexes are unstable at higher temperatures and their existence has been detected by matrix isolation technique. The alkane molecule can be coordinated to the transition metal through one, two or three hydrogen atoms forming C-H...M o--complexes, or it can be coordinated through its C-H bond forming with the electrophilic metal atom a two-electron, three-center bond, as in the triangular species H3+ formed from H2 and H+ in the gas phase. The interpretation was borrowed from the idea that carbonium ions such as C3HF+have non-classical bridged structures with two electron three center bonds (24). When the metal is capable of back donating

electrons to the antibonding C-H orbital, oxidative addition of the alkane to the metal may take place. Many ~-complexes are in equilibrium with their oxidative addition products (25). A growing body of evidence supports the intermediacy of transition metal alkane complexes in solution, C-H activation and reductive elimination. I r, Rh and W complexes may be mentioned. Reactions in the gas phase of neutral transition metal atoms (Pd, Pt) and ions (first row transition metal cations M+) with different alkanes and alkenes have been demonstrated and the metal-alkane complexes have been implicated as intermediates in the dehydrogenation. A question must be raised as to whether this type of mechanism could operate in the activation of hydrocarbon molecules at the surface of oxide catalysts. Very scarce experimental data seem to indicate that alkanes are adsorbed at oxide surfaces weakly and only at low temperatures. Stronger is the adsorption of olefins, which form surface ~-complexes with coordinately unsaturated metal ions characterized by Lewis acid properties. Activation of the CH bond is usually the rate determining step and the reactions are first order in respect to hydrocarbon indicating that the non-dissociative adsorption is absent. At the surface of oxides of transition metal ions in the highest oxidation state, which are usually components of oxidation catalysts, the mechanism of back-donation weakening the C-H bond cannot operate, and indeed pure MoO3 or WO3 are inactive in the conversion of hydrocarbons. They are, however, known to perform the addition of nucleophilic oxygen efficiently once these molecules have been activated. Some other oxide component must be introduced, e.g. Bi203, SnO2 etc. to activate the hydrocarbon molecule and render the catalyst both active and selective (18, 21, 67). In processes of heterogeneous catalytic oxidation, heterolytic C-H bond cleavage on the acid-base pair of sites is usually considered. Two possibilities are taken into account: - abstraction of hydrogen in the form of a proton and formation of a carboanion fragment. ~/ C-(T ----- H+~ i ~ M+n O-2 abstraction of hydrogen in the form of a hydride ion and formation of a positively

charged hydrocarbon fragment ~ !

M+n

0-2

Three mechanisms of homolytic bond rupture are also possible, two with the generation of radicals: - when the transition metal ion has electrophilic properties, and a vacant coordination site, a two-electron three-center bond between the C-H (~ bond and

the metal can be formed, and when the metal contains non-bonding d-electrons which can be back-donated, oxidative addition takes place: C .... H ~' 9 d

Mn+

.~ "~'

"

C M(n+2)+

when an easily reducible cation is present at the surface with basic properties, C-H bond cleavage may take place with the formation of an alkyl radical a proton with simultaneous reduction of the metal cation. C ~

H+ M+n

when atomic or molecular oxygen radicals are formed at the surface, they may abstract the hydrogen atom generating alkyl radical: C'~

~-H

o: Such a mechanism is postulated to operate in the activation of methane at high temperatures in the process of the oxidative coupling (65). Catalysts which are both active and selective for the oxidative coupling of methane may be classified as strongly basic metal oxides. Substitution of Iower-valent cations in their lattice generates oxygen vacancies, which constitute electron acceptor levels and are responsible for the appearance of electron holes in the valence band. These holes diffuse to the surface, because the lone pair orbitals of surface oxide ions are the HOMOs of the oxide and their energy levels form the top of the valence band. Localization of a hole on such lone pair orbital is equivalent to the formation of a surface O- species. On superacidic catalysts the formation of a carbenium ion is suggested to be the first step of the C-H bond activation R\C / R

/i\

H H+H O The data on hydrogen/deuterium exchange in paraffins on oxides indicated that a negatively charged intermediate is formed (26,27). Also the studies of adsorption of ethane and propane on oxide catalysts by I R spectroscopy showed that heterolytic dissociation of the C-H bond takes place with the formation of negatively

charged alkyl fragment (28). Indirect evidence of the abstraction of a proton as the first step in oxidation of hydrocarbons comes from experiments, in which catalytic activity was measured on series of catalysts with different acid-base properties. It has been found that the rates of oxidative dehydrodimerization of propene to 1,5hexadiene and the rates of its oxidation to acrolein increased with decreasing binding energy of O ls electrons and with decreasing difference between binding energy and Auger kinetic energy (29) of SnO2-based catalysts, which were due to the increasing negative charge on the oxygen ions i.e. an increasing basicity. Similar conclusions were derived (30) from studies of the oxidation of propene on SnO2 impregnated with acidic (P205) and basic (Na20) oxides. Optimum basicity was found to exist, reflecting the requirement that the oxide efficiently abstracts protons from hydrocarbon molecules in the first step of the reaction and releases the protons easily in the dehydroxylation step. Recently a distinct correlation was found (31) to exist between the rate of the oxidation of butane on vanadyl pyrophosphate catalysts, doped with alkali and alkaline earth cations, and the negative charge on oxide ions as determined by XPS. On the other hand no experimental evidence exists, except of the case of zinc oxide (32), of the formation of hydride ions in the course of adsorption of hydrocarbons on oxides. The activation of the C-H bond in CH4 has been analyzed on the basis of quantum-chemical calculations using the ASDED-MO (Atom Superposition and Electron Delocatization Molecular Orbital) approach (66). It has been concluded that on oxide catalysts the oxidative addition of the C-H bond to the metal cation requires a very high activation energy because the antibonding orbital, which must accept two electrons, lies very high. When a hole appears in this orbital, the reaction becomes more facile. In the series of consecutive elementary steps of the selective oxidation of a hydrocarbon molecule, the negatively charged alkyl, formed after abstraction of a proton, must become bonded to a surface oxide ion of the catalyst to form the alkoxy-species, generally accepted, to be next intermediates. Their appearance at the surface of oxide catalysts in the course of the selective oxidation of hydrocarbons has been proved experimentally by many techniques, e.g. in-situ IR and Raman spectroscopies (33). No attention was paid to date to the fact that a molecule approaching the surface of an oxide enters into the region of a strong electrostatic field stretching above this surface. Indeed, ab initio HF calculations carried out for V209and V209H8clusters, modelling the surface of V205 catalysts, showed (34) that a strong negative electrostatic field is surrounding the cluster. It may be thus hypothesized that in the oxidation of hydrocarbons at oxide surfaces the activation of the C-H bond starts with the polarization of the molecule by this electrostatic field. Semiempirical quantum-chemical calculations (35) of the interactions developing on approach of a propene molecule to the V10031H12 cluster taken as a model of V205 catalysts showed that at first a bond is formed between the carbon atom of propene and the bridging oxygen atom of the cluster. This is followed by abstraction of a proton with simultaneous injection of electrons into the empty d-levels of the cluster. A general conclusion may be advanced that in the activation of a C-H bond at the surface of a

transition metal oxide with semiconducting properties both parts of the cleaved C-H bond become attached to surface oxide ions. The proton forms an OH group, the hydrocarbon fragment forms an alkoxy group. Simultaneously the two electrons of this bond are injected into the conductivity band of the solid. This process may be represented by a scheme:

-•C+• 0-2

,,,

H+,,

M+

0-2

i

The alkoxy group then loses a second hydrogen and desorbs as an aldehyde or ketone. This series of transformations forms the first part of the Mars-van Krevelen redox mechanism, which is then followed by adsorption of oxygen from the gas phase, transfer of electrons from the solid to adsorbed oxygen molecule and incorporation of oxygen ions into the lattice of the oxide, which completes the redox cycle. The cycle involves two adsorbed redox couples: RH + O21/2 02 -!- 2e-

~ -~

R-O-+ H+ + 2e0 2-

of which the first injects electrons into the oxide, the second one extracts them from the oxide. The conditions of the electron transfer across the gas/solid interface are seldom considered in catalysis, but they are discussed in detail in electrochemistry of semiconductors and the same rules must be valid in the case of a catalytic reaction. The injection of electrons from an adsorbed redox pair into a semiconductor can take place spontaneously only if the redox potential of this pair is situated above the Fermi level and above the bottom of the conductivity band and the extraction of electrons from the solid - when the redox potential is located below the Fermi level and below the top of the valence band (Figure 2).

10

OROCARBON...... . I . . . .

CATALYTIC OXIDATION OF FIYDROCARBON MOLECULE CAN PROCEED

~

OXYGEN O~YGEN.~~~~,,.~'~~ REDOX

CATALYTIC oXIDATION OF HYDROCARBON MOLECULE CANNOT PROCEED, BECAUSE THE MOLECULE IS NOT ACTIVATED

]

CATALYTIC OXIDATION OF H Y D R O C A R B O N MOLECULE CANNOT PROCCED BECAUSE THE CATALYST IS NOT REOXIDIZED

Figure 2: The transfer of electrons across the adsorbate/oxide interface in the course of the heterogeneous catalytic oxidation of a hydrocarbon molecule by gas phase oxygen. The probability of these processes is a function of the density of states in the conductivity and valence band respectively at the potentials corresponding to the redox potential of the adsorbed species. In the electronic theory of catalysis in the later fifties only general thermodynamic rules were considered, but the conditions of the electron transfer were never applied to analyze the behaviour of different oxide catalysts in the oxidation of hydrocarbons. No quantitative data are available to make such analysis, although recently an attempt was made to interpret the changes of the rate of ethanol oxidation in terms of the density of states (36). The relative positions of the energy bands in the solid and the redox potential of the reacting molecule may be adjusted by: a) formation of one or more oxide/oxide interfaces with such values of the contact potentials that the conductivity band will

]! shift to the optimum position, b) doping of the oxide with altervalent ions, which will shift the Fermi-level, c) generation of surface defects, which will create a broad distribution of surface electronic states participating in the exchange of electrons with the reacting adsorbed molecules (37). 5. STRUCTURE SENSITIVITY AND THE ROLE OF DEFECTS

In the analysis of the mechanism of homogeneous catalytic reactions catalyzed by organometallic complexes, the number of vacant coordination sites at the transition metal cation is taken into account, required for the reactants to be coordinated (38). In the heterogeneous catalytic reaction, taking place at the surface of the solid the number of vacant coordination sites can vary either because of the location of the cation at different crystal faces or as the result of the generation of lattice defects. No information is available as to the role of either of these factors. Ample experimental evidence indicates that the habit of crystallites of the catalyst has a profound influence on the activity and selectivity of the reaction, which results in the appearance of structure sensitivity of the selective oxidation reactions at oxide catalysts (39). However, little is known about the origin of this phenomenon and about the differences of the structure of active sites present at various crystal planes. Even less is known about the role of defects present at the surface of an oxide, in determining the catalytic properties. Only recently studies of the properties of (100) surface of a monocrystal of NiO revealed that an ideal surface is chemically inert and the reactivity of the system increases only if defects are introduced in the surface (40). At such a surface, dissociation of molecular water to form hydroxyl groups is observed in contrast to an ideal surface which is inactive in water dissociation. 6. OXIDE

MONOLAYER

CATALYSTS

In recent years considerable attention has been paid to oxide monolayer catalysts, obtained by deposition of a transition metal oxide in submonolayer up to few monolayers coverage on the surface of another oxide of a main group or transition metal playing the role of a support. In the submonolayer range, isolated transition metal-oxygen polyhedra are anchored at the surface of the oxide support. Recent investigations showed (41) that under ambient conditions the surface molecular structures of transition metal oxy-ions are directly related to the surface pH at point of zero charge of the aqueous film on the oxide supports and can be predicted from the corresponding aqueous transition metal ions chemistry (42). Depending on the type of the system and the method of preparation transition metal ions condensation to form clusters, cover completely the support as a surface layer or form small crystallites of a second phase (43). The minority oxide (supported phase) can accumulate entirely at the surface of another oxide (support) when the temperature of annealing is low enough and the miscibility as well as chemical affinity of the two oxides are very limited. In the case of miscibility of the two oxide phases they may also be incorporated into to outermost surface layer of the support (surface framework) or diffuse into the bulk and form a solid solution at a high

12 enough temperature. In the case of chemical affinity of the two oxides new surface or bulk compounds may be synthesized. In the case of preparation by impregnation from the solution the oxide surface may play the role of solvating agent, counter ion or a macroligand (43) as illustrated in Figure 3.

I I

TRANSITION METAL ION IN SOLUTION

TRANSITION METAL ION

1 ISOLATED I OXO SPECIES I

, , I

/

~

g

.-o o-. i

,,% ~

I

SOLVE,NT

/

J

\~ o~)

~

OXIDE SURFACE PLAYS THE

/

(,,-~,,'F~I

",,. 2- 7,-J /

COUNTER ION

\

OF

ROLE

/

J

77

/

~

L,

/

/P

/

IN BIDIMENTIONAL POLYOXO METALATE CLUSTER

I

TRANSITION I METAL ION IN GAS PHASE I I~/L I

I

J

~

[]

<, <-, Y

/

LIGAND

/ / / / J

.,Y,.

,

O-M-O

J

I

IN OXIDE CRYSTALLITE

.,

,

I

J

1

/

FRAMEWORK

-o-' -o-~

o

o o o I

.,

J ~

o

Q

i

1

I

I

~'-,~-'~

J

.

~SOLI'~SOLUTION

J / /

6-~','9~ o

<,% ~' o- o

J

o,~o o;%

1

I

J

..

i

/

~

l

I

/ ./

/

Figure 3. Deposition of transition metal ions by impregnation from their aqueous solution on the oxide surface and their further transformations. As the solid support is a fairly rigid macroligand, mono- or polydentate, it distorts the environment of the transition metal ion and generates new active sites exhibiting specific properties. On increasing the loading of the support with the active phase, islands of the monolayer may at first be formed and three dimensional clusters of the active phase may grow either simultaneously or after completion of the monolayer. It should be borne in mind that the active phase may not be uniformly distributed over the whole surface of the crystallites of the support, but may be

]3 preferentially deposited on certain crystal planes only, giving rise to the phenomenon of structure-sensitivity of impregnation [44]. Properties of the active phase may be in different ways affected by supporting on different crystal planes. It has been shown in recent years that an equilibrium between hydrated (Fig. 3, form III) and dehydrated (Fig. 3, form IV) metal-oxygen polyhedra is established at the surface of a support depending on the water vapour pressure, the hydrationdehydration process being reversible [44]. It may be anticipated that the two types of species would have different catalytic properties, these properties being thus strongly dependent on the water vapour pressure. Moreover, the change of the water vapour pressure shifts the equilibrium of surface hydroxylation, which changes the ratio of Bronsted-to-Lewis acid sites. As yet, little is known about the influence of water on the behavior of the oxide catalysts, although in many industrial oxidation processes water vapour is introduced into the stream of reactants to improve selectivity. Further studies are required to unravel the mechanism of the influence of water on the surface properties of oxides. 7.

WETTING OF OXIDE SURFACE BY OTHER OXIDES

At higher temperatures thermal spreading may occur resulting in the formation of an oxide monolayer [46-48]. Spontaneous spreading of one oxide over the surface of another oxide is the manifestation of the wetting of one solid by another solid due to the operation of the forces of surface tension. It occurs when the free energy of adhesion of the mobile phase to the support is greater than the free energy of cohesion of the mobile phase and continues until the formation of a thermodynamically stable overlayer is completed, characterized by a surface free energy equal to that of the bulk of the active oxide phase. The rate of spreading is limited by surface diffusion described by a parabolic rate equation and depends strongly on the crystallinity of the support and type of the atmosphere [49]. When the monolayer is deposited on a support which is not wetted, coalescence of the monolayer into crystallites of the deposited transition metal oxide takes place on annealing. As the free energies of cohesion and adhesion depend on the oxidation state of the oxide monolayer, reduction and reoxidation entail strong changes of the degree of dispersion. It should be also remembered that the surface free energy is very sensitive to the presence of additives (impurities). Therefore wetting of the surface of the support by the active phase may be controlled by the introduction of appropriate additives, which either remain at the surface or diffuce into the subsurface layer. Unfortunately, practically no experimental data are available on the surface free energy of oxides and their dependence on the type and concentration of defects. A simplified method to calculate the surface free energy on the basis of the ionic model of solids has been proposed some time ago [50]. In recent years attempts are being undertaken to introduce the clusters of transition metal oxides into frameworks, which could control the spacial arrangement of molecules reacting at the active sites provided by the oxide. To this end pillared clays are synthesized, in which either the pillars are built of the active transition metal oxide phase, or this phase id deposited onto the pillars, composed of an inert

14 material. Introduction of clusters of transition metal oxides into the zeolite framework has also been reported. Such "ship-in-the-bottle" catalysts are now the subject of considerable interest. 8.

NUCLEOPHILIC

ADDITION OF OXYGEN

The fundamental question, which has not been answered yet, concerns the properties of the oxide required to perform the nucleophilic addition of oxygen to the hydrocarbon molecule. In the course of the desorption of the oxygenated hydrocarbon molecule as a product of the reaction, an oxygen atom becomes extracted from the surface of the oxide catalyst, leaving behind an oxygen vacancy. Simultaneously, hydrogen atoms abstracted from the hydrocarbon molecules and present at the surface in the form of hydroxyl groups, must be removed by dehydroxylation, which also generated surface oxygen vacancies. It is a well known fact that all active and selective catalysts for partial oxidation of hydrocarbons are based on group V, VI or VII transition metal oxides. These oxides show a strong tendency to eliminate the vacancies by the formation of shear planes, which are nucleated at the surface with the simultaneous release of oxygen by the crystal. A hypothesis was advanced [51-53] that this tendency is a driving force facilitating the desorption of the oxygenated product. An easier and efficient route is thus provided for the addition of a nucleophilic lattice oxygen into the hydrocarbon molecule. It should be however borne in mind that crystallographic shear may not be the only pathway of nucleophilic insertion of oxygen into the hydrocarbon molecule. A number of oxide systems are known, which show good activity and high selectivity in selective oxidation, but do not undergo crystallographic shear. One can mention heteropolyacids and their salts [54] and the important class of oxide monolayer catalysts. With respect to the monolayer catalysts an important question has been raised as to the nature of active sites at the surface of such catalysts. As already mentioned, depending on the surface coverage of the support, the monolayer is composed of isolated MOx polyhedra, regions of a monolayer composed of clusters of condensed polyhedra or three dimensional crystallites of the active phase. Moreover, transition metal ions of the monolayer diffuse into the subsurface region forming a solid solution. Experiments show that part of the monolayer can be dissolved in appropriate solvents, but part is strongly bound [55-57]. In spite of ample experimental data there is no agreement as to which is the nature of the soluble part or the structure of active sites, responsible for the catalytic reaction. If the activity were related to the presence of loosely bound clusters of condensed MOx polyhedra, as indicated by some experiments [58], one could imagine a local rearrangement from corner-linked to edge-linked set of MOx polyhedra, resembling the nucleation of a shear plane, to be responsible for the nucleophilic addition of oxygen. Experiments show that on vandium oxides supported on titania, butane is oxidized to maleic anhydride, whereas on vanadium oxides supported on magnesia it undergoes oxidative dehydrogenation to butene. The mechanism, by which the support changes the properties of vanadium oxides so dramatically remains a fascinating yet unaswered question.

].5 As the phenomenon of crystallographic shear appears in transition metal oxides with anisotropic lattices, pronounced structure-sensitivity of catalytic properties is observed and the habit of crystallites of the catalyst may have strong influence on the selectivity of the reaction [39,59-61]. Multiple examples of the dependence of catalytic properties on the type of exposed crystal plane have been described in literature, but the only attempt to explain this phenomenon in terms of the molecular structure of different crystal planes was undertaken in the crystallogrphic model of active sites [62]. The question awaits more dedicated experiments and deeper theoretical analysis. The first quantum-chemical approach addressing this question has been quite recently published [63]. 9.

CONCLUSIONS

Quantum chemical calculations have already permitted some general conclusions concerning the mechanism of heterogeneous catalytic oxidation [64]. In electrophilic oxidation the type of product formed depends on the direction of the approach and mutual orientation of the hydrocarbon molecule and oxygen, and on the mode of oxygen activation. In nucleophilic oxidation by surface lattice oxygen the site of the hydrocarbon molecule attacked and hence the reaction pathway followed is critically dependent on the structure of the active site and the orientation of the molecule approaching the surface. The desired chemo- and regio-selectivity could be thus achieved by imposing a molecular field appropriately directing the incoming molecule. The incorporation of the molecular recognition function will add the reaction specificity. The ultimate goal of the science of catalysis is the accumulation of knowledge, necessary for molecular design of catalytic systems tailored to the needs of any imaginable catalytic reaction. Answers to the questions raised in this paper will make this goal possible, and hence will change the paradigm of catalysis. REFERENCES

1. G.A. Olah, A. Molnar, Hydrocarbon Chemistry, John Wiley Sons Inc., New York 1995. 2. Solid State Chemistry in Catalysis, eds R.K. Grasselli, J.F. Brazdil, ACS Symp.Series 279, American Chemical Society, Washington, 1985 3. J.P. LaFemina, Surf.Sci.Rep. 16 (1992) 133; Critical Rev.Surf.Chem.3 (1994) 297 4. J.Haber, Materials Sci. Forum 25 (1988) 17 5. J.Haber, Proc. 4th Inern Conf. Chemistry and Uses of Molybdenum, Golden CO 1982, eds H.F.Brry. P.C.W.Mitchell, Climax Molybdenum Co, AnnArbor 1982, p.395 6. H.J.Freund.B.Dillmann, O.Seiferth, G.Klivenyi, M.Bender, D.Ehrlich, I.Hermmerich,D. Cappus, Catal. Today 32 (1996) 1. 7. J. Haber. M. Witko, Catal. Today 23 (1995) 311.

16 8. M. Witko, R.Tokarz, J.Haber, Catal.Today in print. 9. M.J. Gillan, L.N.Kantorovich,P.S.D. Lindan, Current Opinion in Solid State Mat. Sci., 1 (1996) 820 10. M.A. vanHove, G.A. Samorjai, Surf.Sci. 299/300 (1994) 487 11. G.A. Samorjai, Catal.Lett 9 (1991) 311, Surf.Sci. 299/300 (1994) 849 12. K.l.Zamaraiev, Topics in Catal.3 (1996) 1. 13. J. Haber, in "Perspectives in Catalysis", eds J.M. Thomas, K.l.Zamaraiev, Blackwell Scientific Publ., Oxford 1992, p.371 14. M. Che, A.J. Tench. Adv. Catal. 31 (1982) 78;32 (1983) 1. 15. Z.Sojka, Catal. Rev.Sci. Eng. 37 (1995) 461. 16. J. Haber, T. Mlodnicka, J. Molec. Catal. 74 (1992) 131 17. J.Lundsford, Stud.Surf.Sci.Catal., vol 75, Elsevier, Amsterdam 1993, p.103 18. J. Haber, in Proc. 8th Intern.Congr.Catalysis,Berlin 1984, Verlag ChemieDechema 1984, vol.1 Plenary Lectures, p.85 19. J.Haber, in "Solid State Chemistry in Catalysis", eds.R.K.Grasselli, J.F.Brazdill, ACS Symp.Series 279, American Chemical Society, Washington 1985, p.3. 20. F.A.Kroger, The Chemistry of Imperfect Crystals, 2nd Ed., North Holland Publ. Co, Amsterdam 1973. 21. J.Haber, in "Surface Properties and Catalysis by Non-metals", eds. J.P.Bonnelle, B. Delmon, E. Derouane, D.Reidel 1983, p.l. 22. J. Haber, J. Janas, M. Schiavello, R.J.D. Tilley, J.Catal. 82 (1983) 395 23. C. Hall, R.N.Perutz, Chem.Rev. 96 (1996) 3125 24. G.A.Olah, G.U.S. Prakash, R.E. Williams,L.D. Field, K.Wade, Hypercarbon Chemistry, J. Wiley Sons Ltd, New York 1987. 25. R.H. Crabtree, Chem.Rev.95 (1995) 987 26. R.L.Burwell, A. Littlewood, M.Cardew, C.T.H. Stoddart, J. Am.Chem.Soc. 82 (1960) 6272 27. P.J.Robertson, M.S.Scurrell, C.Kemball, J.Chem.Soc. Chem.Comm 20 (1973) 799. 28. V.D.Sokolovskii, S.M. Aliev,O.V. Buyevskaya, A.A. Davydov, Catal.Today 4 (1989) 292 29. E.A. Manedov, Kinet.Katal. (russ)25(1984)868 30. T.Seiyarna, M.Ehashira, M.Iwarnoto, in "Some Theoretical Problems of Catalysis" eds T.Kwan, G.K. Boreskov, K.Tamaru. Univ. Of Tokyo Press, Tokyo 1973, p.35 31. V.Zazhigalov, J.Haber, J.Stoch, V.Batcherikova, Appl. Catal.A:General 134 (1996) 225. 32. R.J.Kokes in Proc.5th lntern.Congr.Catalysis, MiamiBeach FL 1972, ed J.W. Hightower, North Holland/American Elsevier, New York 1973,p.A.1 33. F.Finocchio, G.Busca,V.Lorenzelli,R.J.Viley,J.Catal. 151 (1995) 204 34. M.Witko, Catal.Today 32 (1996) 89 35. J.Haber, M.Witko, R.Tokarz, in print 36. W.Zhang, SA.Desikan, S.T.Oyama, J. Phys.Chem. 99 (1995) 14468 37. I.Manassidis, J. Goniakowski, L.N. Kantoroich, M.J. Gillan, Surf.Sci. 339 (1995) 258 38. R.A.Sheldon, J.K.Kochi, Metal-Catalyzed Oxidations of Organic Compounds,

17 Academic Press, New York 1981 39. J.Haber, Stud.Surf.Sci.Catal.Vol.48 Elsevier, Amsterdam 1989, p447. 40. M.Baumer, D. Cappus, H. Kuhlenbeck, H.J. Freundm G. Wilhemi, A. Brodde H. Neddermeyer, Surf. Sci. 252 (1991) 116. 41. G. Dep. I.E. Wachs, J. Phys. Chem. 95 (1991) 5889 42. I.E. Wachs, G. Deo, D. S. Kim, M. A. Vuurman, H. Hu, Stud. Surf.Sci. Catal., vol.75A, Elsevier, Amsterdam 1993, p. 543. 43. M. Che, L. Bonneviot, Pure Appl. Chem. 60 (1988) 1369 44. K. Bruckman, J. Haber, T. Wiltowski, J. Catal. 106 (1987) 188 45. G. Deo, I.E. Wachs, J. Haber, Critical Rev. Surf.Chem. 4 (1994) 141 46. J. Haber, T.Machej, T.Czeppe, Surf. Sci 151 (1985) 301. 47. H. Knozinger, E. Taglauer, Specialist Period. Reports "Catalysis" vol. 10, The Royal Soc. Chem., London 1993, p.l. 48. J. Haber, Pure Appl.Chem. 56 (1984) 1663 49. J.Haber, T. Machej, E.Serwicka, I.E.Wachs, Catal.Lett. 32 (1995) 101 50. J. Ziolkowski, Surf.Sci. 209 (1989) 536 51. F.S.Stone, J. Solid State Chem. 12 (1975) 271. 52. J. Haber, in Proc 34rd Intern.Conf.Chemistry and Uses of Molybdenum, eds. H.F. Barry, P.C.H. Mitchell, Ann Arbor, MI 1979, Climax Molybdenum, Ann Arbor 1979, p. 114. 53. E.Broclawik, J.Haber, J.Catal. 72 (1981) 379 54. M. Missono, T.Okuhara, N.Mizuno, Adv.Catal. 41 (1996) 113. 55. G.Centi, E.Giamello, D.Pinelli, F.Trifiro, J.Catal. 130 (1991) 220;ibid.238

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

19

The multifunctional properties of heterogeneous catalysts, active and selective in the oxidation of light paraffins F. Cavani and F. Trifir6 Dipartimento di Chimica Industriale e dei Materiali, Viale Risorgimento 4, 40136 Bologna, Italy This review analyzes the properties which are necessary for heterogeneous catalysts to promote the oxyfunctionalization of light paraffins to valuable chemicals. Three catalytic systems are discussed: i) vanadium/phosphorus mixed oxide, the industrial catalyst for the oxidation of n-butane to maleic anhydride, which is here also examined for reactions aimed at the transformation of other hydrocarbons; ii) Keggin-type heteropolycompounds, which are claimed for the oxidation of propane and isobutane, whose composition can be tuned in order to direct the reaction either to the formation of olefins or to the formation of oxygenated compounds; iii) futile-based mixed oxides, where rutile can act as the matrix for hosting transition metal ions or favour the dispersion of other metal oxides, thus promoting the different role of the various elements in the formation of acrylonitrile from propane.

Introduction In recent years a significant interest has been directed towards the study of new catalytic materials able to promote the transformation of alkanes to valuable chemicals by means of selective oxidation (1-8). This interest arises from the possibility of replacing the currently used raw materials, mostly olefins and aromatics, with cheaper and more environmentally friendly organic substrates. The scientific challenge of the activation and selective functionalization of a saturated molecule thus is coupled with the economic and environmental benefit which would arise from the exploitation of natural gas components. These efforts have successfully turned into the commercial application of the process for maleic anhydride synthesis from n-butane, which is replacing the process from benzene, and which is still the only example of a gas-phase industrial oxidation of a paraffin for the synthesis of an intermediate for the petrochemical industry. This has been made possible by the elucidation of the unique properties of vanadyl pyrophosphate, (VO)2P207, which is able to promote the selective transformation of n-butane (9-13). On the other hand, in the last few decades several companies have claimed the development of catalytic systems and processes for the oxyfunctionalization of light paraffins. Table 1 summarizes the various reactions which are possible starting from light alkanes. These processes are in various stages of industrial application applied industrially. Some have been developed but not applied, and others are still at the research stage. In more recent years, a proliferation of papers and patents devoted to the study of catalysts for the oxidative dehydrogenation of paraffins to the corresponding olefins is a clear indication of the interest of industry in these processes, as an alternative to the energy-intensive endothermic steam cracking or dehydrogenation of natural gas or oil components or fractions (5,14-17). It is possible to summarize the important aspects which need to be studied in the development of a catalytic material for the oxyfunctionalization of a paraffin by the following questions:

20 1) What is the mechanism of activation of a paraffin, or what is the best way to activate a paraffin and subsequently transform the intermediate to the desired product ? Which oxygen species is needed to favour this oxidative and selective conversion ? 2) Does the mechanism of oxygen insertion into the activated hydrocarbon involve an intermediate step of formation of an olefin ? Is this step necessary for achieving the highest selectivity to the desired product, or do alternative and more selective pathways exist, which do not involve the formation of olefinic species ? 3) To what extent does the stability of the product affect the selectivity at reaction conditions 9 4) Does the presence of heterogeneously-initiated, gas-phase reactions contribute to the reaction mechanism ?

Table 1. Industrial processes and processes under study or development for the oxyfunctionalization of light paraffins (C 1-C6) in the petrochemical industry. Raw Product Phase Stage of development material Methane Chloromethanes Gas, heterog. Industrial Methane Methanol Gas, het./hom. Pilot plant Methane Syngas Gas, het./hom. Research Methane Ethylene Gas, het./hom. Pilot plant Ethane 1,2 Dichloroethane, Vinyl Gas, heterog. Pilot plant chloride Ethane Acetic acid Gas, heterog. Research Ethane Ethylene Gas, het./hom. Research Propane Acrylic acid Gas or liquid Research Propane Propyl alcohol Liquid, het. or hom. Research Propane Acrylonitrile Gas, heterog. Demonstrative plant Propane Propylene Gas, heterog. Research n-Butane Acetic acid Liquid, homog. Industrial n-Butane Maleic anhydride Gas, heterog. Industrial n-Butane Butadiene Gas, heterog. Industrial, abandoned Isobutane Methacrylic acid Gas, heterog. Pilot plant Isobutane Isobutene Gas, heterog. Research Isobutane t-Butyl alcohol Liquid, het. or hom. Research n-Pentane Phthalic anhydride Gas, heterog. Research Cyclohexane Cyclohexanol, one Liquid, homog. Industrial Cyclohexane Cyclohexanone Liquid, het. or hom. Research Het./hom. indicates the likely presence of a mechanism initiated on the catalyst surface and transferred to the gas phase.

Some of these questions, amongst others, still need to be answered; however, the following points can be considered as definite ones: 1) The very different reactivity of methane and ethane with respect to the heavier paraffins requires completely different classes of catalysts as well as different reaction conditions which are needed for the reactant activation and transformation. Completely different conditions and catalysts imply different mechanisms. With methane and ethane (the

21

transformation of which is catalyzed by alkali and alkaline earth oxides), the contribution of homogeneous radical reactions is fundamental, while it is less important for higher paraffins and transition metal oxide-based catalysts. 2) The generally accepted rule that the olefin is more reactive than the corresponding paraffin is not always true. This is the case for ethylene and also in some cases for propylene. The presence of these olefins may substantially lower the selectivity to oxygenated compounds, in cases where the latter are the preferred compounds. 3) The most selective reactions are those which lead to the formation of stable oxygenated compounds, such as anhydrides (this is the case of the n-butane and n-pentane oxidation). On the contrary, products such as the unsaturated acids and aldehydes are unstable under reaction conditions which are necessary for the activation of the paraffin; and hence, low selectivities are observed. 4) Finally, it is clear that the many aspects related to the successful development of a new process for the oxidation of a light paraffin necessarily involve several disciplines and expertise, to join together the physico-chemical, material, kinetic and technological aspects. From the materials point of view, it is now clear that there are no "simple" catalysts which can be used for the partial oxidation of a paraffin. All of the best systems are either constituted by several different components which cooperate in achieving the final product, or made of multicomponent structures which perform the different functions for selective conversion. These different functions seem, therefore, to be a necessary feature for the activation of the C-H bond in the saturated substrate, and for the multi-electron oxidation and oxygen insertion onto the latter. It is more or less clear that the cooperation of acidic and suitable oxidizing properties is a necessary condition. However, this is not sufficient, and best catalytic performances are obtained only when these properties are structured in such a way to allow the final product to be quickly reached, and rapidly desorbed before further oxidative transformations may lower selectivity. The aim of this review is to describe the reactivity of three catalytic systems which have been widely studied in recent years for the oxidative transformation of light paraffins: i) vanadyl pyrophosphate, which is the industrial catalyst for the oxidation of n-butane, but has also been claimed to be selective in the oxidation of n-pentane to maleic and phthalic anhydrides (18-22), ii) heteropolycompounds, which are currently being studied for the oxidation of isobutane and propane to the corresponding unsaturated acids (methacrylic acid and acrylic acid) (5,23-29), and whose composition can be tuned to change the acidic and oxidizing properties; and iii) futile-based mixed oxides, which can act as the matrix to host various metal components, and which have been claimed as optimal catalysts for the ammoxidation of propane to acrylonitrile (15,30-33).

The oxidation of alkanes on V/P/O catalysts Table 2 reports the maximum yield to the oxygenated products from each paraffin which has been rePorted in the literature over V/P/O-based catalysts. In all cases the catalyst is mainly constituted by vanadyl pyrophosphate, which however, may contain either metal dopants or other vanadium phosphates. The selectivity to the product of partial oxidation is a function of the structure of the reactant. From n-butane and n-pentane the selectivity to maleic anhydride and to maleic plus phthalic anhydrides, respectively, is high, while from ethane the prevailing products are either ethylene or carbon oxides (depending on the reaction conditions); acetic acid is formed in rather low amounts. From propane very low amounts of acrylic acid are formed, and carbon oxides prevail. These differences can be attributed to the formation of very stable products

22 (the anhydrides) from C 4 and from C 5 paraffins, while acrylic acid (from propane) undergoes consecutive reactions of oxidative degradation. The low selectivity to acetic acid from ethane is attributed to the low specificity of the active sites in this reaction, rather than to the instability of the product, since acetic acid can be regarded as a stable compound.

Table 2. Maximum ~,ields to the desired products in the oxidation of light alkanes.

Reactant

Catalyst comp.

Desired product

Max. yield, tool.% (T, ~

Ethane Propane Propane n-Butane

V-P-O/TiO 2 V-P-Te-O V-P-O V-P-O (+Zn, Li, Mo, Zr) V-P-Co-O V-P-O

Acetic acid Acrylic acid Acrylonitrile Maleic anhydride

1.2 (350) 8 (400) 6 (550) 56-60 (380-420)

34 35 36 5

Methacrylic acid, Methacrolein Phthalic anh., Maleic anh.

4.5, 3.1 (280) 23, 22 (350)

37 18

Isobutane n-Pentane

Ref.

It has been proposed that the transformation of n-butane to maleic anhydride involves the following steps at the adsorbed state: 1) n-butane + 1/2 0 2 --~ butenes + H20 (oxidative dehydrogenation) 2) butenes + 1/2 0 2 ~ butadiene + H20 (allylic H-abstraction) 3) butadiene + 1/2 0 2 ~ 2,5-dihydrofuran(1,4 oxygen insertion) 4) 2,5-dihydrofuran + 2 0 2 ~ maleic anhydride + 2 H20 (allylic O-insertion, possibly via y-but-2-enoic lactone) or, 4a) 2,5-dihydrofuran + 1/2 0 2 --~ furan + H20 (allylic H-abstraction) 5a) furan + 3/2 0 2 ~ maleic anhydride + H20 (electrophilic oxygen insertion) Other proposed mechanisms involve either a direct attack of 0 2- species at the 1,4 C atoms of n-butane (38), or an allylic oxidation of an olefinic-like C 4 intermediate to crotonaldehyde, followed by internal cyclization and oxidation (39). A dienic intermediate has also been proposed by Grasselli et al. (40). In any case, the reaction patterns proposed evidence the need for different kinds of active sites on the surface of the vanadyl pyrophosphate which are able to perform with high selectivity each step in the reaction pathway. The polyfunctional nature of the vanadyl pyrophosphate is clearly evidenced by the different classes of reaction which can be catalyzed by this material, as summarized in Table 3. Vanadyl pyrophosphate can catalyze most of these transformations with good selectivity. This also indicates that the multiple steps proposed for the mechanism of n-butane oxidation to maleic anhydride can effectively occur on vanadyl pyrophosphate. The multifunctional properties of the vanadyl pyrophosphate can be summarized as follows: 1) This system possesses sites able to perform the oxidative dehydrogenation of paraffins. This is demonstrated by the formation of benzene with high specificity from cyclohexane, as well as by the formation of olefins and diolefins from n-paraffins, of aromatic compounds

23 from decaline, and of cycloolefins from cycloparaffins. However, this is not the optimal system for the synthesis of light olefins from paraffins.

Table 3. Classes of reaction catalyzed b~' the vanadyl pyrophosphate Reactant Product Reaction type isobutyric acid methacrylic acid oxidehydrogenation cyclohexane benzene succinic anhydride maleic anhydride hexahydrophthalic anhydride phthalic anhydride paraffin olefin olefin diolefin allylic oxidation (Habstraction or O-insertion) 2,5-dihydrofuran furan tetrahydrophthalic anhydride phthalic anhydride benzene maleic anhydride electrophilic oxygeninsertion butadiene furan naphthalene naphthoquinone furan maleic anhydride

2) It possesses centres able to perform the allylic oxidation with high specificity. This is the reason why vanadyl pyrophosphate does not yield olefins with high selectivity in the oxidation of paraffins. In fact the desorption of intermediate olefins is not rapid, since they are quickly transformed to the oxygenated products. Only when molecular oxygen is absent (and the catalyst possesses a low number of O-insertion sites), can the olefin be saved from consecutive transformations, and a good selectivity to the olefin can be achieved. 3) It possesses centres able to insert oxygen into electron-rich substrates. 4) It possesses acid centres. Acidity is recognized to play important roles in the activation of the paraffin, in the desorption of acid products, and in accelerating specific transformations over reactive intermediates (41). 5) It may favour bimolecular condensation reactions which are not acid-catalyzed, but which are accelerated by the proper geometry of sites at which the molecules are adsorbed. This property likely plays an important role in the mechanism of phthalic anhydride formation from n-pentane (19). 6) The oxidation state of vanadium under reaction conditions is a very important parameter in the control of the process selectivity; a certain number of oxidized vanadium sites is necessary to transform the intermediate olefins to the oxygenated products (11,42,43). On the other hand, an overoxidized surface may be responsible for the further oxidative degradation of the desired products. When olefins are used as the reactant, the interaction of the electronrich organic substrate with the vanadium ions leads to a reduction of the latter, and hence to a lower availability of O-insertion sites. The average oxidation state of vanadium under reaction conditions is determined by the nature of the V-P-O phases which have formed during the catalyst activation and ageing (i.e., during the thermal treatment), and by the operative reaction conditions. These properties, however, are not sufficient to lead to the high selectivity obtained in the formation of maleic anhydride from n-butane, and of maleic and phthalic anhydrides from n-

24 pentane. The multifunctionality must be accompanied by other properties, which allow the multi-step transformation of the paraffin to be carried out i) without permitting desorption of any intermediate to occur; and ii) through the proper sequence of dehydrogenation and oxygen insertion reactions. One main characteristics of n-butane oxidation is the substantial absence of by-products of partial oxidation other than maleic anhydride (apart from a low amount of phthalic anhydride). This means that once the alkane has been adsorbed and transformed to the first intermediate species, the latter has to be quickly transformed up to the final stable product. If this requirement is not met, the adsorbed olefinic-like intermediate may desorb. This leads to a lower selectivity to the final desired product, because the olefin may be readsorbed on nonspecific oxidizing sites yielding other undesired products (aldehydes or acids), which can also be precursors for the formation of carbon oxides. Therefore, a rapid transformation of the adsorbed intermediates to oxidized products is necessary in order to obtain high selectivity to the desired product. In order to guarantee this selective pathway, the catalyst surface must provide the required arrangement of specific oxidizing sites: the different functional properties must be arranged so as to provide an ensemble of sites (or, alternatively, sites with multifunctional properties) able to allow the reaction pathway from alkane adsorption and activation up to its transformation to the final product to be completed. On the contrary, when vanadyl pyrophosphate is used to catalyze the transformation of butenes and pentenes, the catalyst deactivates easily and the performance is very poor (11). This occurs because the acidity of the catalyst is responsible for side reactions when olefins are directly used as the feedstock. In this case, in fact, due to its high nucleophilicity, the olefin may easily interact with different types of sites, other than those able to transform it directly to the final anhydrides. Therefore, the surface acidity of the vanadyl pyrophosphate, which seems to be a necessary property for the transformation of n-butane, is a negative feature when the corresponding olefin is the reactant. For the same reason, the synthesis of acrylic acid from propylene must be carried out in two separate reactors, one for the oxidation of propylene to acrolein and one for the oxidation of the aldehyde to acrylic acid. This is due to the fact that the requirements needed for the two steps make the two reactions incompatible. Acidity is needed in the second step, to favour the desorption of acrylic acid and save it from unselective consecutive reaction, while on the other hand, acidity is detrimental for the first reaction, because it favours the transformation of propylene to undesired products. Therefore, the development of a process for the one-step transformation of propane to acrylic acid will be possible when a catalyst is developed which possesses active sites able to perform quickly the complete transformation of adsorbed propane to the acrylic acid, the latter being the only product which finally desorbs into the gas phase. Accordingly, best performances in the oxidation of propane to acrylic acid have been reported to be obtained on heteropolyoxomolybdates (26), which are known to couple tuneable acid and redox properties. In this case, acid properties may facilitate the desorption of acrylic acid. A further requirement for the high selectivity to maleic anhydride from n-butane is the need for a correct sequence of oxidehydrogenation and oxygen insertion reactions. In the oxidation of n-butane the olefinic-like intermediate must be quickly oxidehydrogenated to an adsorbed dienic-like compound in order to favour the selective pathway towards maleic anhydride. In fact, this reaction may occur concurrently with the oxidation of allylic carbon atoms, with formation of aldehydes and acids which can also be precursors of carbon oxides. Thus, the selectivity to maleic anhydride depends on the relative rates of hydrogen abstraction and oxygen insertion. This property can be considered as typical of the vanadyl pyrophosphate; for instance, in the case of the V/Mo/O system the rate of oxygen insertion

25 (represented by the reaction of benzene oxidation to maleic anhydride) is very high in comparison to that of n-butane oxidation. In fact, the V/Mo/O system is selective in benzene oxidation but not in n-butane oxidation. The combination of acidic and oxidizing properties of vanadyl pyrophosphate makes several different transformations possible over paraffins, as illustrated by the scheme in Figure 1 for the reactions which may occur on n-pentane. The relative contribution of the different pathways (i.e., cyclization of intermediate olefin or dienic compound vs. Oinsertion, or dimerization vs. cyclization) is a function of the nature of the reactant and of the availability of surface oxidizing centres or of sites which can favour the dimerization or condensation reactions.

~

I ....

0 aollxYlaClion / ....

COx

,~ allyllc ----b .o.,,.,,o.

/J'~'

X

cox

/

" ~ 1....7 6 0

allylic

\ electrophilic

I

allylic I

~____rtzatlon

.

l,,o.,..,,on o,,.*

~

COx * / z

Figure 1. Possible reactions which may occur on paraffins catalyzed by vanadyl pyrophosphate, as exemplified for n-pentane transformation.

The oxidation of alkanes on heteropolycompounds Heteropolyacids and their salts have been studied as catalysts for oxidation both in the liquid and in the gas phase of several organic saturated and unsaturated substrates (44-50). The main features of these systems, which make them suitable for application as heterogeneous catalysts, can be summarized as follows (51-53): 1) Relatively high thermal structural stability; the Keggin anion in heteropolyacids begins to decompose at temperatures close to 250-300~ Salification leads to a remarkable improvement in the stability, allowing operations up to 400-450~ to be carried out. Salts can be prepared by ion-exchange in aqueous solutions of proton form of the heteropolyacid, or by direct precipitation of the insoluble salt. In some cases, such as the cesium salts of 12molybdophosphoric or tungstophosphoric acids, the structure is stable up to 550~ 2) High intrinsic acidity which arises form the presence of a highly delocalized anion charge and of mobile protons; in both aqueous and organic media the acidity is even higher than that of some mineral acids. The Broensted acidity can be controlled by partial neutralization of the protons.

26 3) Easy reducibility and reoxidizability by means of molecular oxygen. This allows them to be used as catalysts for liquid-phase or gas-phase multi-electron oxidations. The redox properties of these materials can be affected properly by modifying the anionic composition (for instance, by substitution of some Mo 6+ cations by V5+), or the cationic composition. Several different cations can be introduced in the structure, i.e. alkali and alkaline earth metals, ammonium, divalent and trivalent metals such as Cu 2+, Co 2+, VO 2+. Much interest is related to the use of HPC's as catalysts for the oxidative functionalization of light paraffins, since the multifunctional properties of these systems may be of utility for the activation of saturated organic substrates. On the other hand, the high temperatures which are usually necessary to activate the paraffins may be deleterious for the structural stability of HPC's, since the destruction of the primary structure leads to a loss of the unique properties of the compound and to a decrease in catalytic performance. Therefore, it is necessary to use stable salts, which, on the other hand, are less reactive. This seems to leave little room for the design of a catalyst suitable for the activation of paraffins in the gas phase. However, different possibilities exist to resolve these problems: 1) Operation in the liquid phase at high pressure and moderate temperature; this is the case of the liquid-phase hydroxylation of alkanes (54). 2) Use supported HPCs, in order to obtain better spreading of the active phases (i.e., increase the specific surface area of the active compound); a problem may be the strong chemical interaction that develops between the support and the HPC, which can lead to the destruction of the compound itself. Less reactive supports are therefore needed, such as silica. 3) Improve at the same time the structural stability of the HPC and its oxidation potential, by using stable salts and by adding transition metal ions which are known to enhance the oxidation potential of the oxometal, such as low-valence early transition metal ions in the cationic position of the HPC, or vanadium as an additional oxometal. 4) In gas phase oxidations usually water is also added to the stream. This is supposed to guarantee stable catalytic performance and higher activity, likely because the presence of water can favour the surface reconstruction of the heteropolyacid even under conditions at which it would usually decompose. In addition, water may favour the desorption of the products, saving them from unselective consecutive combustion. In recent years heteropolycompounds have been studied for the oxidation of propane to acrylic acid and of isobutane to methacrylic acid. Rohm & Haas Company was the first in 1981 to claim the one-step oxidation of isobutane to methacrolein and methacrylic acid (55). Even though no reference is made to heteropolycompounds, the claimed catalyst compositions correspond to Keggin-type structures. In the patents later issued by Sumitomo (56,57) an important role was claimed to be played by vanadium (in an anionic position), by cesium (in a cationic position), as well as by an excess of phosphorus with respect to the stoichiometric composition. These catalysts gave selectivities to methacrylic acid plus methacrolein close to 70 %, with isobutane conversions in the 10 to 13 % range. Besides carbon oxides, acetic acid was the main by-product. Asahi also has patented Keggin-type heteropolycompounds containing vanadium, copper and alkali metals (27). The necessity for the use of salts instead of acids was also pointed out, in order to increase the catalyst stability. The use of a catalyst suitable for fluid-bed operation was also claimed, so as to allow continuous transport of catalyst from the reaction to the regeneration vessel, and vice versa. Also, alternate feeding of isobutane and of oxygen into the catalytic bed allowed higher selectivities to be obtained. Fundamental features of the claimed heteropolycompounds were i) a partial degree of reduction, and ii) the cubic crystalline structure. This was achieved by calcining the ammonium salt of H3PMo12040 in a nitrogen atmosphere. The heteropolycompounds characterized by the cubic structure, and

27 by a partial degree of reduction (also achieved by in situ treatment with isobutene at 450~ led to superior activity and selectivity to methacrylic acid. The importance of a partial molybdenum reduction has been also recently pointed out by Mizuno et al. (26). The high variability in the compositions of heteropolycompounds makes them potentially useful for the oxidation of different organic substrates, or for addressing the transformation of one reactant into different kinds of products. One example is given by the comparison of the reactivity of Keggin-type molybdenum-based compounds against tungsten-based compounds (58-61). Figures 2 and 3 show the distribution of the products obtained in the oxidation of isobutane at the following conditions: 26% isobutane, 12 % oxygen, 12% water, balance helium, with (NH4)3PMo12040 and with (NH4)3PW12040 catalysts, respectively. The former compound yields methacrolein and methacrylic acid with an overall selectivity around 60%; the main by-products are carbon oxides and acetic acid, with minor amounts of maleic anhydride and of acrylic acid. On the contrary, the W-containing compound exclusively yields isobutene as the product of partial oxidation. The olefin is obtained with a selectivity which is close to 60%. 50

60 methacrylic acid

40

o~ =" 30 0

oxyge

~,

co

L

r 20 -

0

~

o

methacrolein - 20

v

acetic acid

10-

0 320

.>

carbon oxides

>

~

~ 340

n

r

e

360

G)

u~

0 380

400

temperature, ~

Figure 2. Oxidation of isobutane over (NH4)3PMo12040; reaction conditions" residence time 3.6 s, feed composition 26% isobutane, 12% water, 13% oxygen, rest helium.

The very different catalytic behavior originates from the different reducibility of the Mo 6+containing HPC with respect to the W o+ -containing compounds. The Keggin anion in the former compound easily undergoes 2-electron reduction, furnishing nucleophilic 0 2- species which can be incorporated into the activated organic substrate. The PW120403- Keggin anion is much less reducible under the same conditions, and can only perform hydrogen abstraction with water formation. This difference arises for the greater electronegativity of the Mo 6+ and the lower Mo-O bond strength. Only in the case of ethane oxidation, are the Mo-containing heteropolycompounds active and selective catalysts in the formation of ethylene, since this molecule does not possess allylic carbon atoms for O-insertion (62-64). The two different mechanisms which have been proposed for the oxidation of isobutane over i) Mo-containing Keggin-type compounds and ii) W-containing Keggin-type and Wells-

28

Dawson-type heteropolycompounds are illustrated in Figure 4 (24,58). In both cases the first step consists of the formation of an alkoxy species, which is the first reaction intermediate. 4

80 carbon oxides

- 70

_

- 60 isobutane conversion

- 50

o~

0

m 2 -

- 4 0 .~-

G) > tO 0

0

-30

;obutene

u)

- 20

1-

"10 m e t h a c r y l i c acid + m e t h a c r o l e i n + a c e t i c acid

0 330

'~

="

340

"

350

-

360

temperature,

0

370

380

~

Figure 3. Oxidation of isobutane over (NH4)3PW 12040; conditions as in Figure 2. CH3 H3C. /10H3

o.'',.

OH3 H3C~.CI/0H3

," - ~

.

<~..o\

o,.,..

~o.~

GIH3 2Hc~C=CH 2 "---* ~ / H H

o,o.

o,

\ Mo I / O \ 1 Mo / O\ M Io ? H3 , '' H-~c-IC =CH2

/o

--,

~

CH3 H3C-~-CI----OH3 + .H

\

.....

0-%

CH3 2HC_ ~._r oH2 -H./" " "

o.,.o

o CH3 H I - - . ~ c_~C=CH 2 H ---.~ ~ H.~""I

o

Mo./i~

-

O O Vl~/fO...~v,

o

CH3 1 H~c~-C=CH 2

-

I /O. I ~Mo Mo methacrylic acid

s,~, methacrolein T"~ j -H~.c.~C=CH~

:" o/.',..;--,.I ,,--'

' '<.1 / o \ Mo

H3C~_> ~ H3 / - - C H

"'H

.~,Mo ..~.I 1 Mo.8~ Mo

o/

? H3 c f C =CH2 ~ 0~1-1~ /

"/ H~",,-, ~ oI ' , o .~ ,.,I 'i' k_.9. \ M ' 0 " " Mo O\Mo'O"~o

"

'~

Mo

H3C /H H3C .H H3C~C..~C-~.._ H H3C.._NCI~_H C~ H H / H H O *'--" O I O ,V I O ' v, v ~vfO..... ~v W" t ~ W .

'~+1/2 O2 O vI II - O ~ IV W f "~W H3C H /C-- C\~ 4 o i .. + ,, vwIt ~ O~ w v H,C H H20

/ H H3C H,C>C I ( - H . o/ H 6"

,vdV~O...~v

Figure 4. Proposed mechanisms for isobutane oxidation on Mo- and W-containing heteropolycompounds

29 However, with the former catalyst the mechanism is essentially an ionic one, with evolution to a dioxyalkylidene species which can yield either methacrolein, or a carboxylate species, precursor of methacrylic acid. No isobutene is detected among the reaction products since the dioxyalkylidene species is strongly bound to the surface. The overall process involves the transfer of 8 electrons from the organic substrate to the catalyst per molecule of methacrylic acid produced. In the case of W-containing heteropolycompounds the mechanism is essentially a radical one (analogous to that proposed for the radical-like chemistry exhbited by W-containing heteropolycompounds in liquid phase oxidations (46)). In the absence of centres able to insert 0 2- species, the radical alkoxy intermediate converts to isobutene. The process only involves 2 electrons. o~ 12 "~

~

0 ffJ

8

L.

r> cO o

6 4'

e-

=

,I1 0 9"-

2 (

0

0.25

I

I

I

I

I

0.5

0.75

1

1.25

1.5

number of iron atom per Keggin unit

Figure 5. Catalytic performance of KI (NH4)2PMol2040/FexO1.5x in the oxidation of isobutane.

A role can also be played by the cations. Catalysts were prepared which contained iron ions, with the composition, KI(NH4)2PMol2040/FexO1.5x (23-25). Iron was found to substitute for ammonium in cationic positions, possibly in the form of partially hydrolyzed cations, i.e. [Fe(OH)3_x] x+. Figure 5 shows that increasing amounts of iron led to a progressive increase in the isobutane conversion and yield to methacrolein plus methacrylic acid, while the selectivity was negatively affected by the dopant. It is known that the nature of the cation may affect the redox properties of the molybdenum in the primary structure, and this might explain the progressive increase in catalytic activity. However, the addition of iron also was found to increase the Broensted and Lewis acidity of the compound; this acidity might be involved in the rate-determining step of the reaction, facilitating the polarization of the C-H bond.

The ammoxidation of propane over rutile-based mixed oxides The synthesis of acrylonitrile from propane, as an alternative route to the industrial process which employs the olefin as the raw material, is carried out on catalysts which are based on vanadium and antimony mixed oxides. The catalyst contains a large excess of antimony with respect to the stoichiometric requirement for the formation of VSbO 4, and the

30 excess of antimony oxide is claimed to remarkably promote the selectivity to acrylonitrile (15,31-33). This catalytic system, as well as systems based on MoN/Te/Nb mixed oxides which have been developed by Mitsubishi (65), also represent an example of catalyst characterized by multifunctional properties. The rutile structure is the matrix to host vanadium ions as solid solutions, while the antimony oxide is present as a dispersed microcrystalline oxide. Vanadium is the component which is more active in paraffin conversion, while the high selectivity to the desired product is due to the presence of dispersed, separate phase, antimony oxide. In this case the design of the catalytic system is aimed to accomplish the following transformations: C3H 8 + 1/2 0 2 --~ C3H 6 + H20 C3H 6 + NH 3 + 3/2 0 2 ~ C3H3N + 3 H20 Sohio issued several patents claiming catalysts based on vanadium, antimony and some promoters which are able to ammoxidize propane with a completely heterogeneous mechanism (66). These catalysts can be considered intrinsically multifunctional since both dehydrogenation and nitrogen insertion functions are present (67,68). The main problem with this type of catalyst is the low rate of the subsequent ammoxidation of intermediate propylene. Indeed, propylene is always present as a by-product. The mechanism for acrylonitrile formation via intermediate propylene is demonstrated by the results given in Figure 6, which compares the oxidation and the ammoxidation of propane over a V/Sb/O mixed oxide catalyst (69). The conversion of propane and the selectivity to the main products are reported (acrolein was formed only in traces).

~6

t~ ~D m .m

s

0

tO2 0

product Figure 6. Ammoxidation of propane over V/Sb/O catalyst. Reaction conditions" temperature 430~ residence time 2s, propane 25 mol.%, ammonia 10 mol.%, oxygen 20 mol.%.

The rate of conversion of propane is practically the same in the presence and in the absence of ammonia. The oxidation yields propylene and carbon oxides, which are the prevailing products. However, when ammonia is added to the feedstock, the yield to propylene remains unchanged, while the yield to carbon oxides is remarkably decreased in favour of the formation of acrylonitrile. This suggests that in the absence of ammonia the propylene is oxidized to a compound or to an intermediate which under these conditions is burnt to carbon oxides. The addition of ammonia allows this intermediate compound (which might be acrolein or an allyl radical species) to be quickly transformed to the stable

31 acrylonitrile, thus saving it from the unselective subsequent combustion. In any case, the presence of residual propylene suggests that the catalyst is not very efficient in its transformation, likely because of the absence of O or N-insertion sites in proximity to the centres responsible for the formation of the allylic intermediate. A more efficient catalyst will be one where the different functions are organized into arrays of active centres, able to quickly transform the propylene and avoid its desorption, in the same way as it occurs in the oxidation of n-butane over vanadyl pyrophosphate (which does not yield olefins or diolefins at all), and for the oxidation of isobutane to methacrylic acid over Mo-heteropolycompounds (which do not yield isobutene). An important concept, first developed by Callahan and Grasselli (70) is the "siteisolation" theory, which requires that the active oxygen species is present in isolated regions on the catalyst surface, in order to obtain high selectivity to the desired product. In this case, the presence of diluted V sites provides centres for paraffin activation, while bulk vanadium oxide is responsible for side undesired combustion reactions (32).

C o n c l u s i o n s

The main topics which have been emphasized in this review are the following: - the role of vanadium, of molybdenum and of tungsten in some heterogeneous catalysts which are active and selective in the mild oxidation of light paraffins; - the role of the stability of the products; - the role of the multifunctionality of the catalysts; - the role of non-desorption of reaction intermediates; - the role of the relative rates of the reactions of dehydrogenation and of O-insertion; - the role of acidity and basicity of the catalysts; - the role of isolation of active sites. Three catalytic systems active and selective in the (atom)oxidation of saturated organic substrates have been discussed, and the properties which lead to their superior catalytic performance have been examined in relationship to the mechanism of paraffin activation and transformation to the desired products. Common features for these systems are: 1) The necessity of multifunctional properties, arising either i) from the presence of different components, having a specific role in the reaction mechanism but intimately interacting with each other in a complex array of surface centres, or ii) from the presence of monophasic structures characterized by different functions. 2) A weak interaction of the desired product with the catalyst surface, which favours its desorption into the gas phase. The stability of the product, moreover, must be high enough to allow operation to be carried out under conditions which are necessary for the paraffin activation. Properties which are specific for the compounds examined, and which make them rather unique for some peculiar reactions of paraffin oxidation are: 1) The availability of surface arrays of active centres able to cooperate in such a way as to avoid the desorption of olefin intermediates, and to allow the rapid transformation of the latter to the final oxygenated compounds. This is a property of the vanadyl pyrophosphate and of heteropolycompounds; for these compounds the "intrinsic multifunctionality" arises from the molecular-type organization of the structure into well-defined moieties, each one characterized by specific properties. 2) The specificity in the synthesis of anhydrides from n-butane and n-pentane, typical of the vanadyl pyrophosphate, which likely arises from its property of electrophilic O-insertion onto

32 dienic-like intermediates. The heterocyclic compound is then oxidized to the very stable anhydrides. On the contrary, low selectivity to unsaturated acids are obtained in the oxidation of propane and of isobutane; this may be due to the fact that this catalyst does not allow a rapid desorption of these intrinsically unstable products, and thus favours the occurrence of consecutive combustion reactions. One possible hypothesis is the formation of stable esters by reaction between the acids and the surface P-OH groups. 3) The specificity of Mo-heteropolycompounds in the synthesis of unsaturated acids starting from paraffins which can not yield heterocyclic compounds (i.e. propane and isobutane). This may arise from the very strong acidity which is typical of these compounds (even when they are not in a protonic form) which favours the desorption of the organic acids, sparing them from consecutive reactions. These compounds also yield maleic anhydride from nbutane and n-pentane (but with remarkably lower selectivity than the vanadyl pyrophosphate), but they lack the surface properties which are necessary for the formation of phthalic anhydride from n-pentane. 4) The unique possibility of tuning the redox properties (and the acidic properties as well) of heteropolycompounds by chancing the composition of the anion, or by modifying the cationic composition. Changes in composition have dramatic effects in the nature of the products obtained in paraffin oxidation, i.e. oxygenated compounds vs. olefins. 5) The rutile structure is characterized by a very elastic lattice, able to guest foreign metal ions, as well as domains or dispersed amorphous or microcrystalline phases. It is thus possible to build a multifunctionality by properly choosing the nature and amount of the various components which can be added to the rutile matrix. The most studied application for rutile-based materials is the transformation of propane to acrylonitrile, because this structure makes it possible to couple oxidehydrogenating properties with O- or N-insertion characteristics. This is an example of how the use of nanosized technologies for catalyst preparation may substantially enhance the interaction between different phases, favouring phenomena at the interface and promoting the system multifunctionality (71).

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

Selective oxidation compounds

of

hydrocarbons

35

catalyzed

by

heteropoly

Makoto Misono, Noritaka Mizuno, Kei Inumaru, Gaku Koyano, and Xin-Hong Lu Department of Applied Chemistry, Graduate School of Engineering, The University of Tokyo, Bunkyo-ku, Tokyo 113, Japan

Selected characteristic features of heteropoly catalysts for the selective oxidation of hydrocarbons are described based on recent studies from our laboratory as well as from other groups. 1. C O R R E L A T I O N PERFORMANCE

BETWEEN

REDOX

PROPERTIES

AND

CATALYTIC

Fundamental correlations between redox properties and catalytic activity have successfully been established for the hydrogen form and alkali salts of 12molybdophosphoric acid [1]. Provided that the contributions of surface- and bulk-type catalysis are properly taken into account, good monotonic relationships are obtained between the catalytic activity for oxidation and the reducibility (or the oxidizing power) of the catalyst. The rate of oxidation of aldehydes, a surface-type reaction, correlates linearly with the surface reducibility of the catalyst, and the rate of oxidative dehydrogenation of cyclohexene, a bulk-type reaction, with the bulk reducibility [2].

~ a

1

~

2

~ ~0

m

t'~

3 0

N

~m ~o

1 2

L

0

.

|

2O

I

I

I

I

!

4O 6O 80 N MAA Yield/% Figure 1. Effect of V and Cs contents on the yield of methacrylic acid (MAA) at 350~ over H3+xPMo1~.-xVxO40and Cs2.75H0.2~+xPMo12_xVxO40catalysts.

36 The quantitative agreement between the rates of catalytic oxidation observed experimentally and those predicted from the reduction and oxidation rates of the catalysts measured independently demonstrated that the catalytic oxidation proceeded by redox cycles of the catalysts, that is, the redox mechanism or Marsvan Krevelen mechanism [3]. However, attempts to find similar relationships for mixed-metal heteropoly compounds such as molybdovanadophosphates have not been successful. This has been due to the low thermal stability of these compounds. For example, PMollVO4o and PMoloV204o decomposed to PM012040 and VOx above 200~ [4]. We attempted to stabilize the heteropolyanions by forming their cesium salts. Although the possibility of slight decomposition could not be excluded, high yields were obtained for the conversion of isobutyric acid to methacrylic acid (MAA) as shown in Fig. 1 [5]. A crystalline vanadium phosphorous oxide may be regarded in a broad sense to be a heteropoly compound. By applying e x s i t u and in s i t u spectroscopies (Raman spectroscopy, infrared spectroscopy (IR), X-ray photoelectron spectroscopy (XPS), extended X-ray absorption fine structure (EXAFS), X-ray diffraction (XRD), electron diffraction (ED)), we found that the surface of (VO)2P207 was reversibly oxidized to the X1 (5) phase of VOPO4 under reaction conditions in the oxidation of n-butane to maleic anhydride [6]. For example, the in s i t u Raman spectra measured at steady state flow reaction conditions showed that the surface changed reversibly between (VO)2P207 and X1 phase depending on the butane/oxygen ratio in the feed [7]. Correspondingly, the catalytic selectivity varied reversibly (Table 1). Table 1 Changes in the selectivity to maleic anhydride over (VO)2P207 with the partial pressure of n-butane Partial pressure Conversion a/% Selectivity/% Phase b of C4Hlo/% 1.5 c 53.8 52.4 P 0.75 c 52.0 48.8 P 0.25 c 56.8 23.4 P (+ X) 0.25 d 56.0 22.8 P +X 0.75 d 51.3 44.1 P 1.5 d 45.5 61.7 P The flow rate of the feed was adjusted at each step to obtain approximately the same conversion, b Determined by in situ Raman. P: (VO)2P207 and X: XI. r Partial pressure of butane was decreased stepwise from 1.5% (17%, partial pressure of 02 in parentheses) to 0.75 (18.5)% and then 0.25 (19.5)% after N2 treatment of the catalyst at 500~ The balance was N2. ~ Partial pressure of butane (02 content) was increased stepwise from 0.25% (19.5%, partial pressure of 02) to 0.75 (18.5)% and 1.5 (17)% after treatment in air at 460~ ,

,

,

,

37 2. S E L E C T I V E OXIDATION OF ALKANES There have been several a t t e m p t s to obtain oxygenated products from lower alkanes (C2 - C5) by using heteropoly catalysts. It has been reported t h a t the hydrogen form of H3PMo12040 catalyzes the oxidation of lower alkanes to aldehydes and carboxylic acids [8] and t h a t the substitution of V 5§ for Mo 6+ modified the catalytic activity and selectivity [1, 9 - 12]. By optimizing the q u a n t i t y and type of constituent elements of heteropolyanions and counter cations, fairly good yields were obtained for the oxidation of isobutane [13 - 17]. Recently, it was found t h a t acidic cesium salts of Keggin-type heteropolymolybdates can efficiently catalyze the oxidation of isobutane to methacrylic acid with molecular oxygen. The optimal contents of Cs and V were 2.5 and 1, respectively, and the addition of Ni enhanced the catalytic activity even further will be discussed below [13- 16]. The results for CsxH3-xPMo1204o catalysts are shown in Table 2 [13]. The highest conversion was observed around x = 2.5 - 2.85. The main products were methacrylic acid (MAA), methacrolein (MAL) and acetic acid (AcOH). The substitution of Cs for H in H3PMo12040 resulted in a great e n h a n c e m e n t of the MAA production and the yield reached a m a x i m u m at x = 2.5. The sum of the yields of MAA and MAL on Cs2.sHo.sPMo1204o reached 5.1%. The catalytic properties of Cs2.sHo.~PMo~2040 changed by the addition of transition metal ions [16]. The addition of Ni, Mn, or Fe increased the yields of MAA and MAL. In the case of Ni, the yields of MAA and MAL reached 6.5 and 1.5%, respectively. In contrast, Co, Cu, Hg, Pt, and Pd decreased the yields. The results for Cs2.sNioosHo.34+xPMo12-xVxO4o catalysts are shown in Table 3 [16]. The conversions were 10 - 15%. The highest selectivity to MAA was also observed at x = 1. It follows t h a t the substitution of V 5+ for Mo 6+ in Cs2.sNio.0sHo.34PMo~204o resulted in the e n h a n c e m e n t of MAA production and the yield reached a m a x i m u m at x = 1. Table 2 Oxidation of isobutane over CsxH3-xPMo12040 at 340~ a x Surface Conv. Rate Selectivity 5/% area /% /10 .5 mol /m 2 g-1 min -1 m 2 MAA MAL AcOH CO

Sum of

yields of MAA+ C02 MAL/% 0 1.1 7 1.34 4 18 8 44 26 1.5 1 2.1 6 0.60 23 17 10 32 18 2.4 2 5.9 11 0.39 34 10 7 29 21 4.8 2.5 9.5 16 0.36 24 7 7 41 21 5.1 2.85 46.0 17 0.08 5 10 5 44 37 2.4 3c 46.0 8 0.04 0 10 6 32 35 0.8 a Isobutane, 17 vol%; 02, 33 vol%; N2, balance; catalyst, 1.0 g; total flow rate, ca. 30 cm 3 min-1, b Calculated on C4 (isobutane)-basis. c The selectivity to acetone was 17%.

38 Table 3 Oxidation of isobutane catalyzed by Cs2.sNio.osHo.sa+~PMol2-xV~O4o at 320~ a X Conv./% Selectivity 5/% MAA MAL AcOH CO CO2 0 10 27 12 5 30 26 1 15 36 9 6 25 24 2 13 28 8 6 25 33 3 12 10 8 9 35 38 a,b Experimental conditions. See Table 2. Thus, the Keggin-type heteropolymolybdates such as Cse.5Ni0.08Hl.s4PMollVO40 fairly selectively catalyze the oxidation of isobutane into methacrolein and methacrylic acid with molecular oxygen. At 340~ the yield of methacrylic acid reached 9.0%. The 9.0% yield of A A was greater than the highest value of 6.2% reported in the patent literature at similar steady-state conditions [10]. Figure 2 shows a good correlation between the rates of oxidation of isobutane and non-catalytic reduction of catalysts by CO. The correlation noted in Fig. 2 indicates that the catalytic activity is controlled by the oxidizing ability of catalysts. It has been suggested for the case of CsxHs-~PMo1204o catalysts that the factors controlling the catalytic activity are the o 2 oxidizing ability and the O protonic acidity of catalysts [16]. r 9

C s2.5Mn+o.08H 1.5-0.08nPMo I 1VO 40

(M = Ni z+, Fe 3+) also catalyzed the oxidation of propane and ethane [18- 20]. Here, the rate and role of vanadium is of concern [21]. It is interesting that the reduced heteropoly compounds showed higher selectivity to methacrylic acid for the oxidation of isobutane [10, 13, 15, 22, 23]. Ueda et al. applied reduced 12-molybdophosphoric acid to the oxidation of propane and obtained 50% selectivity to acrylic acid and acrolein at 12% conversion [22, 23].

O

O

H

~

Cs2.85~

o~ r

0

S~Cs2 / ~ - - - Cs3

1 2 3 Rate of reduction by CO /10 .6 mol min-1 m-2 Figure 2 Correlation between the rates of catalytic oxidation of isobutane and those of non-catalytic reduction of catalysts by CO. Csx and H represent CsxHs-~PMoy204o and HsPMo12040, respectively.

39 3. SELECTIVE HYDROXYLATION OF BENZENE Various oxidants such as dioxygen, hydrogen peroxide, and alkyl hydroperoxides have been applied for the oxidation of hydrocarbons in the homogeneous liquid phase catalyzed by heteropoly catalysts [1, 24]. Below are presented results on the selective hydroxylation of benzene to phenol with H202 catalyzed by Keggin-type heteropolyanions. It is known t h a t Fenton or related reagents also catalyze this reaction [25]. A research group that includes one of the present authors has previously reported that the reaction proceeded selectively by using H202 and vanadium-substituted heteropolymolybdates or tungstates [26]. The present study is an extension of this earlier study. Recently the same reaction was attempted by using wellcharacterized K salts of vanadium-substituted heteropolytungstates [27]. The selectivity based on benzene was high, but the yield based on H202 was not given. Heteropoly compounds obtained commercially, H3+~PMo12-~V~O40 (x = 0 - 4), were purified by extraction with ether and subsequent recrystrallization. They are abbreviated as PMol2-xVx hereafter. Their IR spectra agreed with those reported in the literature [27]. 51V-NMR spectrum of PMoloV2 in aqueous solution showed that PMoloV2 was a mixture of three to four positional isomers of PMol0V2 and contained PMol~V at less t h a n 30% level. For comparison, NaVOa, V203, V204, and V205 were used. In the case of V2Ox, sulfuric acid was added in order to dissolve the catalyst. This addition resulted in improved yields. The reaction was usually carried out at 20 - 70~ in a four neck flask, by adding dropwise 10 ml of 0.08 M H202 aqueous solution (0.8 mmol) into a mixture of benzene 10 ml and water 15 ml. Catalyst (0.05 - 0.3 mmol) was dissolved in the water phase before the reaction. Hydroxylation of benzene took place in the water phase and a majority of phenol formed was transferred to the benzene phase. The concentration of water and benzene phases were analyzed periodically by liquid chromatography with o-cresol as a standard. The evolution of oxygen gas by the decomposition of H202 was measured volumetrically. The yield of phenol on the basis of H202 consumed tended to increase in parallel with the catalytic activity of each catalyst for the H202 decomposition in the absence of benzene. The sudden introduction of H202 into the reaction system caused the evolution of a significant amount of oxygen. After a short induction period the yield of phenol increased rapidly with time. When the concentration of H202 in the reaction system was kept low by adding dropwise very slowly a diluted H202 solution, the yield of phenol increased remarkably, the unproductive decomposition of H202 being suppressed. The selectivity to phenol was almost 100 % on the basis of benzene and reached above 90 % on the basis of H202. Typical results thus obtained at 65~ are summarized in Table 4. The yields are in the order of PMol0V2 > PMo9V3 > PMosV4 > PMo~V >> PMol2. NaVO3 and V2Ox showed modest performance. The turnover based on the catalyst was about 3 in the case of PMoloV2. It was confirmed that, upon the addition of H202

40 Table 4 Yields of phenol from the oxidation of benzene by hydrogen peroxide and vanadium compounds Yield of Selectivity of Catalyst Selectivity of Catalyst Yield of phenol/% phenol/% phenol/% phenol/% (H202 (benzene (benzene (H202 basis) basis) basis) basis) 65~ 65~ 45~ VO(C5H702)2 46.6 PMo1204o 0 0 92.1 NaVO3 23.8 PMo11VO4o 9.0 0 100 56.0 V205 (I-12804) 52.5 PMoloV204o 92.6 69.2 100 81.6 V204 (H2SO4) 21.8 PMooV304o 90.1 90.7 100 V203 (a2so4) 13.5 PMosV404o 68.1 73.6 100 Catalyst, 0.3 mmol; H20, 15 ml; C6H6, 10 ml. 10 ml of 0.08 M H202 was added dropwise very slowly. Reaction time: 1.5 h. after the reaction, the oxidation proceeded again at a similar rate. Therefore, there was little deactivation of catalyst. The optimum pH range for the phenol yield was 2 to 3 for PMoloV2 and PMo9V3, and 1 to 2 for PMollV and PMosV4. According to the UV-vis spectra, PMoIoV2 and PMogV3 were stable in the pH ranges of 2 - 3 and 1.5 - 3.5, respectively. The temperature dependencies are shown in Fig. 3. The optimum reaction temperature and time appear to depend on the polyanion species. 100 The IR spectra of the reaction solutions showed that 80the structures of the Keggin polyanions remained unchanged 60 r during the reaction. When the concentration of H202 was 40 O increased, however, new b a n d s CD .t-.4 appeared in the 500 - 600 cm -1 20 region possibly due to peroxo I I species. It seems that active 0 "AI " 10 20 30 40 50 60 70 80 peroxo species are formed by the Temperature/~ reaction of vanadiumFigure 3. Temperature dependencies of the substituted polyanions and activity for the oxidation of benzene with H2Oe or that other active oxygen hydrogen peroxide for 1.5h reaction time. species are derived from the Catalyst, 0.3 retool; H20, 15 ml; C6H6, 10 ml. peroxo species. Either or both 10 ml of 0.08 M H202 was added dropwise species are probably active for very slowly. the hydroxylation of benzene. A; H4PMollVO4o, A; H5PMoloV2040, The species tend to deactivate O; H6PMogV304o, 0; H:PMosV404o, by reaction between them (e. g., ["l; H3PM012040, X; V205 (H2804) dimerization), as indicated by

41 the fact that the unproductive decomposition of H202 to oxygen and water became dominant when their concentration was high. The induction period observed when H202 was added suddenly probably corresponds to the period for the formation of the active species. Thus, by choosing appropriate vanadium-substituted heteropolymolybdates and keeping the concentration of H202 low, efficient hydroxylation of benzene to phenol was achieved. The highest yield based on H202 was 93%, where the selectivity with respect to benzene consumed was 100%. 4. CONCLUSION Characteristic features of vanadium containing heteropoly catalysts for the selective oxidation of hydrocarbons have been described. MAA yield from isobutyric acid was successfully enhanced by the stabilization of the vanadiumsubstituted heteropolyanions by forming cesium salts. As for lower alkane oxidation by using vanadium containing heteropoly catalysts, it was found that the surface of (VO)2P207 was reversibly oxidized to the X1 (5) phase under the reaction conditions of n-butane oxidation. The catalytic properties of cesium salts of 12-heteropolyacids were controlled by the substitution with vanadium, the Cs salt formation, and the addition of transition metal ions. By this way, the yield of MAA from isobutane reached 9.0%. Furthermore, vanadium-substituted 12molybdates in solution showed 93% conversion on H202 basis in hydroxylation of benzene to phenol with 100% selectivity on benzene basis. REFERENCES

1. 2.

T. Okuhara, N. Mizuno, and M. Misono, Advan. Catal., 41 (1996) 113. M. Misono, N. Mizuno, H. Mori, K. Y. Lee, and T. Okuhara, Stud. Surf. Sci. Catal., 67 (1991)87. 3. N. Mizuno, T. Watanabe, and M. Misono, J. Phys. Chem., 94 (1990) 890. 4. E. Cadot, C. Marchal, M. Fournier, A. Teze, and G. Herve, in Polyoxometalates: From Platonic Solids to Anti-retroviral Activity (Eds. M. T. Pope and A. Muller), Kluwer, Dordrecht, 1994, p 315. 5. K.Y. Lee, S. Oishi, H. Igarashi, and M. Misono, Catal. Today, 33 (1997) 183. 6. G. Koyano, F. Yamaguchi, T. Okuhara, and M. Misono, Catal. Lett., 41 (1996) 149. 7. G. Koyano, T. Saito, and M. Misono, Chem. Lett., (1997) in press. 8. H. Krieger and L. S. Kirch, Rhom and Haas Co., Eur. Patent No. 0010902 (1979). 9. G. Centi, M. Burttini, and F. Trifiro, Appl. Catal., 32 (1987) 353; J. B. Moffat, ibid., 146 (1996) 65. 10. H. Imai, T. Yamaguchi, and M. Sugiyama, Asahi Chemical Industry Co., JP 145249 (1988). 11. K. Nagai, Y. Nagaoka, H. Sato, and M. Osu, Sumitomo Chemical Co., JP 106839 (1991).

42 12. G. Centi, J. L. Nieto, C. Iapalucci, K. Bruckman, and E. M. Serwicka, Appl. Catal., 46 (1989) 197; F. Cavani, E. Etienne, M. Favaro, A. Galli, F. Trifiro, and G. Hecquet, Catal. Lett., 32 (1995) 215. 13. N. Mizuno, M. Tateishi, and M. Iwamoto, J. Chem. Soc. Chem. Commun., (1994) 1411. 14. N. Mizuno, M. Tateishi, and M. Iwamoto, Appl. Catal. A: General, 118 (1994) L1. 15. N. Mizuno, W. Han, T. Kudo, and M. Iwamoto, Stud. Surf. Sci. Catal., 101 (1996) 1001. 16. N. Mizuno, M. Tateishi, and M. Iwamoto, J. Catal., 163 (1996) 87. 17. F. Cavani, E. Etinne, M. Favaro, A. Gall, F. Trifiro, and G. Hecquet, Catal. Lett., 32 (1995) 215. 18. N. Mizuno, M. Tateishi, and M. Iwamoto, Appl. Catal. A: General, 128 (1995) L165. 19. N. Mizuno, D. J. Suh, W. Han, T. Kudo, and M. Iwamoto, J. Mol. Catal. A: Chemical, in press. 20. N. Mizuno, W. Han, and T. Kudo, Chem. Lett., (1996) 1121. 21. R. Bayer, C. Marchal, F. X. Liu, A. Teze, and G. Herve, J. Mol. Catal. A: Chemical, 110 (1996) 65. 22. W. Ueda, Y. Suzuki, W. Lee, and S. Imaoka, Stud Surf. Sci. Catal., 101 (1996) 1065. 23. W. Ueda and Y. Suzuki, Chem. Lett., (1995) 541. 24. C. L. Hill and C. M. P-M. Charta, Coord. Chem. Rev., 143 (1995) 407. 25. E. g., S. Tamagaki, M. Sasaki, and W. Takagi, Bull. Chem. Soc. Jpn., 62 (1989) 153; S. Ito, A. Mitarai, K. Hikino, M. Hirama, and K. Sasaki, J. Org. Chem., 57 (1992)6937. 26. R. D. Huang, X. H. Lu, B. J. Zhang, and E. B. Wang, Chinese Chem. Lett., 4 (1993) 319. 27. K. Nomiya, H. Yanagibayashi, C. Nozaki, K. Kondoh, E. Hiramatsu, and Y. Shimizu, J. Mol. Catal., 114 (1996) 181.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

The Future Fundamental

of I n d u s t r i a l Advances

Oxidation

43

Catalysis

Spurred

by

B. Delmon Universit~ catholique de Louvain, Unit~ de Catalyse et Chimie des Mat~riaux Divis~s, Place Croix du Sud, 2/17, B-1348 Louvain-la-Neuve, Belgium.

This short review is a response to the invitation of the Organizers to present the state of Industrial Oxidation Catalysis. This is an attempt to sketch a bird's eye picture about fundamental aspects as well as applied perspectives, of catalysts, catalytic mechanisms, reactors and industrial processes. First, a glimpse at the present situation is given, covering both i n d u s t r i a l developments and f u n d a m e n t a l concepts, with particular emphasis on describing trends. The message is essentially that completely new catalytic materials are appearing, new mechanisms in all their complexity are being elucidated and new kinds of reactors and types of unit operations are being designed, all thanks to an exceptionally close cooperation between specialists in different fields. Catalytic oxidation is breaking new ground in the field of catalysis with the development of new concepts and industrial achievements. In view of the limited size allowed for the paper, only very few references will be cited compared to the richness of literature, and many important aspects will not be mentioned. 1. INTRODUCTION The objective of the present contribution is to shed light on how new industrial applications of oxidation catalysis could develop in the future, with account taken of foreseeable fundamental developments. Many excellent review papers and talks have described the new challenges industry is faced with when catalytic oxidation is considered. Equally excellent review papers and lectures have dealt with the scientific aspects of catalytic oxidation. The aim here is not to summarize these articles. It is said t h a t industrial innovation results either from a "push" or a "pull" action. The scientific advances constitute the "push", and the industrial challenges the "pull". Our wish is to provide insight on the way the push-pull process could operate in the next, let us say, 5 or 10 years, in the field of catalytic oxidation. Some scientists, like Professor Y. Moro-Oka, are of the opinion that the science of catalysis, in general, is driven b y applications (strong "pull"). Industry, looking at the situation from the other side, indirectly suggests the same trend, in particular by saying that concepts are lacking (suggesting that the "push" is weak) [1]. However, given the number of novel catalysts, novel reactors and novel processes in the field of oxidation, it would rather seem that neither "push" nor

44 "pull" is determining, but the combination of both. If this is true, it would be useful to attempt to discern the components of both the push and pull actions. By doing so, we may discover that oxidation catalysis, both industrial and fundamental, is in the process of evolving considerably differently from other fields of catalysis. C o n s i d e r i n g i n d u s t r i a l challenges in oxidation, especially the transformation of light alkanes to useful intermediates, the incredible variety of formulations giving equivalent performances suggests that old concepts developed in allylic oxidation are of limited use to guide the search for active catalysts in an effective way [2]. Conversely, new fundamental data open promising perspectives, and new concepts for reactor design lead to surprisingly high performances. This is due to the fact that new academic laboratories in close contact with industry (namely aware of the "pull") are taking radically new approaches in the search for new active catalysts, new reaction conditions, new types of reactors and new types of operation. Research in Academia and Industry is recognising t h a t investigation should be comprehensive, involving many different approaches to solutions. Recent spectacular laboratory results obtained in methane coupling or oxidative dehydrogenations and elegant technical innovations like the moving bed reactor in butane oxidation demonstrate that considerable progress can be made if a multifaceted approach is taken. These are the points we wish to outline shortly in the following pages. 2. P R E S E N T SITUATION IN OXH~ATION CATALYSIS Catalytic oxidation is a strategic part of catalytic industry. Selective oxidation, oxidative dehydrogenation and ammoxidation, represent about 11% in value of the catalysts consumed by the chemical industry. This does not include auto exhaust catalysts, also active in oxidation, which account for 30% in value of catalyst sales (not very far from the 36% share in value of catalysts used by the whole chemical industry). The growth of selective oxidation catalysts is presently about 11% per year, twice as large as the average (for auto exhaust: 13% world-wide, with large differences according to continents). The catalytic elimination of volatile organic compounds (VOC) or organic matter in water by oxidation still represents less than approximately 10% of the total value for selective oxidation catalysts. Auto exhaust purification, the oxidations of SO2 or NH3 for the manufacture of sulfuric acid or nitric acid respectively, and oxychlorination are in the list of the 20 major processes of the chemical industry; formaldehyde and ethylene oxide manufacture are not far down the list. These few figures summarize both the diversity of the catalytic reactions involving oxygen and their industrial importance. With the permission of the author [3], we borrow here data (Table 1) which indicates the production capacity of the major industrial processes using oxygen for functionalizing hydrocarbons. The production of acetic acid should be added to the list, although 60% of its 6.1 million t/year total world capacity (to reach 67% in the next future) is due to the Monsanto process (methanol carbonylation) [4]. Only the rest (2.4 million t/year) is produced by oxidation of butane or other alkanes or acetaldehyde or, for a small proportion, by the Showa Denko process (oxidation of ethylene).

45 Table I. Production capacities. Million tons per year (Mt/year) Japan North America Europe Formaldehyde (pure) 2.00 4.00 0.70 Acrylonitrile 1.40 1.60 0.60 Maleic anhydride 0.21 0.20 0.10 Phthalic anhydride 0.46 1.30 0.30 Acrylic acid 0.71 0.65 0.42 Methyl methacrylate 0.68 0.70 0.46 Ethylene oxide 4.20 2.50 0.75 G. Hecquet, Plenary Lecture, EUROPACAT II, Maastricht, 1995.

World 6.70 4.50 0.65 3.30 2.00 2.20 11.00

Two facts are striking, when considering this table. The first is that, except for maleic anhydride (MAA), which is now produced from butane, and part of acetic acid (in the liquid oxidation route), the feedstocks used in the other processes are olefins, aromatics or molecules already containing oxygen. The ammoxidation of propane is in an advanced state of development but, on the whole, processes based on paraffins are not yet strongly emerging even though their development constitutes a major "pull". Paraffins are, of course, desirable because they are cheaper feedstocks. Besides, based on present knowledge, it seems that in spite of leading to more extensive oxidation to CO2 and H20, their use would be cleaner, environmentally speaking, because less partially oxidized by-products would be produced. A notable exception is the production of methanol or formaldehyde, where the energy wasting route through synthesis gas will remain more attractive for many years. The second fact, linked to the first, is that a much lower yield in valuable products is obtained when paraffins are used instead of olefins or aromatics. The success of the butane-based MAA process occurred at the cost of a dramatic drop of molar selectivity (at most 65-67% compared to 75-77% when starting from benzene) and a drop of productivity by nearly 20% [3]. Compared to the benzene route, butane oxidation gives molar yields per pass of only about 55% instead of 75%. A third r e m a r k can be made when considering the catalysts used in reactions of light paraffins, butane excepted. It is surprising to discover how large the variety of catalysts which have been claimed to activate light alkanes is. This could be attributed to the fact that we are only at the first stages of progress towards really selective systems, and no catalyst has yet emerged as potentially the best. A notable exception is vanadium phosphate: this is the only performing catalyst for butane oxidation. The a r g u m e n t t h a t a very special surface structure is needed for the concerted oxidation mechanism of butane and easy desorption of the product (a pair of oppositely oriented square pyramides) could explain the exception that vanadium phosphate represents. But acrylonitrile, in the ammoxidation of propane, is also very stable and nevertheless many catalytic formulations seem to give similar results. The diversity of formulations is therefore a fact we must take into account in future investigations concerning selective reactions of alkanes with oxygen. As an example of such a diversity, Table 2 lists some of the catalyst f o r m u l a t i o n s claimed to give good r e s u l t s in p r o p a n e oxidative dehydrogenation. As it is only aimed at underlining the variety of formulations proposed for a single reaction, the table is only sketchy, in the sense that the reaction conditions are not reported. The references and a few more details are

45 given in another article [2]. Various magnesium vanadates have been the object of many studies, but other systems seem to have comparable performances (systems based on cerium, niobium, or vanadium, molybdates and noble metals on monoliths used with very short contact time). Table II. Propane oxidative dehydrogenation Reactant Product Catalyst Conversion Yield Selectivity % % % (oxygen not indicated except if defined phases) Propane Propene Nb based catalysts 7 85 Propane Propene VMg, VMg+Ag, 10 84, 86.9 Electrochemical pumping of oxygen Propane Propene VMg and chloride of 23.1 Cu +, Li +, Ag+, Cd2+ Propane Propene noble metals (Pt,Pd) 100 65 (total Ethene on ceramic foam olefins) monoliths at short contact time, 5 ms Propane Propene 19 60 Co0.95MoO4 Propane Propene V/Mg= 2/1 23 46 2/2 23 59 23 49 2/3 Propane Propene 25 60 VMgfrio2 Propane Propene 20 12.5 62 NiMoO4 43 14.8 34.5 Propane Propene 41.3 33.5 81.1 CeO2/2CeF3/Cs20 Propane Propene 40.3 66.2 FeV-supported Nd203 Propane~ Butane I Alkenes Hexane J Propane Propane Propane Propane

Propene Propene Propene Propene Acrylonitrile

Propane

Propene

Propane

Propene

Propane

Propene

Propane Propane

Propene Propene

Vanadate catalysts V-Fe-Nd-A1 VMg CeO2/CeF3 (NH3)3PO4 + In(NO3)3 + Vanadyl phosphate NiMoOx (x=number determined by Ni or Mo valency) A1203-supported Pt/Cs/Sm MgV206 (50% V205+MgO calcined at 610 ~ CoMoO4/SiO2 NaOH/Na3VO4/A1

50 40.3 10 53.4 12 29

50 66.2

26.7 65

36.7 35 36.7 18.1 91 71

4.1 20.9

16.6

77.9 79.8

47 No new formulations seem to have been proposed for propane oxidative dehydrogenation since 1993, the year of the work cited above, and a literature computer survey indicates only 7 articles which deal with mechanisms, one mentioning catalysts containing Bi, Mo, W, V and Ti [5], and 2 r a t h e r general patents. The literature is not richer for the oxidation of propane to acrolein, with mention of the formulation Ag0.ol Bi0.85 V0.54 Mo0.45 04.0, giving a selectivity to acrolein of 32% at a propane conversion of 52% [6], quite comparable to results obtained in 1991 by other authors [7]. Similar remarks can be made for other oxidation reactions. Considering all the recent work on the various selective reactions of light alkanes with oxygen, three remarks must be made: ~ In spite of the "pull" (functionalization of light alkanes), the intensity of academic and industrial research seems modest. The effort perhaps started too short a time ago to have produced patents or articles. 9 Several lines have been followed, but almost all were inspired by former research r a t h e r t h a n by new approaches. It is striking t h a t the proposed catalysts are similar to those tested in the other selective reactions of alkanes with oxygen, principally, oxidative coupling of methane or oxidation of butane to maleic anhydride. Many of them also have compositions comparable to those of catalysts used for the reactions of olefins with oxygen (molybdates or antimonates) or for dehydrogenations in the absence of oxygen (chromium containing catalysts). Because of the success of vanadyl phosphate in butane oxidation, there is a tendency to focus on vanadium containing catalysts. It is not sure, however, that the data available justify this preference, or suggest the exclusion of other formulation. The "push" in this respect seems rather weak. 9 On the other hand, the reaction of methane, ethane, propane, isobutane and pentanes with oxygen described until now are poorly selective at high, or even moderate conversions. The low stability of the products compared to maleic anhydride may be the explanation. But it could also be argued that the reaction with paraffins is more complex than with olefins. In the case of vanadium phosphate in the oxidation of butane, the idea seems to emerge t h a t a special configuration (two adjacent oppositely oriented flat pyramids) is critical for the smooth r u n n i n g of the concerted process leading to MAA [2,8]. This configuration could permit the adsorption of butane in a correct configuration and the successive intervention of the seven oxygen atoms needed in the reaction of a single butane molecule. If such a special configuration of the active site is essential for selective oxidation of alkanes, it could be hoped to find specific s t r u c t u r e s p a r t i c u l a r l y active and selective in other cases. The requirements are perhaps less stringent, however, because a smaller number of oxygen atoms is required than in butane oxidation. The final r e m a r k of this sketchy section should be made cautiously. It seems that surprisingly, light alkane activation, although being heralded as a major area of future development, appears to benefit only from moderate "push" and moderate "pull". Catalytic combustion and the complete oxidation of pollutants seem areas with much more activity presently. But the "pull" by environmental concerns is strong in those cases.

48 3. 'N-IE ,PUSI-r' EXERTED BY CATALYTIC SCIENCE Keeping in line with the objective of this short article, it is useful to analyse the components of the "push" exerted by catalytic science to innovation in selective oxidation. Ideas and concepts have progressively evolved since the time when the oxidation of olefins and butane oxidation to MAA developed industrially. It is useful to examine critically the potential of these ideas and concepts if we wish to discover new catalysts or improve substantially those mentioned in the literature (such as those of Table 2). We summarize here comments made previously [2,9]. It should be remarked incidentally that, according to a literature survey made at the time of writing this paper, only 5 review papers on selective oxidation have been written since 1993. Except a specific one dealing with bismuth molybdate catalysts, we cite all of them in this article. Role of traditional parameters Doping with minute amounts of elements that authors suppose (often without proof) inserted in an oxide lattice or spread atomically on the surface is in principle a good approach for modifying catalysts. It should, however, be emphasized that it has seldom been verified that the doping elements were really incorporated in the host oxide and, if so, did not spontaneously segregate out. In almost all the cases (unfortunately surprisingly rare) where the check has been made, such a segregation was shown to occur. This suggests that the concepts underlying the addition of dopants (e.g. control of valency of active elements, change in the rate of the reduction-oxidation process at the surface, change of "acidity") may not offer a relevant interpretation of the effects observed. More precisely, they do not seem to do so when used in the way they were in the past. This may suggest that reliable lines for further improvements were missing. It is a pity that ideas which are in principle very well grounded could lead only to uncertain conclusions and modest advances for lack of properly planned experiments. Microanalysis allows the determination of composition at the nanometer level. One can expect that accurate analysis of phases in active catalysts in the state they are during use or quenched to room temperature after use, correlated with activity and selectivity measurements, will lead to substantial progress in the future. Supports are presumed either to be inert (e.g. SiO2), or to permit atomic dispersion or formation of a monolayer (e.g. anatase when V205 is the active phase). Hence the idea that epitaxial layers could be the active and selective species. This view is in principle correct, but incomplete. What is still more disappointing is that the existence of the supposed atomic dispersion or monolayer structure has been verified only in a very limited number of cases, and possibly using sometimes inadequate tools. For example, strong doubts begin to be raised concerning the atomic dispersion of V205 in the form of a monolayer on anatase. The most likely structure seems to be that of severallayer thick islets of a suboxide of vanadium (V6013) [10,11]. These islets would be stabilized by epitaxy, thus permitting the suboxide to lose and reincorporate oxygen during the catalytic cycle without irreversible transformation to a different structure. This is in line with a view which is beginning to emerge, namely that the emphasis should be laid on the catalyst in the state it is while it works during the catalytic cycle. Systematic examination of used catalysts could have avoided misleading interpretations. In particular, this could have

49 avoided to attribute a catalytic behaviour to monolayer formation in cases where transformation to islets occurs in the reaction conditions. Epitaxy may reasonably be mentioned for the VOx/anatase system. But there are other cases where epitaxy is doubtful. There is no detectable evidence of Sb204 decorating epitaxially FeSb204. If the picture of oppositely oriented flat pyramids as active sites for butane oxidation is correct, the hypothesis attributing a crucial role to an epitaxy between various vanadium phosphates may lose credibility. On the whole, the parameters which have been generally considered in the last 25 years are certainly relevant in principle, but the conclusions or speculations concerning their real intervention in given catalysts deserve careful verification using modern physico-chemical techniques. The minimum precaution is to check whether the feature mentioned (e.g. doping or epitaxy) survives after the catalyst has worked at realistic catalytic conditions. Are there n e w lines for the search of n e w catalysts ? Several facts suggest new directions for research aimed at the discovery of new catalyst formulations or structures, and new routes for catalyst improvement. Two critical reviews of the role of important parameters in selective oxidation have been presented recently. One of them focuses on factors which have perhaps not attracted sufficient attention, like mode of adsorbate bonding and site isolation, together with new views on more traditional parameters [12]. On the basis of this and the other review [2], we wish to show that the real picture is substantially different from the traditional one. The essential reason is that most academic studies in the past implicitly assumed that the solids submitted to characterization as freshly prepared were already in the catalytically active form. A real and credible "push" is gaining force now. It rests on the extensive characterization of catalysts after they have worked catalytically or during their work (a still better approach). This "push" is constituted of lines of work inspired by newly discovered facts. These are the n a t u r e of the real phases active catalytically, their texture (e.g. the crystallographic phases really exposed to reactants) and the relation between these phases. Three lines of work, among others, can be cited. One is typically represented by the work dealing with the changes in structure and texture of VPO catalysts during activation in the presence of the reacting mixture (butane or propane with 02), by J.C. Volta et al. [13,14]. R.A. Overbeek and J.W. Geus took a complementary approach [15,16]. Another line is that of P.L. Gai-Boyes concerning the defect structure of catalytically active phases during catalysis [17]. The third line deals with the cooperation between phases. One important objective is thus becoming clearer, namely synthesising the phase or phases which are really active during catalysis. Nowadays, this objective can be reached, because we have the tools to identify these phases. This "push" will gain strength in the future, when the impressive inventiveness of catalysis scientists for finding new preparation methods will be directed to the synthesis of these phases. A more extensive characterization of catalysts progressively suggested that the key to activity and selectivity might be the existence of s e v e r a l phases in catalysts. As regards to this aspect, the ideas have been progressing thanks to O.V. Krylov, I. Matsuura, Y. Moro-Oka and U.S. Ozkan. The group of the present author focused on it. In this paper, citations and illustrations will

5o come from other groups. Breiter et al. [18] demonstrated that, at industrial conditions, a phase inactive in the oxidation of acrolein to acrylic acid, namely NiSb206 (or CoSb206) considerably increased the performance of an already optimized catalyst of formula Mol2V3Wl.2Cu2.2Ox. They attributed this effect to the oxygen-donor properties of NiSb206 or the other compounds they used, namely Co85206, CuSb206 and Sb204. These properties were in line with those indicated by the remote control concept [19]. Another set of experiments concerns the reaction of lower alkanes. Those reported in fig. 1 show that MgMoO4 and MoO3 exhibit a cooperative effect in the oxidative dehydrogenation of propane. The result is completely in line with the oxygen donor-acceptor scale resulting from the remote control theory [19]. ~

10

0 100 ",~

~

500~ '

~

.

.

5

5

0

~

(b

500~ 550 ~

#, 90

l

0.5

Moo3 /I.9.oo

1

. .oo3 j

0

10 Propone conversion %

Fig. 1. Oxidative dehydrogenation of propene on MgMoO4-MoO3 mechanical mixtures (C3H8: O2:He=1:1:25) MoO3/(MgMoO3+MoO3): O 5 wt% 9 20 wt% (after M.C. Abello, M.F. GSmez, L.E. Cadus, Actas, XV Simposio Iberoamericano de Cat~lisis, C6rdoba, Argentina, (1996), pp. 233-238.) Last but not least, the discovery of titanium silicalite and its activity in an increasing number of new oxidation reactions not only opens promising prospects of future discoveries with respect to reactions, but also suggests that materials with new catalytic properties will be discovered. Their distinctive features should logically be a control of coordination of the active atom or atoms that zeolites permit more easily than other structures, and shape selectively. Two recent papers review the potential of zeolites as oxidation catalysts [21,22] and very stimulaing comments are made in a third one concerning the oxidation (mainly epoxidation) of large molecules [23]. We mentioned 3 promising lines in this section, namely: catalysts "as they work", catalysts in which remote control is optimised, and zeolites. The first two lines take cognisance of the fact that catalysis is a dynamic process in which the solid, especially in the case of oxidation, cyclically undergoes extensive superficial changes, and its surface is the theater of a continuous

51 movement and transformation of adsorbed species. This suggests t h a t the above advances and others not cited here offer a vast potential for catalyst improvement in selective oxidation. Composition of the gas phase When comparing test conditions in most investigations carried out by Academia with those used in Industry, striking differences may appear. In several processes, for example, steam is injected together with the hydrocarbon and oxygen, but very few university researches included the influence of steam. Practically each time reactions with and without steam have been compared, results turned out to be different. Possible explanations may be: 9 role as a simple diluent (including a possible diminution of the surface concentration of reactants); 9 influence on the thermodynamics, an explanation sometimes quoted for the simple catalytic dehydrogenation of ethylbenzene to styrene in the absence of oxygen (although it is not clear why steam could achieve more favorable thermodynamic conditions); 9 more efficient heat transfer; 9 t u n i n g of catalyst surface acidity; this explanation seems relevant in principle, because other parameters do bring about measurable modifications of the number of BrSnsted sites on the surface of catalysts; water molecules could very easily do the same [19]; 9 other more complicated reactions of water with the catalyst surface [24]; 9 modification of homogeneous oxidation reactions: when comparing the homogeneous oxidative dehydrogenation of propene in the absence and presence of steam, it was found in the second case t h a t a relative increase in the selectivity to propene by more than 25% occurred at low conversions and remained substantial at high conversions [2]. Feed composition does not only concern steam. Industrially, additives are often constantly added to the feed. A reason frequently advanced is t h a t the catalyst could in this way be replenished of elements it loses continuously. This suggests an area of research which should have fascinated Academia, because the use of such additives helps control the degree of doping during the whole life of the catalysts. In this case, the restriction concerning doping expressed above is not valid any more, because the doping elements are fed continuously rather than just introduced in the fresh catalyst. A still more exciting line of research based on modifications of feed composition is the addition of hydrogen in catalytic oxidation. This seems absurd at first sight: one speculates that hydrogen would react immediately with oxygen just to form water! The experiments reported in fig. 2 give surprising results, namely the unexpected catalytic transformation of propene to propene oxide, and the selective oxidation of propane and isobutane respectively to acetone and t-butanol (25,26]. The catalyst is strange (Aufl~i02) and the yields very low. This result and a few other ones mentioned in the literature, however, trigger interesting questions. The role of hydrogen cannot be just to react with oxygen to produce water, because the amount is too small. Is it to produce heat locally on the surface of the catalyst, or to create new surface intermediates? If we accept the view that spillover species can play a role (remote control), a speculation along the following line can also be proposed. In selective oxidation, the remote control operates so t h a t spiUover

52 oxygen keeps catalysts more oxidised than in the state normally resulting from the balance between surface reduction and oxidation during the catalytic cycle. It is believed, and this rightly at least in butane oxidation on VPO catalysts, that the surface must be kept relatively reduced in alkane oxidation, in order to catalyse the initial activation steps. The speculation could therefore be that hydrogen (spillover hydrogen?) could impose a more favorable reduced state to the catalyst. A still stranger reaction involving both air and hydrogen, namely the oxidation of benzene to phenol, will be the object of more comments near the end of this article. A recent review paper stresses in a more general way the importance of the reactive atmosphere in selective oxidation [27]. 0.20 C3

E

{b

E -

1.0

-

0.5

"~

0.15 propylene oxide

~3

oc

0.10

co 2

0.05

acetone co 2

propene (SO~

acetone propane (80~

t-butanol isobutane (80~

0

Fig. 2. Oxidation of propene, propane and isobutane over Aufs in the presence of hydrogen. Feed: H2:O2:hydrocarbon:Ar = 1:1:1:7; flow rate 2000 ml h-l; catalyst AudiO2 (1 wt% Au), 0.5 g). Yields are calculated in mol% on the basis of the starting hydrocarbon (after T. Hayashi, M. Haruta, Shokubai, 37 (1995) 75). Intervention of homogeneous gas phase reactions In this section, we again select the case of light alkanes functionalization in our attempt to discuss fruitful lines of research. It is well known t h a t gasphase reactions play an important role in methane oxidative coupling. This is expected, as this occurs at very high temperature. But this is a general phenomenon. Contrary to the case of olefins, homogeneous catalytic oxidations of alkanes with more than 3 carbon atoms proceed at temperatures similar to those of the catalytic reaction and these are relatively low. This probably has

53 led to m i s i n t e r p r e t a t i o n of data in certain cases, namely a t t r i b u t i n g homogeneous reactions to catalysis. A very interesting contribution in this respect is that of Burch and Crabb, who investigated in detail the role of homogeneous and heterogeneous reactions in the oxidative dehydrogenation of propane [28]. The reaction temperature for the catalyzed reaction was about 130 ~ lower, but differences with the homogeneous reaction depended on the oxygen/hydrocarbon ratio. The unexpected result is that there are similar conversion vs. selectivity relationships for both the heterogeneous reaction and the homogeneous one with most catalysts. Even the best catalysts are not better than no catalyst at all at the higher temperatures. This could seem pessimistic, but obviously does not rule out t h a t new catalysts could give a decisive advantage to catalyzed reactions. Incidentally [2], the homogeneous reaction in the presence of steam seemed to exceed the performances of the catalysts tested in the work cited above [28]. Burch and Crabb noted in the abstract of their paper: "A combination of homogeneous and heterogeneous contributions to the oxidative dehydrogenation reaction may provide a means of obtaining higher yields in propene". This prediction seems to be supported by at least one result concerning a very special catalytic system which may combine homogeneous and heterogeneous processes. This is constituted of lithium hydroxide/lithium iodide melts, which give considerably higher propene yields at higher propane conversion than either homogeneous reactions or reactions catalyzed by solid catalysts [29]. It has been mentioned that both maleic and phthalic anhydrides can be produced in the catalytic oxidation of n-pentane in the presence of VPO catalysts. With n-pentane, however, the homogeneous reaction begins to be significant above 300~ a fact the consequence of which has perhaps been underestimated. By using reactors with empty spaces of different volumes (lengths), it is possible to detect the influence of both the heterogeneous and homogeneous reactions. The non-selective homogeneous reaction increases npentane conversion, but the surprising finding is that the maleic/phthalic anhydride selectivity varies substantially, from 1.8 with as little empty space as possible to over 4 with substantial empty space [2]. This demonstrates that the homogeneous reaction, which begins to play an important role in the oxidation of n-pentane in the range of temperature where catalysts like VPO are active (around 350-400 ~ leads to a modification of selectivity. Conclusion concerning the role of catalytic science Concluding this section, it seems that research on oxidation catalysts and oxidation reactions provides some push to innovation, through (i) the high level of activity in the synthesis of new structures and development of new approaches to catalyst fabrication, (ii) a better knowledge of the state of catalysts during their use, (iii) the role of surface mobility, spillover and remote control, (iv) the use of additives in the feed and (v) a better understanding of the contribution of homogeneous gas phase reactions. 4. 22-IE"PUSH" EXERTED BY THE USE OF NEW TYPES OF REACTORS AND NEW CHEMICAL ENGINEERING CONCEP2~ According to the philosophy of this overview, we shall comment on some of the facets of present science and technology in catalytic oxidation. Because of

54 the importance of light alkane oxidation as a "pulling factor", we shall again consider principally this group of reactions when examining the role of chemical engineering. One must recognise that the headlines concerning selective oxidation of light alkanes in the last years did not concern spectacular advances in catalytic science, in the narrow sense, but rather resulted from the use of reactors which had not been used traditionally in catalytic oxidation. The impressive yield in olefins (65%) using a monolith reactor at very short contact times (5 ms) is not approached by any other result mentioned in Table 2. This result, due to Huff and Schmidt [30] demonstrates that employing a type of reactor not used previously in selective oxidation, but rather in complete oxidation, and very short residence times can lead to promising prospects. It can also be underlined that the active catalytic species used in this work are regarded as among those most active in complete oxidation (platinum in particular). This constitutes an outstanding case where the type of reactor is much more important than the intrinsic selectivity of the catalyst for reaching high selectivities. But recent results also show that the chemical engineering approach as a whole has a key role to play in the development of catalytic oxidation. To illustrate this, the example of methane oxidative coupling is considered. In spite of a wide recognition of the importance of homogeneous reactions, an overwhelming fraction of research has been directed to the discovery of new catalysts (perhaps over 90%) with only a very small number of investigations trying to take these homogeneous phenomena into account. The progress has been deceivingly modest. This has led respected industrial scientists to discourage further research. It was right to recognize t h a t the results published were very far from being economically attractive. But, instead of discouraging research, these scientists should have spurred research and, at the same time, stressed that this should be conducted on different lines. One such possibility is clearly to design reactors taking into account the specificity of homogeneous-heterogeneous reactions (perhaps inspired by the ultrafast monolith reactor of Huff and Schmidt). But considering the whole apparatus or the whole future chemical plant can also be very fruitful, even without changing the reactors. Two recent results convincingly demonstrate that integrating recycle and separation features with a catalytic reactor leads to impressive yield. Tonkovich et al. reached a 50% yield in C2 hydrocarbons in the oxidative coupling of methane. They used a moving bed reactor, thus permitting some sort of a chromatographic separation [31]. The problem remained of the high reactivity of ethylene compared to CH4. One way to deal with this was to recognize that the reaction of methane to ethylene can be extremely selective at very low methane conversions. Considering this, Vayenas et al. achieved an ethylene yield of 85% (calculated on the carbon contained in CH4) [32]. The key was alternating the highly selective adsorption of ethylene, ethane and CO2 on a 5A molecular sieve with periodical release. Conversion was kept very low, and the non-reacted methane recycled. A striking fact was that the catalyst did not belong to the group giving the best performance in simple flow reactors. Besides silver on which oxygen was "pumped" electrolytically, the authors used a Sm203(20%)-CaO(l%)-Ag(79%) catalyst simply contacted with molecular oxygen from the gas phase (both systems gave approximately the same results). Not only did the reactor concept offer new perspectives, but the results proved (as the result of Huff and Schmidt

55 suggested) that a key to success is to adapt catalysts to reactors and operation conditions and vice versa. Another conspicuous illustration of this remark is the success of the recirculating fluidized bed reactor developed by DuPont for butane oxidation (using the so-called "riser" reactor). The catalyst is now more a "carrier" of selective oxygen t h a n a real catalyst. Its characteristics, particularly resistance to attrition combined with easy access of the reactants to particles situated inside the silica shell, resulted from an adaptation to the new type of reactor. The reactor itself necessitated innovative developments for taking advantage of the "oxygen carrier" properties of the catalyst, and to overcome a drop of productivity in MAA per pass. The development of the reactor is a real achievement considering that, contrary to those using classical fluidized beds, this kind of recirculated bed reactor has not been developed very much in industry (except in platinum catalyzed reforming). Another conspicuous innovation is the ultrafast oxidation reactor developed by BASF for the oxidation of isoprenol (2 methyl-butl-ene-4 ol.) to isoprenal for the synthesis of citral:

02

~~...~.CHO

This is a continuous flow reactor operating at 500~ with a residence time of the order of 1 millisecond. The yield in the extremely sensitive isoprenol product is 95% [33]. Other innovative developments have appeared, pertaining to both the process and the reactors. The SMARTsm reactor design derived from the Styro-Plus process of UOP is the key to a more efficient route from ethylebenzene to styrene by non oxidative dehydrogenation. Hydrogen is formed in a first dehydrogenation reactor. The SMARTsm reactor combines 2 catalysts: (i) a platinum oxidation catalyst that reacts the hydrogen produced in the first reactor with oxygen to generate heat (ii) a dehydrogenation catalyst. This innovation permits a large increase in productivity by retrofitting of existing radial flow reactors [34]. Also considering complete oxidation, an interesting tendency appears in catalytic combustion, which can be optimised by using 2 (or several) monoliths in series [35]. Other new concepts are certainly developing, for various simple catalytic oxidations as well as for complex successions of catalytic oxidation and/or catalytic dehydrogenation (new routes to methacrylic acid). Catalytic combustion is the object of many efforts (as the ones of Catalytica Associates and Osaka Gas Co.). The idea of reactors integrating several functions (for example, reaction and separation, as in catalytic membranes; or oxidation coupled with an energy generator, e.g. a turbine) is attracting increased attention. For example, the recently commercialized combined cycle coal gasification technology offers unprecedent efficiency in the transformation of coal to electricity. It is still difficult to discern the driving forces in the development of these new reactors. Three of them are perhaps dominating: (i) control of the homogeneous-heterogeneous processes (methane oxidative coupling, monoliths, catalytic combustion) (ii) full efficient use of the unusual potential of selective oxidation catalysts, which, after all, are reduction-oxidation "reactants" (riser reactor)

56 (iii) integration of the catalytic step in a more complicated network of functions to be carried out by a system (several reactions in a same reactor; reaction and separation; combustion and production of energy). 5. PERSPECTIVES Specialists of chemical engineering and catalysis scientists are certainly ready to agree that a good process needs a good catalyst. The conclusion of the previous paragraphs was that the "push" concerning catalysts is perhaps not strong enough. On the other hand, one discerns a strong "push" in the invention of new reactors or new process designs. It seems likely that the "push", namely the innovation thanks to fundamental concepts, in selective oxidation will essentially be exerted by a close cooperation between catalytic science and chemical engineering in the next years. This necessity of close cooperation clearly appears when considering the recent successes. On the whole, oxidation catalysis conveys a positive image of dynamism and inventiveness. Improvements of already existing oxidation processes are continuously made (in MAA manufacture, with the riser reactor by DuPont, or in oxychlorination, by Montecatini Technologie and ICI). In addition, and still more clearly demonstrating the dynamism of industrial catalytic oxidation, completely new catalysts are discovered, especially with the titanium silicalite which permits the synthesis of hydroquinone from phenol, selective epoxidations, oxidations of alcohols to aldehydes, and the manufacture of cyclohexanoneoxime. New processes have appeared in the last ten years, in addition to the now well established routes from isobutene to methylmethacrylate (Asahi Chemical Ind.). These are: 9 oxidation of methylal to formaldehyde: 02 CH30-CH2-OCH3

~ 3HCHO + H20 300 - 400 ~ FexMoyOz

9 very selective gas phase oxidation of ethylene to acetic acid (Showa Denko). 9 oxidation of 4-methoxy toluene (p-methyl anisole) to paramethoxybenzaldehyde over a catalyst containing vanadium and alkali metals, around 400~ (Nippon Shokubai and Kagaku Kogyo). 9 direct access to methyl isocyanate from methylformamide by oxidative dehydrogenation over a silver catalyst (DuPont). 9 direct oxidation of benzene to phenol catalyzed by a solid catalyst (0.5% Pt/20% V205/Si02) under pressure at 200~ at unusual conditions which will be mentioned further in this section (Tosoh Corp.). It is striking t h a t many new processes get integrated in a new manufacturing chain, as the oxidations of methyl methacrylate were. Other examples are: 9 oxidation of butane to maleic anhydride and tetrahydrofuran (DuPont)

57 acetoxylations, as a route to 1,4 butane-diol and tetrahydrofuran on Pd catalysts containing additives (Te, etc.) (Mitsubishi Kasei): butadiene + 2AcOH + 1/202 -~ CH3COOCH2-CH=CH-CH2OOCCH3+H20 It is remarkable that this reaction and the direct hydroxylation of benzene cited above occur in a 3-phase system: gas-liquid-solid. The oxidation of benzene constitutes an extraordinary case in that the liquid is a mixture of benzene and acetic acid and the gas phase contains both air and hydrogen! This suggests the formidable problems that chemical engineering had to solve to develop these processes. The spectacular results obtained in the most extensively investigated reaction, namely the oxidative coupling of methane, the development of ultrafast catalytic reactors and catalytic combustors and the above examples constitute clear indications that we are still at the beginning of a strongly innovative research period in catalytic oxidation and, in particular, concerning the functionalization of light alkanes. This is a clear indication that the discoveries triggering these innovations will have their roots in all the relevant scientific fields, heterogeneous catalysts as well as homogeneous reactions, discovery of new catalyst formulations as well as new chemical engineering concepts. 6. ACKNOWLEDGMEN'I~ Many colleagues assisted the author in the preparation of this article, by giving suggestions and comments and sending documents. I wish to acknowledge their help here: J. Bongaarts, J.F. Brazdil, M.G. Clerici, G. Hecquet, H.H. Kung, J.J. Lerou, P.G. Menon, S.T. Oyama and P. Ruiz. 7. REFERENCES 1. J. Rabo during his plenary lecture, in "New Frontiers in Catalysis" (L. Guczi, F. Solymosi, P. T~t6nyi, eds.), Akad~miai Kiad6, Budapest, 1993, pp. 1-29. 2. B. Delmon, P. Ruiz, S.R.G. Carraz~n, S. Korili, M.A. Vicente Rodriguez, Z. Sobalik, in "Catalysts in Petroleum Refining and Petrochemical Industries 1995" (M. Absi-Halabi, J. Beshara, H. Qabazard, A. Stanislaus, eds.), Elsevier, Amsterdam, 1996, pp. 1-25. 3. G. Hecquet, Plenary Lecture, 2nd European Congress on Catalysis, EUROPACAT II, Sept. 3-8, 1995, Maastricht (The Netherlands), 1995. 4. Chem. Eng. News, 1996, July 1, p. 7. 5. J. Barrault, L. Magaud, M. Ganne, M. Tournoux, in "New Developments in Selective Oxidation II" (V. Cortes Corber~n, S. Vic Bell6n, eds), Elsevier, Amsterdam, 1994, pp. 305-314. 6. J. Li, W. Song, B.-Sh. Dou, Yongyong Huaxue, 13 (1996), 88. 7. Y.C. Kim, W. Ueda, Y. Moro-Oka, Appl. Catal., 70 (1991), 175. 8. J.R. Ebner, M.R. Thompson, Catal. Today, 16 (1993), 51 (see also in ref. 2).

58 0

10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20.

21. 22. 23. 24. 25. 26. 27. 28. 29.

30. 31. 32. 33.

J.T. Wrobleski, B. Delmon, H.H. Kung, Y. Moro-Oka, J.L. Cihonski, "Catalyst Modification: Selective Partial Oxidation", Catalytica Study nr. 4190CM, Catalytica, Mountain View, CA (1991). T. Machej, M. Remy, P. Ruiz, B. Delmon, J. Chem. Soc. Faraday Trans., 1990, 86, 715. T. Machej, M. Remy, P. Ruiz, B. Delmon, J. Chem. Soc. Faraday Trans., 1990, 86, 723. S.T Oyama, in "Heterogeneous Hydrocarbon Oxidation" (B.W. Warren, S.T. Oyama, eds.), ACS Symp. Ser. 638, Am. Chem. Soc., Washington (D.C.), 1996, pp. 2-19. C.J. Kiely, A. Burrows, S. Sajip, G.J. Hutchings, M.-T. Sananes, A. Tuel, J.C. Volta, J. Catal., 162 (1996) 31. J.C. Volta, Catal. Today, 32 (1996) 29. Eur. Pat. EP942011177.6, ass. to Engelhard De Meern. R.A. Overbeek, PhD Thesis, Univ. Utrecht, 1996. P.L. Gai-Boyes, Catal.Rev.-Sci.Eng., 34 (1992) 1. S. Breiter, M. Estenfelder, H.-G. Lintz, A. Trenten, H. Hibst, Appl. Catal., 134 (1996), 81. L.T. Weng, B. Delmon, Appl. Catal. A, 1992, 81, 141. M.C. Abello, M.F. GSmez, L.E. Cadus, Actas, XV Simposio Iberoamericano de Cat~lisis, CSrdoba (Argentina), Sept. 16-20, 1996 (E.R. Herrero, O. Anunziata, C. Perez, eds.), Unica Ed., CSrdoba, 1996, pp. 233238. E. Hoeft, K.H. Kosslick, R. Fricke, H.-J. Hamann, J. Prakt.Chem./Chem. Ztg., 338 (1996) 1. P. Ratnasamy, R. Kumar, in "Zeolites: a Refined Tool for Designing Catalytic Sites" (L. Bonneviot, S. Kaliaguine, eds.), Elsevier, Amsterdam, 1995, pp. 367-376. A. Baiker, in " l l t h International Congress on Catalysis - 40th Anniversary" (J.W. Hightower, W.N. Delgass, E. Iglesia, A.T. Bell, eds), Elsevier, Amsterdam, 1996, pp 51-61. J.-M. Jehng, G. Deo, B.M. Weckhuysen, I.E. Wachs, J. Mol. Cat., 100 (1996) 41. T. Hayashi, M. Haruta, Shokubai, 37 (1995) 75. M. Haruta, Catal. Today, in press. E.A. Mamedov, Appl. Catal. 116 (1994) 49. R. Burch, E.M. Crabs, Appl. Catal. 100 (1993) 111. I.M. Dahl, K. Grande, K.-J. Jens, E. Rytter,/~. Slagtern, Appl. Catal., 77 (1991) 163. M. Huff, L.D. Schmidt, J. Catal. 149 (1994) 127. A.L. Tonkovich, R.W. Carr, R. Aris, Science, 262 (1993) 221. Y. Jiang, I.V. Yentekakis, C.G. Vayenas, Science, 264 (1994) 1563. W.F. HSlderich, in "New Frontiers in Catalysis" (L. Guczi, F. Solymosi, P. T~t~nyi, eds.), Elsevier, Amsterdam, 1993, p. 127-163.

59 34. T. Imai, R.R. Herber, G.J Thompson, D.J. Ward, AIChE National Meeting, New Orleans (LA), March 1988. 35. M.F.M. Zinkels, S.G. J~ir~is, P.G. Menon, in "Structured Reactors and Catalysts" (J.A. Moulijn, A. Cybulski, eds.), M. Dekker, New York, 1997.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

61

Molecular approach to active sites on metallic oxides for partial oxidation reactions Jacques C. Vrdrine Institut de Recherches sur la Catalyse, CNRS, UPR 5401 associ6e /t l'Universit6 Claude Bernard, Lyon I, 2 avenue A. Einstein, F-69626 Villeurbanne, France

Abstract Through a variety of examples, it is shown that the catalytic oxidation reactions, which operate via a Mars and van Krevelen mechanism, imply ensembles of atoms containing variable valence cations i.e. Lewis acid and redox sites and Lewis bases such as 02% O H or PO 43- able to act in the different elementary steps. The structure and the size of these ensembles, defined as <~, determine their reactional specificity. The examples chosen show that the <<designs >~of such ensembles are strongly dependent on the nature of the material, on its morphology (structure sensitivity as for MOO3, vanadyl phosphates), and on its structure (e.g. vanadyl phosphates, iron phosphates, heteropolyoxometallates, titanosilicates). It is also possible to build such ensembles upon dispersion of an active oxide on a relatively inert support (e.g. MoOx/SiO2, VOx/TiO2, NbOx/SiO2, TiOx/SiO2, etc).

Keywords: Molecular design; Active sites; Partial oxidation; Metallic oxides 1. INTRODUCTION Since the original view in Catalysis of an active site as single atoms various concepts have been developed in particular that of structure sensitivity. In partial oxidation reactions the Mars and van Krevelen mechanism proposed in 1953 is usually occurring. It involves both a redox function of the solid catalyst surface and oxygen insertion into the reagent molecule from lattice oxygen ions. The reaction is rather demanding since it involves for the substrate H atom abstraction, O atom insertion and electron transfer. This occurs on given active sites which should then obligatorily be of a certain size from several atoms up to a complete crystalline face or even a given phase. In the late seventies, early eighties a concept of cristalline faces being active for a partial oxidation reaction has been proposed by J.C. Volta for MoO3 single crystals (1). Many other examples have been further published with the same approach and show clearly that for an oxidation reaction much more than one surface atom is necessary to describe the active site. A new concept has then been developed [2, 3] considering an ensemble of surface atoms insuring the complex reaction mechanism to take place at the catalyst surface. Such a

62 molecular approach concept considers several cations and oxygen ions as constituting the active site. It is known that oxygen species on an oxide surface may exhibit different properties being considered as monoatomic species with a more or less electrophilic or nucleophilic character (O', O 2, ...) or diatomic species (O2", peroxo, ...). The more or less electrophilic character of the oxygen species is important since it may favor the deprotonation of the hydrocarbon molecule (nucleophilic attack) or the allylic attack (weak polarisation of M = O bond) and the direct attack of a double bond or of an aromatic ring (electrophilic attack). Electrophilic oxygen species correspond to O', superoxo (M-O-O) or peroxo (02") while nucleophilic species are rather metal oxo (M = O). The electrophilic attack occurs in the region of high electron density of the reagent molecule as at the YI bond. In such reaction when lattice ions are concerned an anionic vacancy is created and should be replenished by anion transfer process via lattice oxygen anions. It was then suggested that metal oxide materials exhibiting extended defects as shear plane structures should be favorable for such reactions since they exhibit free valency space and thus could favor a redox mechanism to occur during the catalytic reaction. Such a transformation should obviously be reversible and should then correspond to very localized crystallographic modifications like in topotactic type transformations. This also explains why non perfectly crystallized surfaces are more active than well crystallized one. The presence of local defects like anionic vacancies or coordinatively unsaturated cations plays certainly a role which is difficult to characterize due to the lack of ordering necessary in mainly physical techniques to be detectable, in particular in the case of X ray diffraction technique. Majority of the catalysts correspond to metallic oxides with V or Mo as one of the key elements but also cations of variable oxidation states as Fe3+/Fe2+, Cr6+/Cr3+, Cu2+/Cu+, Sb3+/Sb5+, etc. Some metals (mainly Ag for ethylene epoxidation), noble metals (as Pt, Pd), zeolites (titanosilicalite TS-1 from ENI for phenol oxidation) and heteropolyoxometallates (e.g. H4PMOlIVO40for isobutene oxidation to methacrolein) may also be used. Several examples have been chosen in this presentation to illustrate how such a concept may be valid in oxidation reactions and how it can be determining for the choice of the metallic single or mixed oxide and/or for the choice of its preparation procedure. 2. GENERAL REACTIONS

FEATURES

OF OXIDATION CATALYSTS AND OXIDATION

2.1. Oxidation catalysts The oxidation catalysts are usually mixed oxides which operate according to the redox process suggested by Mars and van Krevelen. According to this mechanism the substrate is oxidized by the solid and not directly by molecular oxygen of the gaseous phase. The r61e of dioxygen is to regenerate or to maintain the oxidized state of the catalyst. The oxygen atom(s) introduced into the substrate (or giving H20 for oxidative dehydrogenation reactions) stems from the lattice. This mechanism involves the presence of two types of distinct active species: an active cationic species which oxidises the substrate and another species active for dioxygen reduction. An adequate structure of the material should also facilitate both electrons and oxygen species transfer.

63

2.2. Oxygen species The oxygen atom incorporated into the substrate stems from the lattice and is at -2 oxidation state. Its replacement by molecular oxygen necessitates electrons according to: 02 + 4e--) 202. . This process has its own kinetics related to the reactivity of the sites with oxygen, their concentration, the efficiency of electron transfer, the partial pressure of oxygen, etc.. Usually, it is much faster than the oxidation of the substrate i.e. it is generally admitted that the rate determining state is the substrate activation. The homolytic fragmentation of a C-H bond in the coordination sphere of the acceptor metal ion may occur via a transfer of the hydrogen to the oxygen ion at -2 oxidation state. This is a concerted action with homolytic breaking of metal - oxygen bond which transfers one electron to the metal. Without any hypothesis about the nature of the metal-oxygen bond one can write the reaction with formation of a 7t-alkyl complex (as usually admitted) or a 8-alkyl complex. Depending on the nature, oxidation state of the metal ion and its environment (coordination structure), the metal-oxygen bonds may be more or less polarized and therefore the oxygen ion may exhibit electrophilic or nucleophilic properties. One may distinguish three extreme cases: 8 + 88- 8 + a: M = O (nucleophilic) --~ 02. b: M = O --->Oand c: M = O (electrophilic) --+ Q I Each case corresponds to specific properties. Case a (nucleophilic character) will intervene in activation of a C-H bond in ot of the double bond or of an aromatic ring; Case b (weakly polarized) will favor concerted homolytic - type reaction as for allylic dehydrogenation of olefins. Case c (electrophilic character) allows one a direct attack of a double bond (oxidative breaking). Let's take some examples which will be considered for some of them in more details later m

CH3wCH:CH2 propene

+ 2(O2-)

~

CH3~CH--CH2

+ NH3 + 3 (O2-)

OHC--CH:CHz acrolein

+

H20 + 4e-

NC--CH:CHz + 3 H20 + 6eacrylonitrile

propene

O O

CH3--CH2-CH2--CH 3 + 7 (O2-)

isobutyric acid

4H20 + 14e-

O maleic anhydride

butane

CH3N ,O /CH--C + CH3 OH

+

02-

CH2%c--c ,O CH/

+

"OH

methacrylic acid

H20

+

2e-

64 It clearly appears that a single and isolated metallic ion site cannot take into account all the necessary transformations involved in the reaction since several steps as replenishing of oxygen anion vacancies, H atoms extraction and electrons transfer are concerned. For instance n-butane oxidation reaction to maleic anhydride necessitates 7 lattice oxide ions, 8 hydrogen atoms abstraction from the substrate, 3 oxygen atoms insertion and 14 electrons transfer! 3. STRUCTURE SENSITIVITY OF OXIDATION REACTIONS ON OXIDES Such a concept has been introduced by M. Boudart on metals. It has been introduced for oxides in the late seventies by J.C. Volta et al [ 1], or early 80'ies by J.M. Tatibouet and J.E. Germain [4], J. Haber et al [5], etc and it is widely accepted at present. For instance in the work by J.C. Volta et al, it was shown that single crystal type samples of MoO 3 exhibiting different relative amounts of the different faces (010) basal, (100) side and (101) and (101) apical, exhibit different activities and selectivities in the oxidation of propene to acrolein and COx [6]. The originality of the work was to synthesize crystals of various shapes by epitaxial growth via oxyhydrolysis of MoC15 inserted between the layers of graphite. Table 1 summarizes the main results obtained for propene, but-1-ene and isobutene oxidation on MoO3 crystals . It clearly appears that for propene oxidation the (100) side face is selective for acrolein formation and the (010) for total oxidation. At variance for isobutene oxidation the side face gives both methacrolein and COx while the basal plane has low activity for acetone formation. Such specificity depends on the hydrocarbon molecule. It may thus be proposed that the stereochemistry of the hydrocarbon molecule and that of the oxide face play a determining rSle. Table 1 Structure sensitivity of the different faces of MoO3 crystals in olefin oxidation at 380~ (from ref 6).

Reactant Olefin Propene But-l-ene Isobutene

Products Acrolein CO, CO2 Butadiene CO, CO2 Methacrolein acetone CO, CO2

Basal (010) 0.06 1

Relative Selectivity per face Side Apical (100) (101), (101) 2.3 0.7 0 0

3

9.3

2

1 0 0.06 0

0 0.6 0 1

0 0.1 0.06 0

A more precise analysis and characterization of the MoO3 crystallites shape have shown that in fact the better plane for propene oxidation to acrolein corresponds to the (lk0) plane as schematised in fig. 1 [7]. It is then suggested that the propene activation (H abstraction) into the rt-allyl intermediate occurs on the side (100) plane while the O atom insertion occurs on the (010) basal plane. As a matter of fact it is worth noting that due to the layered structure of MOO3, the lattice oxide ions are much more labile in the (0k0) plane than in the others. Recently such a structure sensitivity was also derived for the partial oxidation of methane to formaldehyde 18].

65

(loo)

[010]

Q

@o13

~

1oo]

",(120) ',,

, ,,

Fig. 1. Cross section view of otMoO3 (100) and (120) planes (projection of the lattice on the (001) plane) (from [7]). 4. VANADYL P Y R O P H O S P H A T E [9] Such a catalyst is well known for the oxidation of n-butane into maleic anhydride. The preparation necessitates the formation of VOPO4, 0.5H20 as a precursor synthesized in an aqueous or better in an organic medium and its activation in a flow of 1 to 2% butane in air at the reaction temperature (ca 380~ Here also, the preparation and the activation of the samples appeared to be particularly crucial in order to obtain a well performing catalyst. In all cases, whatever the catalysts being good or exceptionally good, the (VO)2P207 phase (V 4+ cations) as a main constituent was detected by X Ray Diffraction and by in situ Laser Raman spectroscopy [10] in addition to small amounts of some VOPO4 phases (V 5+ cations) as otii, 13, Y or 8. Moreover VOPO4 pure phases were observed to be active and selective although to a lesser extent. It turned out that the presence of some V 5+ cations on the V 4+ catalyst surface of (VO)2P207 was necessary although an excess was observed to be detrimental. Moreover the catalyst surface was observed to be richer in P than the bulk by a ratio of about 2. The (100) face of (VO)2P207 corresponding to edge sharing dimers of VO6 octahedra bonded to the following chain by PO4 tetrahedra was shown to be the active and selective face [ 11]. One has one oxygen of V = O bond pointing away from the surface and the second one pointing downward in the form of a dimer as schematized in fig. 2. In a recent paper Grasselli et al [12] have proposed a mechanism with activated 02. (peroxo) species on unsaturated V ion at the surface and have proposed that the active site is composed of an ensemble of four dimers isolated one from the other by excess of phosphate species as schematised in fig. 2. Note that if such an assumption is true it is hard to understand how such activated oxygen species may be stable at high temperature and to understand the moving bed technology recently developed by the Dupont researchers [13]. In the later case one can expect surface oxygen species to be consumed and thus the VS+/V4+ ratio value to decrease with the reaction time. However the concept of four dimers as an active site seems reasonable and coherent with the scheme proposed by Ziolkowski [14] and Bordes [15] and with the presence of excess P species on the surface.

66

,' ",

iiiiol 1..i l~176 ',," ,"~',, ',P..' ,", ', .'

6;,

",L'

.",

....

', ,,"

',,~,,' ,,';',, ",,Y .,~,, ,,K, ,,;,, 'X' ,~, %; .,';',,

..".."....'..",,....,..--7.... '.;:;..;-"..,'.:--)'--:~.;-

t::)ol.L'L.,ioi;')oi411.... i i oi ..

iit~

"

I

1oi

,,

,L,

0

o~11..o~

1o

I -- o~V.o~bo

0 s/p'~ 0 o

.---.

o

o..,,.o

0

0

Fig. 2. Schematic representation of the surface structure of one polytype of (VO)2P207. The arrows represent the possible pathways for facile exchange of surface bound oxygen, either monoatomic or diatomic, between the active sites. The <<site-isolation >> due to the diffusion barrier created by the pyrophosphate groups is clearly shown by these arrows (from [ 12]). 5. IRON PHOSPHATES AND HYDROXYPHOSPHATES [16-19] Such catalysts appeared to be potentially important catalysts for the oxidative dehydrogenation of isobutyric acid (IBA) to methacrylic acid (MAA). This reaction is important as a first step to form methyl methacrylate monomer used for plexiglass or altuglass formation by polymerisation. The industrial type catalyst contains iron hydroxyphosphate of uncertain nature and Cs or NH4 as additive and unfortunately necessitates a large amount of water in the feed (namely 10 to 12 mol. H20 per tool. of IBA) to remain stable with time on steam. This makes this process difficult to be develop ,ed industrially. Taking into account the ph ase diagramme Fe203-FeO-P205 it could be possible to select and study several phases wt~ch contain Fe 2 , Fe 3+ or both catio,,s able to insure the redox mechanism necessary for the reaction to take place as it was shown to proceed via a Mars and van Krevelen mechanism [20]. It was also shown that the active phase was a phosphate with iron ion at a two oxidation states +2 and +3, ctFe3(P207)2, composed of trimers of face sharing FeO6 octahedra (as schematized in fig. 3a) separated one from the other by PO4 tetrahedra. The active sites were shown to correspond to a group of two trimers facing each other on their respective layers. The r61e of redox couple and of hydroxylation is shown below:

H20 02 Fe23+Fe2+(P207)2 ~---~ Fe23+Fe 2+ (PO3OH)4 > Fe2+x3+Fel-x2+(PO3OH)4-x(PO4)x 400oC The way the octahedra are connected to each other is an important parameter since it has been shown that the other polymorphic form of Fe3(P207)2 namely J3Fe3(P207)2 was poorly active and selective. In this phase, the central iron atom is in fact occupying a FeO6 trigonal prism (fig. 3b). A study of many different iron hydroxyphosphates exhibiting clusters of face sharing FeO6 octahedra of different sizes has been carried out. It turned out that all catalysts were belonging after catalytic testing to the same solid solution of the type Fe3+4.~Fe2+3x(PO4)3(OH)3-

67

3xO3xwith 0_<x _<1. Depending upon their composition these phases contain clusters of different sizes ranging from dimers (Fe4(PO4)3(OH)3) to continuous chains (J3Fe2(PO4)O) of FeO6 octahedra. The results showed that all samples were active and selective for the reaction whatever the different sizes of the octahedra arrangements but the optimum activity and selectivity were observed for the phase Fe3(PO4)z(OH)2 called barbosalite and containing limited size clusters, namely trimers as schematized in fig. 3a. The added water plays a role as for the active phase of the industrial catalyst in stabilizing the hydroxylated catalysts and the redox couple is the same but implies O/OH instead of PO4/POaOH couple. a

b ;3o

,

Fig. 3. Arrangements of FeO6 octahedra in trimeric clusters isolated one from the other by PzO7 groups: (a) otFe3(P2Ov)z; (b) 13Fe3(P2OT)2; grey circles, FEZ+;black circles, Fe 3+. In a theoretical study using extended Hfickel molecular orbital calculation it has been shown that for iron octahedra assembled as dimers [21], trimers [22] or larger clusters there may occur fast electron exchange between Fe z+ and Fe 3+ cations. This occurs in the case of trimers but neither for dimers nor for trimers in 13Fe3(P207)2 (fig. 4). This fast electron exchange between iron cations which very probably favors the redox mechanism and thus the catalytic properties, should be limited to the closest neighbours to give the most performant catalysts [23 ]. Such clusters of FeO6 octahedra also exist in other inorganic compounds. For instance they exist in ilvaite (CaFe3+Fe2+2Si2OvOOH) where silicate layers replace phosphate anions and FeO6 octahedra form ribbons. It is interesting to note that such a material is active and selective for the reaction although to a lesser extent than the previous hydroxyphosphates [24]. The lower catalytic behaviour may be due to the presence of Ca 2+ cations in the structure and/or to the silicate counter anion whose basicity in the sense of Pearson is different from that of the phosphate anion and/or the infinite size ofFeO6 octahedra clusters. The main idea one has to keep in mind from this study is that inorganic clusters of iron octahedra with iron at two oxidation states are active for the reaction studied and that one has to consider the active sites as these clusters, preferentially as ensemble of two trimers although clusters of other sizes (dimers, tetramers, pentamers...) are also active and selective but to a lesser extent. The iron oxidation state is changing during the reaction in a similar way as the

68 V4+/V5+ redox couple on VPO catalysts (vide supra w This again shows that metallic oxides have to be considered with a dynamical view during the oxidation reaction.

X=2

X=3

X>3

Fig. 4. Electron exchange between iron cations in (FeO6)n octahedra clusters calculated by extended Huckel molecular orbital theory (from [21] and [22]. 6. SUPPORTED OXIDE CATALYSTS Supports are very often used in catalysis for several reasons, namely: i. Dispersion of the active phase in order to increase the surface to volume ratio since heterogeneous catalysis is occuring at the solid surface. ii. Heat transfer: this is a particularly important aspect in oxidation reactions because of the high exothermicity involved. This holds particularly true at industrial scale since the hot spot problem is very crucial and should be monitored with precision within a few degrees to avoid local overactivity and subsequently overheating with presumably irreversible phase transformation of the catalyst. iii. Attrition this aspect is of main importance for fluidized or solid transported beds. One very usually uses silica or carborandum as a binder or as a coating to limitate attrition or to facilitate heat transfer respectively. iv. Formation of new catalytic sites this point will be emphasised below in some examples. One will see that well dispersed species of limited size could be formed and exhibit peculiar catalytic properties. v. Modification of the active phase properties due to its chemical interaction with the support, including epitaxial induced modifications. In some cases the chemical effect of the support will be determining for catalytic properties. For example several oxides (V205, MOO3, RezO7, Cr/O3...) deposited on several oxide supports (SiO2, TiO2, AIzO3) were shown by I. Wachs et al [25] to exhibit very different catalytic properties in methanol oxidation reactions as summarized in table 2.

69

Table 2 Turn over number values in (S)"1 for methanol reaction at 230~ deposited on several supports (from ref. [25]). Oxide support Supported oxide V205 MoO3 CrO3 SiOz 2 39 160 A1203 20 2 Nb205 700 32 58 TiO2 1800 310 300 ZrO2 2300 92 1300 6.1. M o l y b d e n u m

on lwt % metallic oxide

Re207 20 12 1200 170

oxide s u p p o r t e d on silica

Different procedures may be used to prepare such samples as impregnation of silica with a molybdate salt or grafting molybdenum chloride or molybdenum based organo metallic compounds as Mo carbonyls on the hydroxyl groups of silica or at last solid-solid reaction between MoO3 and SiO2 at temperatures near or above 500~ A parameter important for the impregnation method is the pH of the molybdate solution. As a matter of fact the following equilibrium has been well established: MO70246" + 4H20 6-~ 7 M0042 + 8H +. The monomeric tetrahedral MoO42 species is favored at high pH and vice versa for the polymeric heptamolybdate anion. In a study of ammonium heptamolybdate impregnated silica [26], the molybdenum loading was varied up to about 20 wt%. The hydroxyl groups of the silica were observed by infrared and/or UV vis-NIR spectroscopies to decrease with Mo loading and to disappear at 7 wt% Mo loading while two UV bands were appearing at 245 and 340 nm. The former band may be assigned to tetrahedral MoO 24 monomeric species and the latter one to octahedral polymeric (polymolybdate) species. The former band was observed to increase in intensity proportionally to Mo loading at low Mo content and to saturate at roughly 3 wt% Mo. The latter band was observed to appear near 2 wt% Mo, to increase then proportionally to Mo content up to 7 wt% Mo and then to decrease slowly with Mo loading. The formation of the former species was described by: \ / /Si x I

2

-SiI

O

OH

+ MOO]- +2H +

--)

/. O

\ Mo "/ / \ /0 \0 / Si \

+ HzO

Catalytic properties were studied for two reactions namely isopropanol conversion and propene partial oxidation. The first reaction is a test reaction which allows to characterize acidic, basic or redox properties of a catalyst. One gets dehydration to propene or diisopropylether for acid catalyst, acetone for basic catalyst in absence of air and acetone and water for redox type catalyst in presence of air. The experimental results at 100~ clearly show that at low Mo loadings acidic features are favored while redox features are favored at higher loadings. This indicates that monomeric MOO]- species are acidic (presumably as in silicomolybdic acid) while polymeric species exhibit redox properties.

70 The second reaction studied is propene oxidation at 380~ Weak activity was observed for low Mo loading. At higher Mo loading, propanal was the major product while acrolein was also observed. At very high Mo loadings and for MoO3 one gets almost exclusively acrolein. Propanal is known to stem from propene by an electrophilic attack while acrolein corresponds rather to a nucleophilic attack. These results indicate that monomeric MoO~- species does not oxidize propene in our conditions while polymeric (polymolybdate) species exhibit redox properties, with O 2 species being rather electrophilic. Molybdenum oxide exhibits redox properties with 02. species being rather nucleophilic. One can thus realize how the size of the active site is important in oxidation catalysis. This is typical of structure sensitive reaction. 6.2. Vanadium oxide on TiO2 support

Such a catalyst is well known for several reactions, such as o-xylene oxidation to phthalic anhydride and selective catalytic reduction (SCR) of NO by ammonia. The anatase form of TiO2 appears to be better than the rutile form. Such catalysts with 1 and 8 wt% V2Oflanatase was prepared by Rh6ne-Poulenc (S ~ 10 m2 g-l) for an exercise of characterization by 25 different european laboratories. All results are assembled in one issue of Catalysis Today published in May 1994, vol. 20 n~ Surface vanadium species were observed to exist in three different forms: monomeric VO43 species, polymeric vanadate species and V205 crystallites [27], the relative amount of which depended on initial wt % V2OJanatase and on the subsequent selective dissolution treatment. The following conclusions could be drawn considering that a polyvanadate species occupies a ca circular zone of a diameter of 0.38 nm per V atom (i.e. 0.165 nm2) and an isolated monovanadate species 0.66 nm per V atom (i.e. 0.43 nm2): i. Monomeric VO43 species exhibit acidic character with OH groups (Br6nsted acidity) and result in propene formation for isopropanol conversion and in total oxidation for o-xylene oxidation. Maximum coverage equals 0.43 wt% V205. ii. Polyvanadate surface species exhibit redox properties, namely give rise to acetone for isopropanol conversion and phthalic anhydride for o-xylene oxidation. iii V205 crystallites exhibit low activity for o-xylene conversion and high selectivity in total oxidation. Here too the size of the VOx species is of primary importance for the catalytic behaviour. 6.3. Niobium oxide species on silica support [28]

Deposition of NbOx species of different sizes on a silica support has been performed starting from organo metallic complexes as Nb(~I3-C3Hs)4 for a monomer, [Nb(qs-CsHs)H-ta(~5, ~1C5i_h)]2 for a dimer and Nb(OC2Hs)5 for a monolayer. All species were characterized by EXAFS, XANES, FT-I~ Raman, XPS and ESR techniques. Catalytic performances for conversion of ethanol are summarized in table 3. The structure of the species was well characterized particularly the NbO, NbNb, NbSi bonds lengths and coordination numbers. The catalytic data show clearly that dehydrogenation occurs majoritarily on monomeric species and dehydration on monolayer (bidimensional) structure for ethene (intra molecular reaction) and at last on dimeric species for both intra and inter molecular dehydration reaction.

71 Table 3 Catalytic performances of Nb monomeric, dimeric or monolayered species NbOx on SiO2 support for dehydrogenation (---> AA) and dehydration (intra ---> E, inter ~ DE) of ethanol (from ref. 29). Initial rate mmol. min~ g(Nb) ~ Selectivity % Catalyst Total AA E + DE AA E DE 1.25 1.2 0.049 96.1 2.8 1.1 monomer* 0.18 0.004 0.176 dimer* 2.1 24.2 73.7 0.11 0.001 0.106 monolayer** 0.9 99.1 0.0 0.17 0.052 0.118 impregnated* 30.5 20.2 49.3 0.0026 0.0004 0.0022 14.8 46.9 38.3 Nb205 bulk** AA acetaldehyde, E Ethene, DE Diethylether Reaction: Ethanol 3.1 kPa *at 523 K **at 573 K 6.4. Various oxides deposited on different oxide supports The idea is here to determine how a support may modify the catalytic properties of the different oxide species (as inorganic clusters of different size). It has been described above how the size of the deposited oxide species (monomefic, polymeric, bulk-type) results in different catalytic properties. Vanadium oxide has been deposited on several supports as silica, alumina, titania but also zirconia, niobia, zirconium hydroxyphosphates, etc. Its reductibility was studied by reduction with hydrogen at 400~ It was observed that V on TiO2 and yA1203 was reduced rather fast and one reached a consumption such as an O/V atomic ratio near to 1 for TiO2 and 0.65 for A1203 [30]. At variance for silica the reduction was much slower and an O/V ratio of 0.57 was obtained. Such differences were interpreted by the authors as due to different species on the surface: mainly monomefic VO43 species for TiO2, dimefic species V2074"for y.AI203 [30] and V205 crystallites for silica [31 ]. The size of the polyvanadate species in solution corresponds to V3093, V40124"with tetrahedrally coordinated V at pH = 7 and 4.5 respectively and V100286 or VI0028H 5 with octahedrally coordinated V at pH = 2.5. Deposition of such polyvanadates of different sizes was performed carefully on yA1203 support and their initial structures despicted above were shown to remain stable even after calcination in flowing air at 500~ and to be only partly modified after catalytic reaction of oxidative dehydrogenation of propane in the 350 to 450~ temperature range. Moreover in such cases, the selectivity towards propene was observed to be the same at the same conversion level, indicating that in this size domain of polyvanadate species the catalytic selectivity was not changed, only the activity was observed to increase with vanadium loading [32]. In a study of VO 2§ and Cr 3§ cations exchanged or impregnated on the surface of zirconium hydrogenophosphates and compared to vanadyl pyrophosphate and chromium (III) phosphate, the best catalytic properties for ethane to ethylene oxidative dehydrogenation reaction at 500~ were observed when the material present either continuous chains of vanadyl or CrO6 octahedra as in (VO)EP207and et CrPO4 respectively or chains of limited size [33]. By comparison with bulk oxides exhibiting similar environment, it was suggested as for supported VOx and CrOx species on zirconium hydrogenophosphates that the catal~ic performances were related to the stronger basic character ofPO4 3" anions with respect to O anions.

72 7. VMGO CATALYSTS [34-42] Such a system was interesting to consider since it is basically composed of an acidic (V205) and a basic (MgO) oxides. One may then expect either to have well defined phases of given catalytic properties or to have such phases deposited on a basic support (MgO). Three phases are well known namely the orthovanadate Mg3V2Os, the pyrovanadate Mg2V207 and the metavanadate MgV206. The first phase exhibits isolated VO4 tetrahedra (separated by MgO6 octahedra), the second one has comer sharing VO4 tetrahedra and the third one has an octahedral structure. The pyrovanadate phase was found [36] to be the more selective phase at equal conversion with respect to the other two phases for oxidative dehydrogenation of propane at 500~ At variance the Mg3V208 phase (exhibiting isolated VO4 entities separated by octahedral MgO6 entities) was found by H. Kung et al [35] to be more efficient for propene formation in propane oxidative dehydrogenation than the other two phases. A. Pantazidis and C. Mirodatos [37] found recently for the same reaction that propane selectivity was about the same (65 to 75 %) within a large range in V205 content (10 to 80 V205 wt %) and at similar conversion levels. A. Corma et al [36] concluded that tetrahedral isolated VO4 species are more selective for oxidative dehydrogenation of Ca and nC4 alkanes. It then appears than contradictory results have been obtained resulting obviously in contradictory interpretations concerning the role of atom arrangements. The discrepancies between authors may arise from differences between samples, supposedly similar, and from different catalytic reaction conditions particularly the alkane to oxygen ratio values. Taking the samples prepared in our group we have tried to go deeper in the characterization of the catalysts. Reducibly and electrical conductivity measurements [39] have shown that the MgEV207 phase was more easily reduced than the other two while anionic vacancies were created according to: C3H8 q- O2"surf > C3H6 (g) + H20 + [-7 + 2e. Moreover, basicity of surface oxygen ions acts on the ability to create anionic vacancies and thus on the catalytic activity. Moreover if the sample contains more MgO than the above stoichiometries for pure phases one may imagine that its basicity may also play an additional r61e. As a matter of fact propene as a base will be more easily eliminated from the surface of a basic catalyst. The acid-base features of the catalysts were studied by the reaction of isopropanol conversion to propene (acidic feature) and acetone (basic feature) under N2 in the feed and the redox features by the reaction under air in the feed. It was observed at 230~ (table 4 from ref 40) that the pyrovanadate sample was much more basic than the other two pure phases and that excess MgO with respect to crystallized phase stoichiometry induced even more basic character (table 4). As the olefins and to a lesser extent the alkanes are basic one may expect the desorption to be favored by surface basic sites. In other words oxidative dehydrogenation of alkanes is expected to be easier on surface exhibiting basic properties. As a matter of fact the results given in table 5 from ref. 41 show that Mg2V207 which is more basic as shown in table 4 is more selective for olefins in propane conversion and to a lesser extent for n-butane and isobutane oxidation reactions than the other two phases. Such a feature is even more pronounced for the samples with excess MgO at least for propane oxidation, samples which were also shown to present higher basicity (table 4).

73 Table 4 Catalytic data for isopropanol conversion under nitrogen (A) or air (B) in the flow at 230~ alter activation of the sample at 230~ for lh30 under N2 flow, values taken (from ref 40). S t2g"1

Samples

V205 (wt%)

n

MgVzO6 Mg2V207 Mg3V208 MgO 40VMgO" 60VMgO" V205

82.1 70.1 6O.8 0 38 58.5 100

0.1 1.7 0.9 1 J,2 43 19 2.1

Conversion % 7.5 1.5

2.2 0.2 1.2 3.6 64

m= 100 mg; flow rate 18.6 cm 3. min"l 9 Pi= 2.0 x 103 Pa

A Propene %

B

Acetone Conversion Propene % % %

90 43 69 0 14 24 94

10 57 31 100 86 76 6

8

1.7 2.8 0.2 1.7 4.3 4.8

90 40 74 0 17 24 75

Acetone % 10 60 26 100 80 76 25

" mixture ofMgO, Mg3V208 and otMgzV207 as determined by XRD analysis.

Table 5 Catalytic data for oxidative dehydrogenation of propane (A) n-butane (B) and isobutane (C) at 540~ (from ref 41).

Samples MgV206 Mg2V207 Mg3V208 40VMgO 60VMgO

Conversion % 7.4 6.9 8.3 8 8

A Propene" % 15 53 6 65 65

B

COx' % 53 28 94 35 35

Conversion % 10 11

15 15c 42 c

C Dehydr. Conversion i butene % products b 33 14 19 40 9 40 13 10 20 20 23 33 6 13 30

flow rate = 50 cm 3min1, P(C3--)=2 x 103pa, C3 or C4: air = 2:98 "" balance to 100 % corresponds to oxygenates as ethanal, acrolein, propanal and acetic acid. b. dehydrogenates correspond to butenes (major) (but l ene and but 2 ene) and butadiene (minor). c taken at 450~ instead of 540~ because of its high conversion level due to its high surface area. It follows that one may conclude that basic properties as well as atomic arrangements at molecular level play an important r61e in alkane oxidative dehydrogenation reactions. Such a conclusion could also be reached from the study of VMgO catalysts [42] for oxidative dehydrogenation of several alkanes as ethane, propane and butane under similar conditions (see e.g. fig. 5 in ref. 42).

74 8 CONCLUSION Some general conclusions may be drawn from this general presentation: i. Oxidation reactions in gas-phase heterogeneous catalysis usually proceed via Mars and van Krevelen mechanism i.e. involve lattice oxygen ions. Such ions exhibit an electrophilic or a nucleophilic character and therefore present different catalytic properties. As a matter of fact the electrophilic oxygen has been suggested to interact with a double bond or an aromatic ring and the nucleophilic oxygen to interact with a C-H bond in ~ of the double bond or of an aromatic ring. ii. Oxidation reactions are structure sensitive and therefore greatly depend on the local and surface structure of the oxide catalysts. A peculiar fitting between stereochemistry of the solid surface and of the reactant molecule(s) is to be obtained to get the best catalyst. Parameters such as reducibility and reoxidability features of the oxides are very important for catalytic reactions. iii. Active sites for oxidation reactions appear to be molecular "inorganic ensembles" of metallic oxide atoms whose size greatly influences the catalytic properties [2, 3, 43]. In some examples the number of atoms constituting the active sites could be established. For instance double trimers of face sharing FeO6 octahedra are particularly active and selective for oxidative dehydrogenation of isobutyric acid to methacrylic acid on iron hydroxyphosphates; ensembles of four dimers of VO6 octahedra are suggested to be the active sites for butane oxidation to maleic anhydride in vanadium pyrophosphate catalysts. Usually monomeric species as MoO~or VO43" exhibit acidic features and then total oxidation properties. At variance low size polymeric species exhibit better selectivity for many partial oxidation reactions than large size species or bulk-type oxide. It is thus obvious that the size of these r inorganic molecular ensembles )) as active sites is important depending on the reaction considered and the support used. The characterization of such ~ clusters )~ laying on the surface is rather difficult and the above statements for supported oxide catalysts are only qualitative. To better characterize such clusters, techniques as XANES (sensitive to local symmetry and oxidation state of transition metal cations), EXAFS (sensitive to the coordination number and nearby elements distances) and radial electron distribution fRED) of X ray diffraction peaks (sensitive to the nearby elements distances) have to be used. Already data have been obtained for MoO3 and V205 deposited on SiO2, Al203 or TiO2 support and for NbOx species on SiO2 (28, 29). This holds true only if one has one type of cluster species to analyse. The reader interested may read a chapter devoted to XANES and EXAFS techniques by B. Moraweck in the book by B. Imelik et J.C. Vrdrine (1988 and 1994) [44, 45]. iv. Many other examples may have been given. For instance heteropolyoxometallates as H~Mo~VO40 material or zeolite materials as titanosilicalite constitute other materials where the isolated or more or less condensed active species results in very different catalytic behaviour. The case of the titanosilicalite from ENI Co is typical since isolated Ti species are active for phenol oxidation to catechol and anthraquinone while condensed Ti species results in H202 decomposition. v. Oxidation catalysts have to be considered with a dynamical view under reaction conditions. This is related to the Mars and van Krevelen mechanism which involves a redox mechanism and also to the mobility of the oxide lattice. This dynamical phenomenon results in the wetting effect observed under catalytic reaction conditions for multicomponent and supported oxide catalysts [46, 47]. It follows that for many catalysts a certain time on stream is necessary before the catalyst reaches its steady state. It is frequent that in an industrial plant a

75 steady state is reached only after one or two hundreds hours, the catalysts lasting several years before having to be replaced. For simple catalysts as doped vanadyl pyrophosphates used for butane oxidation to maleic anhydride the fight size of the active sites (e.g. tetramers of vanadyl dimers) is monitored by the reactants in catalytic reaction conditions leading to the fight VS+/V4+ ratio, by the preparation procedure to change the material morphology (the (100) face of (VO)2P207 being developed), and by the adequate addition of additive elements which regulate the site size and VS+/V4+ ion ratio. The view of an oxidation catalyst as dynamical under catalytic reaction conditions is essential for our understanding of its functioning. REFERENCES

[1] J.C. Volta, W. Desquesnes, B. Moraweck and G. Coudurier, React. Kinet. Catal. Lett.,12 (1979) 241; J.C. Volta and J.L. Portefaix, Appl. Catal., 18 (1985) 1. [2] J.C. V6drine, J.M.M. Millet and J.C. Volta, Catal. Today, 32 (1996) 115. [3] J.C. V6drine, G. Coudurier and J.M.M. Millet, Catal. Today, 33 (1997) 3. [4] J.M. Tatibouet and J.E. Germain, C.R. Acad. Sci., 290 (1980) 321; J. Catal., 72 (1981) 37. [5] J. Haber, in R.A. Sheldon and R.A. van Santen, Editors, Catalytic oxidation, Principles and Applications, World Scientific, Singapore, 1995, p. 17. [6] J.M. Tatibouet, J.E. Germain and J.C. Volta, J. Catal., 82 (1983) 240. [7] M. Abon, B. Mingot, J. Massardier and J.C. Volta in <<Structure-activity and selectivity relationships in heterogeneous catalysis )), R.K. Grasselli and A.W. Sleight (Ed.), Stud. Surf. Sci. and Catal., Elsevier, Amsterdam, 67 (1991) 67. [8] R.M. Smith and U.S. Ozkan, J. Catal., 141 (1993) 124. [9] G. Centi, Editor, Vanadyl Pyrophosphate Catalyst, Catal. Today, 16(1) (1993). [10] G. Koyano, T. Okuhara and M. Misono, Catal. Lett., 32 (1995) 205; G.J. Hutchings, A. Desmartin Chomel, R. Olier and J.C. Volta, Nature, 368 (1994) 41. [ 11] F. Trifiro and F. Cavani, Chem. Technol., April (1994) 18. [12] P.A. Agashar, L. de Caul and R.K. Grasselli, Catal. Lett., 23 (1994) 339. [ 13] G.S. Patience and P.L. Millo, Stud. Surf. Sci. Catal., 82 (1994) 1. [14] J. Ziolkowski, J. Catal., 100 (1986) 45. [ 15] E. Bordes, Stud. Surf. Sci. Catal., 67 (1991) 21. [16] J.M.M. Millet, J.C. V6drine and G. Hecquet in <
76 [25] I. Wachs, G. Deo, M.A. Vuurman, H. Hu, D.S. Kim and J.M. Jehng, J. Mol. Catal., 82 (1993) 443. [26] T.C. Liu, M. Forissier, G. Coudurier and J.C. V6drine, J. Chem. Soc., Faraday Trans. I, 85 (1989) 1607. [27] J.C. V6drine, Editor, Eurocat oxide, Catal. Today, 20(1) (1994). [28] Y. Iwasawa, Catal. Today 18 (1993) 21-72. [29] Y. Iwasawa in ~( l lth International congress on catalysis-40th anniversary )), J.W. Hightower et al (Ed), Stud. in Surf. Sci. and Catal., ser., Elsevier, Amsterdam, 101 (1996)21. [30] J. Haber, A. Kozlowska and R. Kozlowski, J. Catal., 102 (1986) 52. [31] J. Le Bars, J.C. V6drine, A. Auroux, S. Trautmann and M. Baems, Appl. Catal. A: Gen., 88 (1992) 179. [32] J.G. Eon, R. Olier and J.C. Volta, J. Catal., 145 (1994) 318. [33] M. Loukah, G. Coudurier, J.C. V6drine and M. Ziyad, Microporous Mater., 4 (1995) 345. [34] M.A. Chaar, D. Patel, M.C. Kung and H.H. Kung, J. Catal., 105 (1987) 483. [35] M.A. Chaar, D. Patel and H.H. Kung, J. Catal., 109 (1988) 463. [36] D. Siew Hew Sam, V. Soenen and J.C. Volta, J. Catal., 123 (1990)417. [37] A. Pantazidis and C. Mirodatos in ~(Heterogeneous hydrocarbon oxidation )), S.T. Oyama and B.K. Warren, Ed., ACS Sympos., ser., Washington, 638 (1996) 207. [38] A. Corma, J.M. Lopez Nieto, N. Paredes, A. Dejoz and I. Vazquez, in (( New Developments in Selective Oxidation II )), V. Cortes Corberan and S. Vic Bellon (Ed.), Stud. in Surf. Sci. and Catal. ser., Elsevier, Amsterdam 82 (1994) 113. [39] A. Guerrero Ruiz, I. Rodriguez Ramos, J.L.G. Freno, V. Soenen, J.M. Herrmann and J.C. Volta in ~(New Develpments in Selective Oxidation by Heterogeneous Catalysis ~), P. Ruiz and B. Delmon (Ed.), Stud. in Surf. Sci. and Catal. ser., Elsevier, Amsterdam, 72 (1992) 203. [40] A. Ouquour, Doctoral Thesis N ~ 105-91, University C. Bernard, Lyon, 1991. [41] V. Soenen Lebeau, Doctoral Thesis N ~ 173-91, University C. Bernard, Lyon, 1991. [42] P. Concepci6n, A. Galli, J.M. L6pez Nieto, A. Dejoz and M.I. Vazquez, Topics in Catalysis 3 (1996) 451. [43] J.C. V6drine in ~( Catalytic Oxidation, Principles and Applications )~, R.A. Sheldon and R.A. van Santen (Ed.) World Scientific, Singapore, 1995, pp. 53-78. [44] B. Moraweck in ~(Les Techniques Physiques d'Etude des Catalyseurs )~, B. Imelik et J.C. V6drine (Ed.), Technip, Paris, 1988, pp. 587-598. [45] B. Moraweck in ~ Catalyst Characterization, Physical Techniques for Solid Materials )), B. Imelik and J.C. V6drine (Ed.), Plenum, New York, 1994, pp. 377-416. [46] J.M.M. Millet, H. Ponceblanc, G. Coudurier, J.M. Herrmann and J.C. V6drine, J. Catal. 142(1993) 381. [47] H. Ponceblanc, J.M.M. Millet, G. Coudurier and J.C. V6drine in ~ Catalytic Selective Oxidation )), S.T. Oyama and J.W. Hightower (Ed.), ACS Sympos. ser., Washington, 523 (1993) 262.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

In situ e l e c t r o c h e m i c a l l y oxidation catalysts

controlled promotion

77

of c o m p l e t e

and p a r t i a l

Constantinos G. Vayenas and Symeon I. Bebelis Department of Chemical Engineering University of Patras GR-26500 Patras, Greece

Oxidation catalysis on metal catalysts can be affected significantly by electrochemical means. The catalytic activity and selectivity of metal films interfaced with solid electrolytes can be varied in situ in a very pronounced and reversible manner via electrical potential application (typically +1-2 V) between the catalyst film and a counter electrode also in contact with the electrolyte. The effect is particularly pronounced for complete and partial oxidation reactions. Catalytic rates can be varied by up to a factor of 200. The induced steady-state change in catalytic rate is up to five orders of magnitude larger than the steady-state rate of electrochemically supplied ionic species from the electrolyte onto the catalyst surface. This novel effect of Electrochemical Promotion or non-Faradaic electrochemical modification of catalytic activity (NEMCA effect) has been studied in over 45 catalytic reactions on Pt, Rh, Pd, Ag, Ni and IrO 2 catalyst-electrodes using a variety of solid electrolytes (O 2-, F-, Na § H § conductors) and more recently mixed ionic-electronic conductors and aqueous alkaline solutions. In addition to possible technological applications, this new effect allows for a systematic study of the role of promoters in heterogeneous catalysis. In this paper the main features of electrochemical promotion are summarized with emphasis on complete and partial oxidation reactions and the origin of the effect is discussed on the basis of work function measurements, recent surface spectroscopic investigations and ab initio quantum mechanical studies. 1. INFRODUCTION Controlled modification of the activity and selectivity of complete and partial oxidation reactions on metals is a long-sought goal in oxidation catalysis, [ 1]. Metal alloying [2], metalsupport interactions [3] and the use of promoters [4] have been investigated extensively for this purpose. The use of electrochemistry for this purpose, i.e. to activate and precisely tune heterogeneous catalytic processes is a new development [5-15] of considerable theoretical and practical interest [ 16-18]. The electrochemically induced catalytic rate enhancement can be several orders of magnitude higher than that anticipated from Faraday's Law [5-15]. Furthermore the product selectivity of catalytic processes can be affected in a pronounced and reversible manner [9,10,14,15]. Most studies in this area have focused on oxidation reactions using solid electrolyte cells [5-15]. Recent work, however, has shown that similar phenomena occur in aqueous electrochemistry [ 19]. The experimental setup is quite simple:

78 The gaseous reactants (e.g. C2H4 plus 02) are cofed over the working electrode of a solid electrolyte cell: Catalyst working ] solid electrolyte [ counter auxiliary gaseous reactants electrode [(e.g. ZrO2-Y203) electrode gas (e.g.C2H4+O2) (e.g. Pt,Rh,Ag,IrO2) (e.g. Au) (e.g. 02) The working electrode serves simultaneously as an electrode and as the catalyst for the catalytic reaction under study, e.g., C2H 4 or H 2 oxidation. The auxiliary gas can be the reactive gas mixture itself in the, so called, "single-pellet" design [9,10]. Upon varying the potential of the working electrode/catalyst it is found that not only the electrocatalytic (net charge-transfer) reaction rate is affected, as anticipated from Faraday's Law, but also the catalytic (no net-charge transfer) reaction rate changes in a very pronounced, controlled and reversible manner. The increase in the catalytic rate can be up to a factor of 100 higher than the open-circuit catalytic rate and up to 3x105 times larger than the change in the electrocatalytic rate, e.g., each 02- supplied to the catalyst electrode can cause the catalytic reaction of up to 3x105 chemisorbed oxygen atoms [9,14]. This novel effect has been termed non-Faradaic electrochemical modification of catalytic activity (NEMCA effect [5-15])or electrochemical promotion [16] or in situ controlled promotion [20]. Its importance in catalysis and electrochemistry has been discussed by Haber [ 18], Pritchard [16] and Bockris [17], respectively. In addition to the group which first reported this new phenomenon [5-7], the groups of Lambert [12], Hailer [10], Sobyanin [8], Comninellis [ 13], Pacchioni [21] and Stoukides [11] have also made important contributions in this area, which has been reviewed recently [14,15]. In this review the main phenomenological features of NEMCA for oxidation reactions are briefly surveyed and the origin of the effect is discussed in the light of recent kinetic, surface spectroscopic and quantum mechanical investigations. 2. EXPERIMENTAL The experimental setup for kinetic electrochemical promotion studies is shown schematically in Fig. 1a. The electrically conductive working catalyst electrode, usually in the form of a porous film 3-20 gm in thickness with a roughness factor 3 to 500 [9,14,15] is deposited on the surface of a ceramic solid electrolyte (e.g. Y203-stabilized-ZrO 2 (YSZ), an O 2- conductor, Na-13"-AI203, a Na § conductor, CaZr0.9In0.103-a, a H § conductor [22], or TiO 2, a mixed electronic-ionic conductor [23]). Catalyst, counter and reference electrode preparation and characterization details have been presented in detail elsewhere [9,14] together with the gas analysis system for on-line monitoring of the rates of catalytic reactions via gas chromatography, mass spectrometry and infrared spectroscopy. The superficial surface area of the metal working electrode-catalyst is typically 1-2 cm 2 as measured via reactive titration of oxygen with CO or C2H4 [9,14] or via reactive titration of CO with O2 [9,14]. The catalyst-electrode is exposed to the reactive gas mixture (e.g. C2H4+O2) in a continuous flow gradientless reactor (CSTR). The counter and reference electrodes are usually exposed to air when using the "fuel cell" type design and to the reactive gas mixture itself when using the "single-pellet" type design. In the latter case the counter and

79 reference electrodes must be catalytically inert (e.g. Au) while the reference electrode is only a monitoring (pseudoreference) electrode [24]. A galvanostat or potentiostat is used to apply constant currents between the catalyst and counter electrode or constant potentials, VWR, between the catalyst-working electrode (W) and the reference (R) electrode. In this way ions (02- in the case of YSZ, Na + in the case of Na-I3"-AI20 3, H + in the case of CaZr0.9In0.103_ ~) are supplied from (or to) the solid electrolyte to (or from) the catalyst electrode surface. The current, I, is defined positive when anions are supplied to or cations removed from the catalyst surface. As shown under Results and Discussion there is concrete evidence obtained from several techniques, including work function measurements, cyclic voltammetry, temperature programmed desorption (TPD), Xray photoelectron-spectroscopy (XPS) and scanning tunneling microscopy (STM) that these ions (accompanied by their compensating (screening) charge in the metal thus forming surface dipoles) migrate (back-spillover in catalytic terminology) onto the gas-exposed, i.e. catalytically active, catalyst electrode surface. Consequently the electrolyte acts as an electrically activated catalyst support and establishes an "effective electrochemical double layer" on the gas-exposed, i.e. catalytically active, electrode surface. The experimental setups and procedures for using XPS [25] (Fig. l b), TPD [26] and cyclic voltammetry [27] under ultra-high-vacuum (uhv) conditions and work function measurements [7,27], cyclic voltammetry [26,28] and STM [29] under atmospheric pressure conditions to investigate the origin of electrochemical promotion, have been described in detail recently [25-29]. X-Ray Source

/

~=,o

cou

II

ER

Phoroel ec trQn Energy Analyzer

I

CATALYST

I

??

R~FE~RENCE

~~-~7 --J

v~" cE Q

,

Figure 1. Schematic of the experimental setup for electrochemical promotion studies using the fuel-cell type design (a) and for using x-ray photoelectron spectroscopy (XPS) (b) to investigate the catalyst-electrode surface; G-P: Galvanostat-Potentiostat; WE: Working electrode, RE: Reference electrode, CE: Counter Electrode (adapted from refs. [6], [25]). 3. RESULTS AND DISCUSSION 3.1. Catalytic oxidation rate modification A typical transient NEMCA experiment, carried out in the setup depicted in Fig. la, is

80 shown in Fig. 2. The solid electrolyte is YSZ. The catalytic reaction is the complete oxidation of C2H 4 on Pt [6]. The Pt catalyst film has a gas exposed surface area corresponding to N=4.2x10 .9 mol Pt and is exposed to Po2--4.6 kPa, PC2H4=0.36 kPa, T=370~

in the

gradientless continuous flow CSTR reactor of Fig. la. Initially (t<0) the electrical circuit is open (I=0) and the open-circuit catalytic rate, r o, is 1.5xl 0 .8 mol O/s. The corresponding turnover frequency (TOF), i.e., oxygen atoms reacting with C2H4 per surface Pt site per s is 3.57 s-1. At t=0 a galvanostat is used to apply a constant current between the catalyst and the counter electrode. In this way oxygen ions, O 2-, are supplied to the catalyst-gas-solid electrolyte three-phase boundaries (tpb) at a rate I/2F=5.2xl 0 -12 mol O/s. The catalytic rate starts increasing (Fig. 2) and within 25 min gradually reaches a value r=4.0-10 -7 mol O/s which is 26 times larger than r o. The new TOF is 95.2 s -1. The increase in catalytic rate Ar=rro=3.85xl 0 -7 mol O/s is 74,000 times larger than I/2F, which is the maximum rate increase anticipated from Faraday's Law. Thus each 0 2- ion supplied to the Pt catalyst causes at steady-state 74,000 oxygen atoms chemisorbed on the Pt surface to react with C2H4 and form CO 2 and H20. For this reason this novel effect has been termed non-Faradaic electrochemical modification of catalytic activity (NEMCA effect).

50 t i = 0

30--

0!

~-'-"I = + l ~ t A

40-

C)

20--

I=O ;t... 800 r 0 = 1.5xlO s mol O / s 'i Ar = 38.5x10 -s mol O / s I / 2 F = 5.2x10 -12 mol O / s

~ o~, =26 0

-~

30

---" ~

I ~-

_...

74000 -

-4oo~

Ji

~2

J i

10--

-o

lO-

0 "-I

I

0

I 0

"--T-20 Time, min

40

I:, 110

1 130

Figure 2. Electrochemical promotion: Catalytic rate and catalyst potential response to step changes in applied current during C2H 4 oxidation on Pt [6]. T=370~ Po2=4.6 kPa, PC2H4=0.36 kPa. The experimental (x) and computed (2FN/I) rate relaxation time constants are shown on the figure. The steady-state rate increase Ar is 74,000 times higher than the steady-state rate of supply of 0 2- to the catalyst-electrode (A=74,000); (adapted from ref. [6]).

81 The Faradaic efficiency, or enhancement factor, A, is defined [6,9] from: A =Ar/(I/2F)

(1)

In the experiment of Fig. 2 the maximum A value is 74,000. A reaction exhibits the NEMCA effect when IAI > 1. Depending on the observed sign of A, catalytic reactions are termed electrophobic (A> 1) or electrophilic (A<-I); A values ranging from 3xl 05 [6,9,14,15] and down to -5•

[9,14,15] have been measured. Relatively safe predictions about the

order of magnitude of A can be made as discussed below. A second important parameter is the rate enhancement ratio, p, defined [6,9] from p =r/ro

(2)

In the experiment of Fig. 2, the maximum p value is 26; p values up to 100 [30] or even higher [ 12] and down to zero [31-33] have been obtained. The NEMCA rate relaxation time constant, x, is defined [9,14] as the time required for the catalytic rate increase to reach 63% of its final steady-state value in galvanostatic transient experiments, such as the one depicted in Fig. 2. As shown in this Figure, "c is of the order of 2FN/I. This is a general observation in electrochemical promotion studies utilizing YSZ: x = 2FN/I

(3)

The parameter 2FN/I equals the time required to form a monolayer of O on a surface with N adsorption sites, when O is supplied as 02. at a rate I/2F as is the case here. This provided the first, kinetic, evidence that NEMCA is due to the electrochemically controlled migration (backspillover) of an oxidic species from the solid electrolyte onto the gas-exposed catalytically active catalyst surface [6,9]. As shown below via XPS, TPD and cyclic voltammetry this electrochemically supplied oxygen forms a new strongly bonded ionic adsorption state on the Pt surface and is far less reactive than normally chemisorbed oxygen formed via gas phase adsorption. It severely modifies (increases) the catalyst surface work function [7,27], forces normally chemisorbed oxygen into a weakly bonded adsorption state [26] and acts as a sacrificial promoter by reacting with C2H4 (or other oxidizable molecules) at a rate which is A times slower than the weakly bonded atomic oxygen [26,30]. Figure 2 also shows this point: At steady-state the rate, rc, of consumption of the promoting 0 2- species via reaction with C2H4, has to equal its rate of formation I/2F. Consequently, since A=Ar/(I/2F) and Ar=r, it follows A=r/rc=TOF/TOF c where TOF is the turnover frequency of the catalytic reaction in the NEMCA-promoted state and TOF c is the turnover frequency of the reaction of the promoting oxygen species with ethylene. It thus follows for the experiment of Fig. 2 that TOFc=TOF/A= 1.3• 10 -3 s. This implies that that average lifetime of the promoting species on the catalyst surface is TOFcl=770 s in excellent qualitative agreement with the catalytic rate relaxation time constant upon current interruption (Fig. 2). This observation provides strong support for the oxygen backspillover mechanism of electrochemical promotion.

82

Q

100 -

A

9 9

"

" 250

T

ate

80

200

o

T=370~

," 60 ,~ c "" 40 "

D

Regular Open Circuit Rate

20

150

er ~ r.r.,

100

~ ol> g: ~" ~

PC2H4=0"65kPa

00_ ~..,~.-c~c_,5 ' r I0~ po 2/Pc2H4

n

15' o,

50

2u~0

Figure 3. Effect of gaseous composition on the open-circuit (unpromoted) catalytic rate of C2H4 oxidation on Pt/YSZ and on the electrochemically promoted catalytic rate when the Pt catalyst film is maintained at VWR = 1V (adapted from ref. [6]).

Figure 3 shows the steady-state effect of constant positive potential application VWR=(=+ 1V) on the rate of C2H4 oxidation on P t ~ S Z as a function of the PO2fPC2H4 ratio. The promotion is much more pronounced under oxidizing gaseous compositions, where a sixty-fold enhancement in catalytic rate and turnover frequency is obtained. The stronger electrochemical enhancement under oxidizing conditions is a general observation in NEMCA studies of oxidation reactions [ 14,15] and is due to the fact that the promoting ionic oxygen species forms, upon 02- supply, only after the coverage of normally adsorbed oxygen is near completion [25,26]. An example of using Na-lY'-AI203, a Na § conductor, as the solid electrolyte to induce NEMCA is shown in Fig. 4. The figure depicts the effect of catalyst potential VWR,

,v,

2O

s

Figure 4. Electrochemical promotion of Pt for CO

t.....

"

"

oxidation using Na-13"-AleO3: Effect of P c o , catalyst potential and corresponding linearized sodium coverage on the rate of CO oxidation at

-0.4

~. "'Q'

---.v.. ~ o.oz -"'i 0

5

T = 3 5 0 ~ and Po2=6 kPa (reprinted with permission from ref. [20]).

83 corresponding coulometrically measured [20,32] Na coverage ONa, and Pco on the rate of CO oxidation on Pt/Na-[Y'-A1203 at 350~ [20]. When the Pt surface is Na-free, the rate goes through a sharp maximum with respect to Pco, indicative of competitive adsorption of CO and O (Langmuir-Hinshelwood type kinetics). For high Pco values the surface coverage of O is very low and thus the catalytic rate is also low. Increasing Na coverage (via negative current or potential application) under these conditions causes a 6-fold enhancement in the rate of CO oxidation due to Na-assisted enhanced O chemisorption. The promotion index of Na, PNa, defined from: PNa=(Ar/ro)/AONa

(4)

takes values up to 200 under these conditions. Higher Na coverages (Fig. 4) poison the rate due to the formation of surface Na-CO complexes [20].

3.2. Selectivity modification: Ethylene epoxidation Electrochemical promotion can be used to modify significantly the product selectivity, of catalytic oxidation reactions. An example is presented in Fig. 5 which shows the effect of catalyst potential and corresponding work function change on the selectivity to ethylene oxide (Fig. 5a) and acetaldehyde (Fig. 5b) of ethylene oxidation on A g ~ S Z at various levels of gas phase chlorinated hydrocarbon moderators [31] (The third, undesirable, product is CO2). As shown in the Figure a 500 mV decrease in catalyst potential causes the Ag surface to change from selective (up to 70%) ethylene oxide production to selective (up~to 55%) acetaldehyde production. The same study [31 ] has shown that the total rate of ethylene oxidation varies by a factor of 200 upon varying the catalyst potential. (Fig. 6)

80

A(e@), eV

-0.4

,

,

-0.z

,

',-

0.0

,

,

0.z

,-

A(e@), eV 8o,-q.4

,

-q.~.

o.o

0,.2

l

8.5% 0 2, 7.8%C , H 4 o O'x T=270~ P=500~cPa 50 ~ ~ k~ ~ 0.0 ppm C2t-I,C12 [~ ~, 9 0.4ppm L ~ \ [] 0.8 ppm 40j" / % 9 1.2 ppm I ~ x~lm 1.6 ppm

5O 9 r~

30

~ ~ , . / ~

I0

/~

o600

~-4oo

T__'52~/0Ooc~8_~

iiii m

a

(b)

z0

A 1.2 ppm [] 1.6 ppm @ 2.0 ppm

-~oo

V~ , mV

6

-~

",s

u

-

-z~o, mY

'

a

Figure 5. Effect of catalyst potential and work function and of the gas-phase addition of various levels of 1,2-C2H4C12 on the selectivity to ethylene oxide (a) and acetaldehyde (b) of ethylene oxidation on Ag supported on YSZ (reprinted with permission from ref. [31]).

84

-0.4

-0.2

0.0

0.2

[

10 1

m

0

M

o G

1-

O

I

7

T=260~C, P = 5 O O k P a

//

a.5~

l~

0.1

~/-/ ~w/" ,'~

-700

o~, 9.8~

9 CO,_

-500

~

I

t

-300 V~ , rnV

I

C,_H,

* CJI40 [] C H a C H O

I

-0.1

I

-i00

T

0.01 !00

Figure 6. Effect of Ag/TSZ catalyst potential and work function on the rates of formation of ethylene oxide (A), acetaldehyde (~) and CO2 (O); T = 260 ~ ; P=500 kPa, Po2 = 17.5 kPa; PCzH4=49 kPa.(reprinted with permission from ref. [31])

Figure 7 refers to ethylene epoxidation on a Ag film deposited on [Y'-A1203 [24] and shows the effect of catalyst potential, VWR, and partial pressure of gas phase chlorinated hydrocarbon moderator on the selectivity to ethylene oxide. For VWR--0 and VWR=-0.4 V the Na coverage on Ag is nil and 0.04, respectively. As shown in the Figure, there is an optimal combination of VWR(0Na) and PCzH4CI 2 leading to a selectivity of ethylene oxide of 88%. This is one of the highest selectivity values reported for the epoxidation of ethylene. Figure 7 provides an example of how in situ controlled promotion can be used for a systematic investigation of the role of promoters in technologically important partial oxidation reactions. Sm~x=88~

90ES

r~ 7S

70

t0 ~s J'(""~ 1"J~~~ PP'~ G ~ ~.a ~ . . r "

~c/2

Figure 7. Effect of catalyst potential and gas-phase 1,2C2H4C12 partial pressure on _ / ~ , ~ 3 the selectivity of ethylene / ' ~ epoxidation on Ag/]Y'-A1203. ~ ~ ~ (reprinted with permission ~:~ ~t ,~,~' from ref. [24])

85 3.3. Aqueous electrolyte NEMCA: An example of NEMCA in an aqueous alkaline solution is shown in Fig. 8. The catalytic reaction is the oxidation of H 2 by gaseous 0 2 on finely dispersed Pt supported on graphite [19]. Similar results have been obtained with Pt black supported on a Teflon frit [34]. The figure shows the effect of positive and then negative current application on the rates of consumption of hydrogen (rH2) and oxygen (ro=2ro2) and on the catalyst-electrode potential. It can be shown easily that at steady-state rH2-ro=I/2F [19,34], as also confirmed by experiment (Fig. 8). The induced increase in the catalytic rate is clearly non-Faradaic with A values up to 7.2 and 9 values up to 3.5. Similarly to the case of solid electrolyte induced NEMCA, a detailed kinetic study [34] has shown that increasing catalyst potential weakens the Pt=O chemisorptive bond and weakens the Pt-H chemisorptive bond. This is due to the electron donor and electron acceptor character of chemisorbed hydrogen and oxygen, respectively. !t

Figure 8. Electrochemical

tO

t ~

-

-

~

~

t

promotion of H2 oxidation on Pt in contact with 0.1 M KOH solution ] [19]. Transient effect of .8 applied positive and negative current on the ~ rates of consumption of 4 _~ H2 (rH2) and ~ ( r o = 2 r 0 2 ) and on Pt

!

9 8 7

,, . . . . 4

"

' 6

/--'4,, / - - '~ ~ ' ~ - - - - - - . 4

,'

t,~ ~

~

":~"23y/I(2F ) ~.[ 0 -t L !

I

" "~"'-

t.2

.... t / . ' ~ O " %

1

~ - .....

[r t t , l i ~ , ,. , I . : , ~ : i't~ . . . .

0 ~ I 2 5 7 9 11 1.3 15 25 2FN/7 t, rain

electrode;Vs.Catalyst_electrodea reversiblepH2 =0.75potentialkpa,Ha

_ rt'0.4

20

P~176 total gas flowrate 280 cm3min -1 at STP(adapted from ref. I19]).

3.4. Electrochemical kinetics and the magnitude of A: Table 1 provides a list of the catalytic oxidation reactions studied so far from the view point of electrochemical promotion and of the measured A, P and Pi values. As shown in this table measured IAI values range from 1 to 3x10 $. It has been shown both theoretically [9] and experimentally [9,14] that the order of magnitude of the absolute value IAI of the Faradaic efficiency A can be estimated for any reaction, catalyst and solid electrolyte from the following approximate expression:

= ro/(I0/2F)

(5)

where r o is the open-circuit catalytic rate and I0 is the exchange current of the catalyst-solid electrolyte interface. The latter can be obtained from standard current-overpotential (Tafel)

86 plots [9,14]. Equation (5) indicates that in order to obtain a non-Faradaic rate enhancement (IAI>I) it is necessary for the intrinsic catalytic rate, r o, be higher than the intrinsic electrocatalytic rate I0/2F.

Table 1 Catalytic oxidation reactions investigated in electrochemical promotion studies (Adapted from Table 3 in ref. [ 14] which provides specific references to each reaction) I. Electrophobic reactions ~/~e~)>0; ~r/bVwR>0; ~r/~I>0; A>0 Reactants

Products

Cata- Electrolyte lyst (Promoting ion)

C2H4, 02 C2H4, 02 C2H4, 02 C2H4, O2 C2H6, 02 C3H6, O2 CH4, O2 CH4, O2 CH4, 02 CO, 02 CO, 02 CO, 02 CH3OH, 02 C2H4, 02 CO,O2 H2, 02 H2, 02 CO, 02 C2H4, 02

CO2 Pt CO2 Rh CO2 IrO2 C2H40, CO2 Ag CO2 Pt C3H60, CO2 Ag CO2 Pt CO2 Pt CO2,C2H4,C2H6 Ag CO2 Pt CO2 Pd CO2 Ag H2CO, CO2 Pt CO2 Pt CO2 Pt H20 Pt H20 Pt CO2 Pt CO2 Pt

YSZ (O2-) YSZ (O2-) YSZ (Oz-) YSZ (O2-) YSZ (02-) YSZ (O2-) YSZ (O2-) YSZ (O2-) YSZ (02-) YSZ (02-) YSZ (02-) YSZ (02-) YSZ (02-) ~"-A1203(Na+) 13"-A1203(Na+) Nafion (H§ KOH-H20 (OH-) CaF2(F-) T102(T1Ox,O-) 9

9

+

2

T(oC)

A

P

Pi

260-450 250-400 350-400 320-470 270-500 320-420 600-750 590 650-850 300-550 400-550 350-450 300-500 180-300 300-450 25 25-50 500-700 450-600

3.105 5.104 200 300 300 300 5 50 5 2.103 103 20 104 5.104 105 20 20 200 5.103

55 90 6 30* 20 2* 70 3 30* 3 2 5 4* 0.25 0.3 6 6 2.5 20

55 90 5 30 20 1 7 3 30 2 1 4 3 -30 -30 5 5 1.5 20

-100 -3-103 -500 -60 -104 -3 -1.2 -105

7 6 6 3 15" 3 8* 8

250

II. Electrophilic reactions ~9r/3(e~)<0; ~r/~VwR<0; ~/~)I<0; A<0 C2H6, 02 C3H6, 02 CO, 02 CO, 02 CH3OH, 02 CH4, 02 CH4, 02 CO, 02

CO2 Pt CO2 Pt CO2 Pt CO2 Au H2CO, CO2 Pt CO2 Au C2Hn,C2H6,CO2 Ag CO2 Pt

YSZ (02-) YSZ (02-) YSZ (02-) YSZ (02-) YSZ (02-) YSZ (02-) YSZ (02-) 13"-A1203(Na+)

* Promotion-induced change in product selectivity

270-500 350-480 300-550 450-600 300-550 700-750 700-750 300-450

87 3.5. Work function modification An important step in the understanding of the origin of NEMCA was the realization that solid electrolyte cells with metal electrodes are both work function probes and work function controllers for the gas-exposed surfaces of their electrodes [7,27]. Both theory [9,14] and experiments via the Kelvin probe (vibrating capacitor) technique [7,27] and more recently via UPS [35] have shown that: eV~

= e~-e~R

(6)

eAVwR =A(e(I~v)

(7)

and

where eC~w is the catalyst surface work function and e(I)R is the work function of the reference electrode surface. Equation (6) provides an interesting additional physical meaning to the EMF of solid electrolyte cells, in addition to its well-known Nernstian meaning [9,14]. Equation (7) is equally important as it shows that the work function of the gas-exposed, i.e., catalytically active, surface of solid electrolyte cell electrodes can be varied and controlled via current or potential application. Positive currents increase e~ and negative currents decrease it. Physically this variation is brought about at the molecular level by the spillover of ions from or to the electrolyte to or from the gas-exposed catalyst surface [ 14]. 3.6. XPS studies The use of X-ray photoelectron spectroscopy (XPS) has provided conclusive evidence that spillover/backspillover phenomena are real and that electr, ochemically controlled backspillover of oxide ions, O ~, is the origin of electrochemical promotion. The first XPS investigation of Ag electrodes on YSZ under electrochemical 02- pumping conditions was published in 1983 [36] and provided strong evidence for the creation of backspillover oxide ions on Ag (O l s at 529.2 eV) upon positive current application. These results were confirmed by G6pel and coworkers who used XPS, UPS and EELS to study Ag/YSZ catalyst surfaces under electrochemical bias conditions [35]. A similar detailed XPS study of Pt films interfaced with YSZ [25] has shown conclusively that: I. Backspillover oxide ions (Ols at 528.8 eV) are generated on the gas-exposed Pt electrode surface upon positive current application (peak ~5in Fig. 9 top). II. Normally chemisorbed atomic oxygen (Ols at 530. 2 eV) also forms upon positive current application (peak 7 in Fig. 9 top). The maximum coverages of the ], and ~5states are comparable and of the order of 0.5 each. III. Oxidic backspillover oxygen (5-state) is significantly less reactive with the reducing (H 2 and CO) ultra high vacuum background than normally chemisorbed atomic oxygen. These observations provide a direct explanation of electrochemical promotion when using O2--conducting solid electrolytes [25]. The use of XPS has also confirmed recently that electrochemically controlled Na backspillover is the origin of electrochemical promotion when using Na+-conducting solid electrolytes such as ~"-A1203 [12,37,38].

88

534

Eb , eV 530 528

532

I

I

1

526

t

653K 0.1 kPa

-

1

--

~-5

800

_ oo

,/

.-/0 -/5

52/,

I

-

6

!

-7OO

t

I

i

i

~

"400

!

0

~R,

mV

;o

4O lO0 2OO SO0 4O0 600

8OO I ,

400

Figure 9. XPS (top) and linear potential sweep voltammographic (bottom) investigation of oxygen adsorbed on Pt films deposited on YSZ following positive overpotential application. Top: O 1s photoelectron spectrum of oxygen adsorbed on a Pt electrode supported on YSZ under UHV conditions after applying a constant overpotential AVwR=1.2 V, corresponding to a steady-state current I=40~tA for 15 min at 673 K (ref. [25]). The same O ls spectrum is maintained after turning off the potentiostat and rapidly cooling to 400K (ref. [25]). The ),-state is normally chemisorbed atomic oxygen (Eb=530.2 eV) and the 8state is backspillover oxidic oxygen (Eb=528.8 eV). Bottom: Linear potential sweep voltammogram obtained at T=653 K and Po2=0.1 kPa on a Pt electrode supported on YSZ showing the effect of holding time t H at VWR=300 mV on the reduction of the ),- and ~5 -states of adsorbed oxygen; sweep rate: 30 mV/s (ref. [39]).

3.7. Cyclic voltammetric studies The two types of electrochemically formed chemisorbed oxygen on Pt films interfaced with YSZ are also clearly manifest via solid state linear potential sweep voltammetry (Fig. 9 bottom, Ref. [39]): The first oxygen\ reduction peak corresponds to normally chemisorbed oxygen (y-state) and the second reduction peak which appears only after prolonged positive current application [39] corresponds to the 8-state of oxygen, i.e. backspillover oxidic oxygen, which is significantly less reactive than the ),-state. 3.8. Temperature p r o g r a m n ~ desorpfion The creation of two types of chemisorbed oxygen on Pt surfaces interfaced with YSZ and subject to NEMCA conditions is also manifest clearly by temperature-programmed-desorption (TPD) [26] as shown in Fig. 10. The strongly bonded backspillover oxygen species (peak desorption temperature Tp=750-780 K) displaces the normal chemisorption state of atomic oxygen obtained via gas phase adsorption (Tp-740 K) to a significantly more weakly bonded state (Tp=680 K). The pronounced rate enhancement in NEMCA studies of catalytic

89 oxidations with positive potentials (electrophobic behaviour) is due to the very fast oxidative action of this weakly bonded oxygen. The strongly bonded backspillover anionic oxygen is significantly (A times) less reactive and acts as a sacrificial promoter. 20 Pt-WE ... Au-CE

f~

16

~12

E

Z

8

0 50O

600

700 T,K

800

900

Figure 10. Oxygen TPD spectra after gaseous oxygen adsorption at 673 K and Po2=4.x10 -6 Ton" for 1800 s (7.2 kL) followed by electrochemical 0 2- supply (I=151.tA) for various time periods t/s comparable to 2FN/I (2570 s). Gaseous oxygen supply creates a single adsorption state (Tp=740K), but additional electrochemical oxygen supply creates two adsorption states. The weakly bonded state is highly reactive [26], while the strongly bonded (backspillover) state (Tp=750-780 K) acts as a sacrificial promoter for catalytic oxidations [26]. (Adapted from ref. [26]).

3.9. Scanning tunneling microscopy (STM) The very clear demonstration of ion backspillover as the cause of NEMCA when using Na-~"-A1203 as the solid electrolyte was recently obtained via atmospheric pressure scanning tunneling microscopy (STM) [29]. APt monocrystal with an exposed Pt(111) surface was interfaced with a Na-~"-A1203 component using a Pt paste electrode along the perimeter [29]. Negative current application was found to cause Na backspillover on the Pt(111) surface forming at low surface coverage (<0.01) a Pt(111)-(12xl2)-Na adlattice on the previously existing Pt(111)-(2x2)-O adlattice. Positive current application was found to totally remove the Na adlattice leaving the Pt(111)-(2x2)-O adlattice intact [29].

90 This study, in addition to explaining NEMCA with Na+-conducting solid electrolytes, provided the first STM confirmation of spillover/backspillover phenomena.

3.10. Ab initio quantum-mechanical calculations A very frequent feature of electrochemical promotion studies is the observed linear variation of catalytic activation energies with varying catalyst work function [9,14,15]. It had been proposed that this is due to a linear variation in chemisorptive bond strengths with catalyst work function [9,14], a proposition recently supported by TPD studies for oxygen chemisorption on Pt/YSZ [26]. A recent first-principles investigation of electrochemical promotion using cluster models of oxygen adsorbed on Cu and Pt metal surfaces with coadsorbed positive and negative ions or point charges has confirmed the experimentally observed linear dependence of the O bond strength on metal work function [21 ]. This linear relationship was also found at the first-order perturbation theory level by taking into account only the purely electrostatic interaction between the field induced by the ions and the polar metal-oxygen bond. This suggests that the observed pronounced variation in oxygen desorption energy is largely due to electrostatic effects [21]. In addition to providing a direct explanation for the effect of electrochemical promotion, this study provides a theoretical methodology for the investigation of the role of promoters in heterogeneous catalysis.

4. CONCLUSIONS Electrochemistry can be used to affect oxidation catalysis on metals and metallic oxides [13] in a very pronounced and reversible manner. The observed promotional phenomena are due to an electrochemically driven and controlled backspillover of ionic species on the catalyst surface. These species, which in some cases cannot form via gas phase adsorption, alter the catalyst work function and affect the binding strengths of chemisorbed reactants and intermediates in a pronounced and theoretically predictable manner. This electrochemically controlled variation in the binding strength of adsorbates causes the observed pronounced modification in catalytic activity and selectivity. The ability of solid electrolytes to act as reversible promoter donors to influence oxidation catalysis is of considerable theoretical and, potentially, practical interest.

Acknowledgement: We thank the JOULE CEC programme, and the EPET and PENED programmes of the Hellenic Secretariat of Research and Technology for financial support.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

93

Reductive and oxidative activation of oxygen for selective o x y g e n a t i o n of h y d r o c a r b o n s Kiyoshi Otsuka Department of Chemical Engineering, Tokyo Institute of Technology, 2-12-1 Ookayama, Meguro-ku, Tokyo 152, Japan

A simple method for reductive activation of dioxygen by applying electrochemical O 2 - H 2 cell reactions is proposed for the oxygenation of alkanes and aromatics at cathodes. The active oxygen species for most of the cathodes with and without metal compounds was suggested to be OH radicals. However, the active oxygen species in the case of a SmC13/graphite cathode was not an OH radical. The electrolysis of water formed oxidatively activated O 2-, which caused the epoxidation of various alkenes including propene to propene oxide on a PtO 2 phase for a Pt black anode. A catalytic system composed from Eu-salts, carboxylic acids and zinc powder caused the partial oxidations of light alkanes (CH4, C2H6, C3Hs) at room temperature and atmospheric pressure. In this system 02 was reductively activated by Eu(II) which was generated from Eu(III) through reduction by zinc powder. The active oxygen formed in this manner showed a strong electrophilicity as well as a radical character.

1. I N T R O D U C T I O N Partial oxidations of alkenes to epoxides, light alkanes (CH4, C2H6, C~Hs) to their alcohols and aldehydes or hydroxylation of aromatics in one-step using 02 as the oxidant are the most attractive but difficult targets in the field of selective oxidation catalysis. Monooxygenase or its mimic systems are often applied for these oxygenations under mild conditions, using a reductant such as ascorbic acid, NADPH, NaBH 4 or zinc for reductive activation of oxygen making possible the monooxygenations [1-6]. Although the mechanisms for the activation of oxygen are complicated and the active oxygen species suggested are quite different, the reductive activation of O2 in acidic media and its suggested reaction may be represented by equations 1 and 2. 02 + 2H + + 2e-----O* + H 2 0 RH + O* --~ R OH

(1) (2)

94 I n this scheme e- is provided by the reductant added to or present in the system and O* is the active oxygen species responsible for the oxygenations of hydrocarbons. The attempt to reductively activate O2 using the cheapest reductant H 2 [7-11] is attractive but careful handling is required for the explosive O2-H 2 gas mixture. Moreover, the deep reduction of O2 to H20 must be controlled. 1 . 1 R e d u c t i v e a c t i v a t i o n o f 0 2 by O2-H 2 c e l l Recently, we have reported a simple method for the reductive activation of 0 2 by applying an electrochemical O2-H 2 cell reaction for the oxygenation of alkanes and aromatics at the cathode [12-14]. With an acidic electrolyte in the O2-H z cell, the stoichiometric anode and cathode reactions are written simply as follows, (Anode) (Cathode)

H 2 ---, 2H* + 2 e 1/2 0 2 "[" 2H* + 2 e ---~ H 2 0

(3) (4)

Although the elementary steps for the cathode reaction must be complicated and dependent on the nature of the electrocatalyst, the reduction of O2 at the cathode may be written schematically as follows. O2

e-

~-

02-

e-

~- O222HO" MO2H

e-

~

023-

e-

~- 2H20

(5)

HO; H20 M-O; H20

If the reduced oxygen intermediates including the protonated ones in parentheses have a finite lifetime in the presence of a suitable catalyst (M), these reduced oxygen species might be able to attack the hydrocarbons in the cathode compartment, resulting in their oxygenation during the O2-H z cell reactions. In fact, we succeeded in the oxygenation of alkanes and the hydroxylation of aromatics on the basis of this concept [12-14]. 1 . 2 O x i d a t i v e a c t i v a t i o n o f H20 by In contrast to the reductive activation described above, the reverse reaction, should oxidatively activate the oxygen follows. electrolysis H20 H2+ O** RH + O** -~ ROH

electrolysis of 0 2 during the O2-H 2 cell reaction i.e., the electrolysis of acidic water of water at the anode (equation 6).

(6) (7)

The so-formed oxidatively activated oxygen O** can be expected to also activate hydrocarbons producing oxygenated products under mild reaction conditions. In fact, the epoxidation of olefins proceeded selectively by the oxygen generated at the anode during water electrolysis at room temperature as will be demonstrated later.

95

1.3 Design of catalytic systems based on the zinc-air battery The zinc-air battery (ZnIKOH]Air) is a commercialized, small b u t t o n - s h a p e d battery that uses KOH as an electrolyte. If we use a q u e o u s acid s o l u t i o n as an electrolyte, zinc is oxidized into zinc salt (equation 8) at the anode and o x y g e n is r e d u c e d into water at the cathode (equation 4). (Anode)

(8)

Zn + 2 H A --~ ZnA 2 + 2H + + 2 e

Similar to the concept d e s c r i b e d for the O2-H 2 cell system, the r e d u c t i v e activation of o x y g e n can be expected at the cathode in the p r e s e n c e of a suitable catalyst. On the basis of this zinc-air battery model, we have d e s i g n e d a n u m b e r of catalytic s y s t e m s from mixtures of zinc p o w d e r , c a r b o x y l i c acid and v a r i o u s metal chlorides for o x y g e n a t i o n s of alkanes and alkenes. In these catalytic s y s t e m s , zinc p o w d e r w o r k s as the reductant as well as the electron c o n d u c t i n g medium. The c a r b o x y l i c acid w o r k s as a p r o t o n - c o n d u c t i n g m e d i u m . O x y g e n is reductively activated on the metal cations by p r o t o n s from the c a r b o x y l i c acid and electrons from zinc p o w d e r .

2. E X P E R I M E N T A L

2.1

Oxygenation

0 2 Vent in

an

~~HT~ Vent O 2 - H 2 cell reactor The O2-H 2 cell reactor and the principle of the method for the o x y g e n a t i o n of hydroc a r b o n s are d e m o n s t r a t e d in F i g u r e 1. A detailed description of the cell setup has been -, H2 given elsewhere [ 13]. A --~ S ' ~ ~""J silica-wool disk ( t h i c k n e s s /"<x:~" I / ]~'\~ ~Anode 2.0 mm and diameter 26 ram) [ Cathode HaPO4 aq impregnated with aqueous Spinbar H3PO 4 (1M, 1 ml) as an electrolyte separates the Figure 1. Reductive activation of dioxygen using O2-H 2 a n o d e and cathode compartH 2 cell reaction for oxygenation of hydrocarbons. ments. The anode was prepared by a h o t - p r e s s m e t h o d from a mixture of graphite, Pt-black and Teflon p o w d e r . The cathodes w e r e p r e p a r e d from a mixture of carbon w h i s k e r s with v a r i o u s metal oxides or metal salts. The superficial s u r f a c e area of both electrodes was 3.1 cm 2. The o x y g e n a t i o n of b e n z e n e was carried out in the f o l l o w i n g s t a n d a r d manner: (i) O x y g e n (101 kPa) was b u b b l e d into the liquid s u b s t r a t e s (40 ml) in the c a t h o d e c o m p a r t m e n t , (ii) h y d r o g e n (49 kPa) and water v a p o r (4 kPa, to keep the electrolyte wet) w e r e p a s s e d t h r o u g h the anode c o m p a r t m e n t , (iii) the reaction was initiated by s h o r t i n g the circuit at 300 K and c o n t i n u e d for 3 h. In the case of o x y g e n a t i o n of light alkanes (CH4, C2H6, C3H8) , a gas mixture

/Ioo- 9

22~[ n'~l~

:.

96 of alkanes (50 kPa) and oxygen (51 kPa) were passed in the cathode compartment. Such an explosive gas mixture must be handled carefully in the reactor barricaded by thick (2 cm) acrylic plates. The oxidation efficiency (O.E.) for the formation of oxygenates was defined as follows, where F=96500 C. amount of the oxygenates(mol) (O.E.)=

x100%

(9)

charge passed(C)/2F 2 . 2 Oxygenation by H20Ar or electrolysis Propylene The reactor for H20electrolysis used in the epoxidation of alkenes is shown schematically in e Figure 2. The anode was prepared from metal blacks (70 mg) mixed with Teflon powder by the hot-press - R-CH=CH-R' method. The cathode was I-hO/Ar prepared from a mixture of Pt ~R.CH.CH.R? ( black, graphite and Teflon \O/ ~. powder. Propylene was bubbled into CH2C12 (40 ml used as a solvent) in the "T--'~-"~node H3[~O4aq. cathode anode compartment. In the magnetic case of 1-hexene epoxidation, spinbar 10 ml of the alkene was Figure 2. Schematic diagram of the reactor and dissolved in 30 ml of the principle of the epoxidation of alkenes duCH2CI 2. Argon (98 kPa) and ring electrolysis of water. water vapor (4 kPa) were passed through the cathode compartment to prevent the electrolyte from drying out as well as to remove hydrogen during the electrolysis. The oxidations of propylene and 1-hexene were performed under the following conditions unless otherwise stated" T=303 K, P(C3H6)=101 kPa, applied v o l t a g e = l . 7 V and reaction time=2 h. The oxidation efficiency for the formation of the oxygenates concerned was calculated according to equation 9.

J;LI

_

2.3 Design of catalytic systems for oxygenation The standard procedurc for thc oxygcnation of hydrocarbons using the catalytic systems dcsigncd from thc zinc-air battery was as follows. Mctal salts wcrc dissolvcd in a stirred solution of CH3COOH (or CF3COOH ) and CHzC12 in a three-necked flask with a reflux condcnser. Then a substratc and zinc powder were added to the solution. Thc oxygcnation of the substratc was started by stirring thc mixturc undcr a strcam of 02 (101 kPa) at 313 K, and the rcaction was continucd for 1 h.

97 3. R E S U L T S

AND DISCUSSION

3 . 1 P a r t i a l o x i d a t i o n d u r i n g O2-H 2 r e a c t i o n s . Effective cathode catalysts for partial oxidation 50 of hydrocarbons by the ~] cyclohexanone method in Figure 1 were 1 cyclohexanol looked for using cyclo40

hexane as a model substrate. The metal salts to ~" be tested as electro~ 30 catalysts were added to ,~ graphite by impregnation ~ 20 from aqueous solutions of their metal chlorides and were dried at 373 K. The 10 reaction products were cyclohexanol and cyclohexanone and no CO 2 was 0 Gr. Mg Sr Sc La Pr Sm Gd Dy Er Yb produced. The active caFe Ca Ba Y Ce Nd Eu Tb HoTm Lu thode catalysts contained Figure 3. Product yield in the oxidation of cyclohexane rare earth metal cations. for the cathodes containing various metal chlorides. Among them, S m 3+ was T=298K, reaction time=20h, most active for the partial Cathode: Metal chloride (25pmol)/graphite (70mg) oxidation of cyclohexane cyclohexane (40ml), P(O2)=101kPa, (Figure 3). The integrated turnover number for the Anode: Pt-black/graphite, P(H2)=98kPa , P(H20)=3.3kPa. Sm 3+ exceeded 10 after six successive reactions. The SmC13/graphite cathode was also effective for oxygenations of other alkanes (n-hexane and adamantane) and hydroxylations of aromatics (benzene, toluene and naphthalene) at room temperature during O2-H z cell reactions. The treatment of graphite in oxidizing agents such as potassiumpermanganate, boiling HNO 3 before the impregnation of SmC13 enhanced the catalytic activity of SmC13/graphite. This favorable effect of oxidation pretreatment of the host graphite suggests that surface functional groups such as -COOH, -OH or -CHO are required for coordinating the active S m 3§ ions. The addition of H202 during reaction under standard experimental conditions did not enhance the oxygenation of cyclohexane, suggesting that H202, which should have been generated at the cathode during the O2-H 2 cell reaction, could not be the active oxygen species in the case of rare earth metal-embedded cathodes. Comparison of the results of toluene oxidation using SmC13/graphite with those using Fenton's reagent indicated that the active oxygen species O ' g e n e r a t e d on the SmC13/graphite was quite different from the OH radical. The O* s h o w e d more electrophilic character than the OH radical.

98 Cyclic-voltammetry studies of the SmC13-doped glassy carbon in the presence and absence of cyclohexane suggested that a hydroperoxy radical generated on Sm 3+ was directly or 50 r indirectly responsible for oxygenation of cyclohexane. 40 The h y d r o p e r o x y radicals o could be further activated on ~ 30 Sm 3+ giving a strongly perturbed h y d r o p e r o x y species ~ 2o Sm(3-~)+(O2H) ~§ which had electrophilic character and 0 ~9 lO could activate the C-H bond of alkanes. 0 The use of carbon ~ 0 whiskers instead of graphite CH4 C2H6 C3H8 Alkanes as the host cathode material markedly improved the Figure 4. Oxidation of light alkanes during O2-H2 cell rate of oxygenation of reactions 9 cyclohexane and benzene, Cathode: Carbon whisker; P(alkane)= 51, P(O2)=50kPa. especially for cathodes Anode: Pt-black/graphite; P(H2)=50, P(H20)=4 , containing iron compounds P(He)=47kPa. and Pd black [15]. The 1 , acetaldehyde for ethane and acetone for propane. incorporation of iron I--1 , carbon dioxide. compounds together with Pd black showed a m a r k e d synergism for the formation of oxygenates. In this case, hydroxyl radicals were suggested to be the active oxygen species responsible for the h y d r o x y l a t i o n of benzene and oxygenation of cyclohexane [15]. The method of Figure 1 without solvent was also effective for the oxidation of methane, ethane and propane at room temperature (Figure 4). The carbon whisker cathode without additives was the most effective for the partial oxidation of propane. The selectivity to useful oxygenates (acetone) in this case exceeded 65% on the basis of the propane reacted [16].

3 . 2 E p o x i d a t i o n of aikenes during H20 e l e c t r o l y s i s Materials effective as anode catalysts for epoxidation of 1-hexene by the method in Figure 2 were screened. Among various metal oxides, metal salts and metal blacks tested, the most active and selective anode catalyst for the formation of 1,2-epoxyhexane was Pt black (Table 1). The oxidation efficiency for the formation of epoxide defined by equation 9 was about 26% and its selectivity was 66%. Pt black samples obtained from different producers or prepared in this work showed quite low electrocatalytic activity. However, the calcination of these inactive Pt blacks in air at 673 K substantially enhanced the catalytic activities of these samples. XPS studies on various Pt black samples suggested that a PtO 2 phase was associated with the active oxygen for the epoxidation.

99 Table 1. Epoxidation of 1-hexene on various anodes of noble metals Amount of product / mmol 9m -2 Anode

~O

Others

C.P./ 104 C 9m -2

Total O.E. /%

Epoxide O.E. /%

SEpo. /%

Pt-black

1066

394

170

80.8

42.4

25.5

65.6

Pd-black

43

24

25

118.6

1.8

0.7

47.3

Rh-black

3

3

21

142.1

0.6

0

9.8

Ru-powder

2

1

6

1.9

9.4

2.3

30.8

Au-powder

1

2

4

1.7

9.3

1.0

12.5

Reaction conditions: T-303K, time 2h, applied voltage 1.7V, H3PO4 1.0M. Anode: noble metals, P(Ar)=101kPa; Cathode: Pt-black/graphite, P(Ar)=98. P(HzO)=3kPa. The Pt black sample active for the epoxidation of 1-hexene and 2hexenes was also tested in the oxidation of propylene. Figure 5 shows the results of p r o p y l e n e oxidation as a function of the applied voltage across the cell. The oxidation of propylene was initiated at an applied voltage higher than ca. 1.1 V. The formation of p r o p y l e n e oxide and acetone were remarkably enhanced at an applied voltage >1.1 V. The m a x i m u m oxidation efficiency for the propylene oxide was 25% and the selectivity to propylene oxide was 53% at 1.7 V. These results indicate that the epoxidation of propylene by the oxidative activation of H20 proceeds with fairly good current efficiency and selectivity on the Pt black anode [17].

100

600 i

O 400

.,.,~ ~3

5o ~

r

O 9,,,-i .,..a

~" 200 (D

~

o 0 2.0O

rJ3 ,,,,o

o

1.0-

0

0

0.0

1

2

3

Applied voltage / V Figure 5. Partial oxidation of propylene during electrolysis of water as functions of applied voltage. Standard reaction conditions: T=303K, reaction time=2h, concentration of H3POg=I.0M. Anode: Pt black, P(C3H6)=101kPa. Cathode: Pt black/graphite, P(H20)=3, P(Ar)=98kPa. V , O.E. for propylene oxide and acetone; II, O.E. for propylene oxide; 0 , propylene oxide; ~ , acetone; [-7, COa;/X, acetic acid; X, propionic acid 9

100 The product ratio in cisand trans-2,3epoxyhexanes observed in the epoxidation of cis- and trans-2-hexenes indicated the retention of the cis-trans configuration of the starting 2hexenes. This observation demonstrates that the rotation around the C(2)-C(3) axis of the activated complex is strictly prohibited. Kinetic results and tests using propylene oxide and acetone as the s tar tin g s u bs trate suggest the reaction mechanism indicated in Figure 6 [17]. 3.3

Design

of catalytic

H

H

XO"

(1)

_ O _pt4_+O _pt4_+O _

~ -2H+, -2eo*

(2)

A

"......O*

(3)

: ..

J

~9O ,

O - P'r

(4)

- P'4+ ~

~,H,,~,O,(6)

.O'~*"'a~O, (8)

+

+

(5) O- Pt4-+O-Pt4+

(7)

O- Pt4-+0- Pt4+

4+

COl (9)

"~OX~H(10)

O-Pt4+-O- Pt4+

Figure 6. Reaction mechanism for oxidation of propylene. systems

based

on m i c r o c e l l

models.

Based on the zinc-air battery, a new catalytic system designed was composed of metal chlorides (MClx), carboxylic acids and zinc powder for oxygenation of alkenes to epoxides and partial oxidation of light alkanes (CH 4, CzH 6 and C3H8) into alcohols and aldehydes. These catalytic systems are likely to activate oxygen as schematically illustrated in Figure 7, where CF3COOH is used as a proton-conducting medium. Oxygen is reductively activated on the metal cations by protons from CF3COOH and electrons from zinc. Among the many metal salts tested in this work, Eu Salts "Microcell Catalyst" (EuC13, Eu2(CO3)3, Eu(AcO)3 ROH, RO RH Eu(C104)3, and Eu(NO3)3) 2 CF3COOH ~/// showed catalytic activity for the oxygenation of i + H20 hydrocarbons at room (CF3COO_)2Zn Mn+ temperature [18-20]. The /-~ 2 H + " - - ~ \ results for the oxidation of " ~ " C1 C1 (cathode) cyclohexane over various metal chlorides and Eu-salts 2eare compared in Figure 8. Zn The mixture of Eu saltsFigure 7. Reductive activation of dioxygen for light CH3COOH-Zn powder with alkane oxidation. CH2CI z as a solvent catalyzed the epoxidation of alkenes (hexenes, butenes, propene)

/

~

101 Yield based on cyclohexane / % 1

2

i

I

~ LaC13 ! CeCI3 " ! PrC13 NdCI3 SmC13

CyOH

CyO EuC13

" 1 GdCI3 1 TbC13 ~ DyC13 HoCI3 ErC13 "1 TmC13 YbCI3 LuCI3

Eu(NO3)3 Eu(CIO4)3 Eu2(CO3)3

"! Eu2(SO4)3

Eu(AcO)3

" FeCI3 " ! CuCI2 "1 PdCI2

0

5 TON based on catalyst / h -1

10

Figure 8. Partial oxidation of cyclohexane using various metal chlorideCH3COOH-Zn powder systems. Reaction conditions: T=313K, rection time=lh, catalyst=30gmol, Cy-C6H12=2ml, O2=101kPa, CH3COOH=2ml, Zn powder=l.0g, CH2C12=2ml. and partial oxidation of alkanes (adamantane, cyclohexane, n-hexane, propane). The mixture of Eu salts-CF3COOH-Zn powder in the absence of solvent catalyzed the oxidation of methane and ethane into their alcohols with turnover numbers of 4 and 8, respectively, in one hour at room temperature [20]. The important role of CF3COOH is to stabilize the alcohols formed by converting them into their esters. The reductively activated oxygen generated in this catalytic system showed strong electrophilicy as well as radical character [21]. If the catalytic system of Eu salts-CH3COOH (or CF3COOH)-Zn powder is deprived of one of its components, the partial oxidation of alkanes and epoxidation of alkenes occur very slowly or not at all. These observations suggest that the concept of designing catalytic systems based on the microcell model in Figure 7 may be justified.

102 REFERENCES

1. 2. 3. 4.

5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15.

16. 17. 18. 19. 20. 21.

J . T . Groves, T. E. Nemo and R. S. Myers, J. Am. Chem. Soc., 101 (1979) 1032. Ortiz de Montellano (ed.), "Cytochrome P-450, Structure, Mechanism and Biochemistry", Plenum Press, New York, 1986. F. Montanari and L. Casella (eds.), "Metalloporphyrins Catalyzed Oxidations", Kluwer Acad. Pub., Dordrecht, (1994). D . H . R . Barton, M. J. Gastiger and W. B. Motherwell, J. Chem. Soc., Chem. Commun., ( 1 9 8 3 ) 4 1 ; D. H. R. Barton et al., J. Chem. Soc., Perkin Trans. I, ( 1 9 8 6 ) 9 4 7 . N. Kitajima, H. Fukui and Y. Moro-oka, J. Chem. Soc., Chem. Commun., (1988) 485. H. Dalton and J. Green, J. Biol. Chem., 264 (1989) 17698; J. Colby, K. I.Stirling and H. Dalton, Biochem. J., 165 (1977) 395. I. Tabushi and A. Yazaki, J. Am. Chem. Soc., 103 (1981) 1771. N. Herron and C. A. Tolman, J. Am. Chem. Soc., 109 (1987) 2837. A. Kunai, T. Wani, Y. Uehara, F. Iwasaki, Y. Kuroda, S. Ito and K. Sasaki, Bull. Chem. Soc. Jpn., 62 (1989) 2613. A. Sato, T. Miyake and T. Saito, Shokubai (Catalyst), 34 (1992) 132. Y e W a n g and K. Otsuka, J. Catal., 155 (1995) 256. K. Otsuka, I. Yamanaka and K. Hosokawa, Nature, 345 (1990) 697. I. Yamanaka, and K. Otsuka, J. Chem. Soc., Faraday Trans., 89 (1993) 1791. I. Yamanaka, and K. Otsuka, J. Chem. Soc., Faraday Trans., 90 (1994) 451. K. Otsuka, M. Kunieda and H. Yamagata, J. Electrochem. Soc., 139 (1992) 2381; K. Otsuka, M. Kunieda and I. Yamanaka, Stud. Surf. Sci. Catal., 82 (1994) 703. Q. Zhang and K. Otsuka, Chem. Lett., No 4, (1997) 363. K. Otsuka, T. Ushiyama, I. Yamanaka and K. Ebitani, J. Catal., 157 (1995) 450. I. Yamanaka, K. Nakagaki and K. Otsuka, J. Chem. Soc., Chem. Commun., (1995) 1185. I. Yamanaka, T. Akimoto, K. Nakagaki and K. Otsuka, J. Mol. Catal. A., 110 (1996) 119. I. Yamanaka, M. Soma and K. Otsuka, J. Chem. Soc., Chem. Commun., (1995) 2235. I. Yamanaka, K. Nakagaki, T. Akimoto and K. Otsuka, J. Chem. Soc., Perkin Trans. H, ( 1 9 9 6 ) 2 5 1 1 .

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

103

T h e S e l e c t i v e O x i d a t i o n of M e t h a n o l : A C o m p a r i s o n of t h e M o d e of A c t i o n of M e t a l a n d O x i d e C a t a l y s t s D. Herein, H. Werner, Th. Schedel-Niedrig, Th. Neisius, A. Nagy, S. Bernd, R. SchlSgl Fritz Haber Institut der Max-Planck Gesellschaft, Faradayweg 4-6, D-14195 Berlin i

1. A b s t r a c t The reactivity of methanol towards oxygenated silver and copper substrates and towards the molecular oxide H4PVMoiiO40 has been investigated by a variety of insitu techniques. The intention was to find the chemical origin of the selective action of atomic oxygen and to discuss the influence of the metal species on the oxygen reactivity. The exceptionally high reaction temperature over silver was traced back to the difficulty of forming one active oxygen species. Surface spectroscopies found evidence for three chemically inequivalent atomic oxygen species in the metal-oxygen systems in analogy to the three oxygen forms in the molecular oxide characterized by a single crystal structure analysis. The large body of surface science experiments in these systems has contributed to our understanding of the reaction possibilities but needs to be treated with reservation when a ,,high pressure" reaction mechanism is considered as these experiments describe only part of the system in the static limit of low chemical potential of the gas phase.

1. I n t r o d u c t i o n The selective oxidation of methanol to formaldehyde is a technologically relevant [1,2] reaction carried out over metallic silver or iron-molybdate catalysts. The reaction is also suitable as a model reaction to understand the chemical principles of modifying the oxidizing potential for dissociated oxygen on various surfaces [3,4,5,6,7]. The methanol molecule is a special organic substrate as it contains only hydrogen atoms in alpha position to the functional group. Chemically speaking

The work was supported by the Bundesministerium ftir Bildung und Forschungthrough its catalysis programme. Special supportand helpful discussionscamefromthe BASFAG, Ludwigshafen.

104 methanol is more a derivative of the water molecule r a t h e r t h a n an organic molecule with a functional group. The acidity of the methyl protons is thus exceptionally high allowing their easy activation in selective oxidation. All larger molecules contain additional hydrogen atoms much less acidic than those of methanol and these are much more difficult to activate for elimination or substitution. Methanol is thus a poor model molecule when specific acidity or delicate chemisorption properties of e.g. olefin functions are of interest. The surface chemistry of methanol [5,8,9] and its co-adsorption systems with oxygen over single crystal surfaces have been reviewed extensively [3,10]. Mechanistic studies on single crystalline metal oxide surfaces under well-defined conditions are much more scarce. The purpose of the present paper is to present a cumulative view on methanol oxidation under ,,high pressure" reaction conditions over a variety of surfaces. Experiments probing the geometric surface structure, its electronic structure and its reactivity will be presented which complement the significant body of ultrahigh vacuum (UHV)-oriented studies. Insitu experiments in which the surface was exposed to a flow of methanol-oxygen mixtures for prolonged times will be presented. The property of interest was either determined directly during the experiment or measured after a transfer into UHV. In contrast to m a n y preceding studies the attention is focused on the state of the active surface and not on the observation or indirect identification of short living reaction intermediates which are very difficult to characterize under insitu conditions. In this way the dimension of the material gap will be reduced between well-defined single crystals of metals (single crystals of relevant oxides of size and defect concentrations as applied for element metals do not exist as yet) and technical catalysts.

2.The Activation of Molecular Oxygen Molecular oxygen is a difficult-to activate chemical species and requires the presence of an electron donating catalytic surface such as a metal or a partly reduced oxide. The sticking coefficient of oxygen is usually low (about 10 .2) on clean metals and on fully oxidized oxides but rises to values of about 10 -2 on oxygen pre-covered metals [11,12] or partly reduced [13] substrates. A general reaction pathway [8] is depicted in Scheme 1. From a molecular precursor [14] (not shown) oxygen reacts to the peroxo state at which the oxygen-oxygen bonding is still present in a weak form. This state is extremely active in oxidation [15,16] but usually unselective. It occurs on silver [17] at about 300 K. At higher temperatures the peroxo species dissociates completely and forms atomic oxygen which can be either present as surface-adsorbed (a) species or as surface-intercalated (7) species [18]. The overall sticking coefficient into these selectively reacting species is low as the likely processes for depleting the precursor states are desorption or bulk-dissolution [12]. The chemical bonding [19] of the two species to a metal substrate is significantly different and can be characterized as O 2p-metal sp hybrid for (a)

105 and O 2p-metal dsp hybrid for (7) oxygen. On silver the (7) state is the most stable form of oxygen complex and can be formed either directly, or peroxo by conversion from the (a) state or by segregation from bulk-dissolved oxygen. The chemical reactivity of the two surface species is quite different + and can be characterized as oxidizing for the more weakly held (a) species and as basic (,,proton activating") for the strongly bound (7) species. Evidence for these assignments comes from a suite of spectroscopic experiments [7,17,18,20,21,22,23]. oxidizing basic Figure 1 shows the thermal desorption spectroscopic (TDS) BijerrumBr~ electrophilic j / ~ responses of a polycrystalline electrolytic silver catalyst [24] from the three species (a), (7) and bulk dissolved with their respective T+ desorption temperatures and line profiles. The superimposed conversion curve (atmospheric Scheme 1" Products of the activation of molecular pressure, methanol : oxygen ratio x = oxygen on metals. 1.0, SV 80.000 h -1, methanol to water ratio 1"1, reaction time 130 h) illustrates that under technical reaction conditions the reaction is driven by (7) and by bulk-dissolved oxygen converting during segregation into (a). The fact that the conversion curve TDS 100 .-. alpha exhibits at 50% the plateau in 10 g exact agreement with the :5 8o ~ changeover from bulk 8 m eAg- poly Q~ dissolved to (7) oxygen is seen 60 E 6 as qualitative indication (see 0 r also Figure 7) for the change 0 4 40 .~ in the dominant reaction III g pathway from oxidative 20 ~ dehydrogenation to dehydrogenation together with 400 500 600 700 800 900 the dominating surface Temperature (K) species. Figure 1" Deconvoluted TDS responses from three different Figure 2 exemplifies the oxygen species on silver spectroscopic fingerprints for

I

A

106 the different bonding schemes mentioned above. The method of near edge X-ray absorption spectroscopy (NEXAFS) at low photon energies requires a complex UHV system with a reaction cell AgxO O K-edge high-pressure ....o 4 allowing treatments at 10 mbar oxygen for 300 s t o 3600 s on an -o .>.. - 3 Ag (111) single crystal between 600 K and 900 K in order to 0 ~2 prepare the relevant species. NEXAFS [19] can be used as a highly surface-sensitive tool with about 80 % of the information 011 arising from less than 2 nm depth 520 530 540 550 560 570 of the sample. The unique Photon Energy (eV) resolving power of the method for the details of the hybridisation is Figure 2: NEXAFS data of three atomic oxygen evident. The sharp first structure species on silver is the n* resonance arising from transitions from O ls into unoccupied d-sp hybridised states exhibiting a crystal-field splitting fine structure. The broader features above 536 eV arise from the o* ' transitions from O ls into unoccupied sp r states. The significant differences in chemical bonding between oxygen atoms ~ 2 2o and substrate are much more clearly seen than in the more conventional X-ray 0 photoelectron spectroscopy (XPS) in which ((z) and (?) give rise to a chemical shift of s about 1.5 eV without disclosing, however, 6 ~ ~ / ~ ~ the origin of the chemical shift. The absence of multiple n* resonances and of any polarization effects on well ordered Ag surfaces (data not shown, see [19]) prove . ~ 1 together with Raman spectra [25,18] without any doubt the atomic nature of all three species. The peroxo species existing 40o 600 soo at low temperatures and being not Temperature(K) relevant to selective methanol oxidation is difficult to prepare in a spectroscopically Figure 3: Isotopic exchange TDS with 18 02 pure form as its reactivity leads always to of fully oxidized HPA at 1 K/s heating rate. interference with carbonate [17,8] H2CO was obtained in a TPRS run. formation. Observation of reactive oxygen species on oxides is significantly more difficult. Figure 3 shows an isotope exchange experiment with the molecular oxide HnPVMollO40-14 H20, a heteropoly acid (HPA) [26,27] and oxygen 18 in the gas ,

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107 phase. The sample stored molecular oxygen together with the crystal water which exchanged readily at low temperatures due to the acidic environment. Attempts to chemisorb 18 02 onto the HPA surface after dehydration were unsuccessful even at atmospheric pressure. In an atmosphere of 18 Oz it was possible, however, to observe a dynamic exchange between gas phase and lattice oxygen as could be seen from the increase of the scrambled oxygen partial pressure with temperature. At 685 K the structure began to collapse giving rise to an increased scrambling rate and at slightly higher temperatures to oxygen evolution from the sample bulk. The correlation with the formaldehyde formation in an atmosphere of methanol is useful for an assignment of oxygen reactivity. A sustained production of formaldehyde was observed only after the constitutional water was lost lasting up to the temperature at which the breakdown of the structure finally gave rise to an increased (but not lasting) activity. The data show that the HPA is only active in this reaction in its dehydrated state and that the useful temperature window at which steady state oxygen activation from the gas phase is observed is limited to the interval between full dehydration and beginning decomposition. The activation of molecular oxygen is likely to be a complex process requiring a significant amount of activation energy to overcome the 493 kJ/mole bond energy in molecular oxygen which can be moderated by the ease at which electrons are transferred from the catalyst to the chemisorbed oxygen molecule. As the whole process needs to be reversible it is of course not possible to use strong reductants such as early transition metals or alkalis as these materials will not be reducible by the organic substrate. Rather inert metals such as silver and gold are a good choice as well as oxides with multiply valent states (defects).

3. Selective Oxidation Pathways The term ,,selective oxidation" can be defined as a set of reactions in which the formal oxidation state of the organic substrate is increased but not to the level of t h a t of CO or CO2. Selective oxidation is then not only the addition of an oxygen atom to the organic molecule , the abstraction of hydrogen via dehydrogenation and oxidative dehydrogenation (or frequently oxidehydation) is also enclosed in ,,selective oxidation". The dehydrogenation is an endothermic process due to the high energy content of C-H bonds without a suitable compensation by the exothermic water formation. For this reason there is a high tendency of the hydrogen product to react irreversibly with the metal-oxygen complex of the active site and to irreversibly damage it. Dehydrogenations can thus only occur with very stable catalysts and under conditions where the reductive chemical potential is modified by the presence of e.g. a large excess of molecular water [28](e.g. in styrene formation over iron oxide). Different redox reactivities with atomic oxygen are the consequence of differing oxygen-substrate interactions. The (a) species is either transferred to the substrate (oxidation) and/or to hydrogen (oxidehydrogenation). The (~,) species reacts as a basic center and is strongly attached to the catalyst which needs not to be a metal such as silver or

108 platinum [11,12] but may well be a partly reduced metal ion in a sub-oxide catalyst [29].

4 . O x i d a t i o n of M e t h a n o l Methanol can be transformed into formaldehyde in principally two different ways known as dehydrogenation and oxidative dehydrogenation. The main side products are hydrogen and water respectively. A cumulation of the mechanistic considerations drawn from model experiments on metal surfaces [3,6,30] is given in Scheme 2. The catalyst surface containing pre-adsorbed oxygen of either ((z) or ('1I) type transforms methanol into methoxy (1). This can react along three main pathways. Reaction (2) is the simple oxidative dehydrogenation leading to formaldehyde and water. Reaction (3) occurs with (~) oxygen and is a thermodynamically unfavorable dehydrogenation path. At long residence times and under excess (a) oxygen a series of complex reactions occurs which begins with the surface oxidation of methoxy (4) to formate. The labile formate + will either decompose into carbon dioxide and water (5) or even into carbon dioxide and hydrogen in the absence of sufficient chemisorbed oxygen. The desorption of formic acid (6) or the formation of dimethyl ether or dimethyl ketone are alternative side reactions. The relative values of the relevant rate constants will control the selectivity to the side reactions although they rarely dominate the overall conversion. The reaction pathways denoted by the intermediates (2) and (4) have been subject to extensive fundamental studies [3,6,31,32] whereas the dehydrogenation pathway was rarely found in such studies. The obvious reason for this problem Scheme 2: Reaction pathways of methanol on metals in the presence of different surface oxygen species. For the is that (?) oxygen cannot be generated from the gas phase numbers se text. under conditions accessible in UHV systems [20]. Only after prolonged use of a single crystal the bulk-dissolved

CH3OH

(1)~O--HH O--CH H

~ HH~1 methoxy HC

(2

H--O HC--O--~

H H

i

H

(3)

(4) 0

(5~0

formate I H--d-- I

/~'~

HC. (6) IH-o

109 species (unintentionally prepared) will segregate to the surface during heat t r e a t m e n t s and create some (7) oxygen. The reactions described in Scheme 2 require the existence of metal sites besides chemisorbed oxygen. The existence of active sites containing both metallic and oxygen centers was imaged directly in atomic resolution on silver saturated with (7) oxygen. An insitu (111) faceted grain of electrolytic silver was used for the image shown in Figure 4. Within the hexagonal a r r a y of silver atoms (normal interatomic distance of 0.285 nm, white) oxygen atoms are intercalated in the top atomic layer of the catalyst. The large contrast arises from the profound change in local electronic work function [19] on going from metallic to anionic atom sites. The acidity of intermediate OH groups formed on metal Figure 4: STM image of Ag (11 l) with (7) catalysts is assumed to be insufficient to oxygen. The interatomic distance of silver interfere with the elementary steps of measured along the white line is 0.285 nm. oxidative dehydrogenation. Oxide surfaces are significantly more For the ,,hole" contrast see text. complex in the range of reactivity of their t e r m i n a t i n g atoms. On HPA strongly acidic OH groups with pka values of mineral acids exist [33]. Surface hydroxyl groups on nominally pure oxides are also known to react acidic with a wide distribution of pka values. ,,Metal" centers a r e , in contrast to frequent colloquial designations, not existent on most catalytic oxides. Cation sites from regular metal-oxygen polyhedra are also difficult to access by adsorbates due to the inherently large size of oxygen ions relative to all cations. Defects in closed-packed surfaces and incompletely coordinated oxide polyhedra (e.g. octahedra) are, however, accessible for chemisorption of oxygen containing adsorbates such as methoxy. These sites are Lewis acids in contrast to metal sites on the metal-oxygen catalysts which are all in contact with the conduction band and hence Lewis bases. A variety of defect types is conceivable and has also been observed experimentally on binary and on complex oxide surfaces [34]. Figure 5 shows an atomically resolved STM image of a FeaO4 (111) single crystalline surface to give an impression about the complexity of a defective oxide surface. The oxygen atoms of 0.26 nm a p p a r e n t size form a regular a r r a y of a hexagonal closed packing as expected from the bulk structure. A hexagonal Moiree pattern indicates slight lattice mismatch between surface and bulk of the oxide film [35]. Point defects of various geometry and the modified local electron density on the adjacent oxygen atoms (lighter contrast) can clearly be seen. A fundamental problem for experimental [36,37] and theoretical [38] studies is our still very limited knowledge about geometry and defect disposition [34] on most oxide surfaces even when they are considered as

110 single crystals by diffraction techniques. The view that oxides contain two types of oxygen sites termed frequently as ,,terminal" (7) and ,,bridging" ((z) is certainly not incorrect [39] but maybe oversimplifying. As a consequence of the oxygen species distribution and the considerable acidity of oxide surfaces alternative reaction pathway for oxi-dehydrogenations have to be taken into account. The cleavage of the C-O bond in methanol becomes now a feasible reaction. The strong acidity allows further dehydration and the formation of dimethyl ether and of the acetal of formaldehyde as a prominent side reactions at low conversions [37, 40]. The kinetics of the oxidative dehydrogenation may exhibit four rate determining processes involving the formation of the methoxy intermediate, the cleavage of the C-O bond, the activation of the methyl hydrogen combined with the desorption of the formaldehyde product or the re-oxidation of the active site under evolution of water. It is the overwhelming conjecture [32] that for alcohol dehydrogenations the methyl activation step should be rate-limiting and neither the re-oxidation [40,41] nor the formation of alkoxy [42,43,44] control the overall kinetics. The methyl protons should be quite acidic considering the bonding situation in methanol and hence should be activated by a basic oxygen site [44]. This is in conflict with the observation [41] that basic oxide surfaces tend to totally oxidize the alcohol (multiple proton abstraction) and that acidic oxides exhibit a correlation between acidity and formaldehyde production. On suitably acidic surfaces the reaction mechanism may be dominated by the electrophilic action of protons on the methanol leading to an ,,embedded" methoxy which is easily activated at the methyl positions. On HPA catalysts this reaction path was claimed to be dominating [27] as long as the catalyst was sufficiently acidic. For the dehydrogenation of the methyl proton a radically different view was developed recently from theoretical Figure 5" STM image of Fe304 (111). The surface is considerations of HPA oxygen-terminated with each oxygen atom being resolved catalysts [38]. A direct 9LEED and ISS were used to verify the termination. ,,metal"-hydrogen bonding interaction between the proton

111 and one Mo 6§ center was claimed from an orbital analysis of an extended Hfickel calculation of transition geometries of this reaction step. It will be necessary to spectroscopically assess the covalency of the ,,metal"-oxygen interaction at the active sites in order to support this chemically surprising interpretation. The complexity of the methanol-oxide surface interaction is responsible for the still inconsistent picture about this reaction which may well proceed via several p a t h w a y s on the same bulk oxide. The reaction scenario in Scheme 3 is a reflection of the site complexity depicted in Figure 5 allowing for a wide spectrum of chemical reactivity of ,,metal"-oxygen groups. This spectrum is directly related Mo to the local geometric structure (,,bond I + CH3OH distance") and it is thus not surprising that a very pronounced structure H3CO sensitivity of the selective oxidation of Ogasphase O--..~ ~--~O methanol was found [36,37,45]. The silver-oxygen system restructures under reaction conditions into a uniformly active surface regardless of its initial O~ljO ~ O~ /O Mo Mo + CH30 H orientation [46,47] and acts as a structure-insensitive catalyst. - H2CO t + Olattice Y " H20 In the following we will investigate H H tO---.. ~ O t the methanol oxidation process over H Mo ] H silver, copper and heteropoly Mo molybdates in order to identify the + CH3OH occurrence of the possible reaction l H2CO pathways from Schemes 2 and 3 on l~3COH polycrystalline surfaces and at atmospheric pressure. The main O~ jO Mo Mo emphasis in these experiments will be on the source of active oxygen. This focus was chosen to better u n d e r s t a n d Scheme 3" Reaction pathways of methanol the involvement of bulk and sub-surface [10,48] chemistry in selective oxidation oxidation over early transition metal oxides. catalysis. Such an u n d e r s t a n d i n g is required when the catalytic performance is compared between series of chemically different systems which are chosen to investigate only one surface property such as acidity. These experiments will naturally only cover a small selection of the problems discussed with the reaction Schemes.

5. A n , , I n t e r m e d i a t e "

Experiment

The surface science studies of methanol oxidation over metals were carried out u n d e r conditions where no (y) oxygen was present on the surfaces. A t e m p e r a t u r e - p r o g r a m m e d reaction experiment was performed [49] in order to

112

prove that the reaction paths (2) and (3) can coexist on a surface. Electrolytic silver was loaded with molecular oxygen at 900 K and 10 mbar for 10 min. The generation of two species of atomic oxygen ((z) and (7) was confirmed by a TDS scan. After reloading the sample was heated in methanol vapor. The results are collected in Figure 6. The TDS trace exemplifies the well separated responses from (a) and bulk-dissolved oxygen (670 K) and from (y) oxygen at 990 K. In the TPRS traces the formaldehyde generation coincides exactly with the desorption peaks for the two oxygen species. This confirms the reactivity of both species and shows that the reaction is of the Langmuir-Hinshelwood type as the formaldehyde knows the bonding of the - ~ IIII" oxygen involved. The trace for carbon dioxide shows a parallel selectivity I ~ I I ~ I with partial oxidation for the (a) state II ~ iI but no total oxidation from the (7) state. 02 I A "'~" /"% This allows to assign the (a) state as the species for oxidative H2 dehydrogenation and the (y) state as the center specific for dehydrogenation. This assignment is supported by the m/e 18 trace exhibiting no intensity at the high temperature formaldehyde signal. The two sharp desorption peaks in the m/e 44 trace above 400 K arise from I I t i 400 600 800 1000 methoxy and carbonate (formate) Temperature[K] species which hardly produce Figure 6: TDS (dashed) and TPRS (full) of formaldehyde. This finding strongly oxygen/methanol on electrolytic silver. Heating underlines the notion that methoxy and rate: 1.3 K/s formate prepared in the surface science experiment as long-lived species [50] are different from methoxy and formate existing as reaction intermediates in the steady state conversion situation. It is speculated that minor differences in the chemical bonding of the chemically identical ad-species decide over their character as stable spectator species or unstable intermediate structures. Vibrational spectroscopic data [44] of methoxy and the analysis of the homologue ethoxy system [51] strongly support this view.

6. E x p e r i m e n t s w i t h Silver and C o p p e r The existence of the two reaction pathways to formaldehyde shown in Scheme 2 is documented in principal with the data from Figure 4. At steady state and at atmospheric pressure the contribution of the two reaction channels needs to be shown independently. In Figure 7 a section through the reaction parameter space of electrolytic silver under isothermal conditions is shown. The molar

113 stoichiometry was varied from oxygen-rich (in the explosion limit!) to methanolrich and the effect on conversion and selectivity was monitored. It is significant t h a t the optimum formaldehyde production was found at slightly methanol-rich conditions precluding the exclusive [2] contribution of oxidative dehydrogenation to the practical conversion. About 50 % of the formaldehyde production is observed at strongly understoichiometric feed compositions indicating t h a t dehydrogenation contributes significantly to the total activity. The selectivity to total oxidation is "-. "~ strongly coupled with 6O the excess of oxygen in the gas phase indicating ',/ \ w , t h a t the total oxidation II / I \ 10 . 20 o 40 of methoxy via formate is a rapid process when f, the surface is oxygen20 covered. The follow-up ,,............................ reaction to formic acid .............. i...................................... m ~I I I (see Scheme3) occurs not 50 1O0 150 200 250 in parallel with the Ratio CH30H : 0 2 (%) production of the Figure 7: Conversion over electrolytic silver at atmosphereic precursor formaldehyde pressure. A reactor holding 5.0 g catalyst in a 6 mm high bed was indicating that the used. The catalyst load was 10 g/h cm2. All data in mole % oxidation of reactive methoxy to formate is a unlikely process and that formic acid m a y be produced from the strongly adsorbed form of methoxy which occurs under conditions of lean reactive surface oxygen. From this only moderately reactive methoxy a finite selectivity leads to carbon dioxide in full accord with Scheme 2 as indicated by the increase of the CO2 production at very methanol-rich compositions. A quantification of the contribution of dehydrogenation to the total activity can be obtained from determination of the hydrogen gas in the tail gas. This value is s o m e w h a t uncertain, as reaction (5) from Scheme 2 will positively contribute and the consecutive hydrogen + oxygen reaction to water will negatively contribute. The data in the inset of Figure 7 indicate the contribution of several effects to the hydrogen yield. The change in slope at 873 K coincides with the occurrence of the (y) oxygen species (see Figure 4). The high t e m p e r a t u r e branch is attributed to the dehydrogenation with (~,) oxygen (reaction 3 in Scheme 2). The low t e m p e r a t u r e branch with low total conversion and significant evolution of CO2 contains significant contributions from formate decomposition to the hydrogen production. Evidence for the presence of several intermediates with significantly different residence times on the surface of silver was gained from a non-steady state conversion experiment. Electrolytic silver after 200 h time on stream was loaded at 870 K with pure oxygen [25,24]. After purging with nitrogen a stream of methanol was injected in the nitrogen stream from time 0 to 100 s and the

114 e v o l u t i o n of products was monitored every 40 s. The normalized response curves in Figure 8 indicate t h a t metallic silver cannot only store oxygen but also activated methanol in CO2 significant quantities X - " 6.4 T= 680 K ~. . . . HOHO~ on its surface. This is unexpected in the light of the failure to analyze the bonding I state of adsorbed methanol on silver at ! ! high temperatures ! I [17] with surface analytical tools where ! molecular methanol was found to desorb I I I 1 above ca. 200 K. The 32 33 34 35 product curves show Time (min) t h a t the Figure 9: Rate oscillations over copper chips. The methanol : oxygen oxygen-covered ratio was 6.4. Set temperature 680 K. SV 250 h~. surface is highly active for total oxidation. With progressing removal of the surface oxygen species [19] the selective partial oxidation wins with most likely a sizable contribution from dehydrogenation which does not require a stoichiometric a m o u n t of oxygen and can thus produce formaldehyde in a 100 situation lean in ~ ~ , ~ k H2CO surface oxygen (after ,, 80 ca. 300 s in Figure 8). The reaction product formic acid / / \ / \ \ sToK 60 from methoxy-toformate oxidation 40 does not form in parallel with the 20 abundance of methoxy (formaldehyde) nor C with the abundance of 0 1O0 200 300 400 500 Time (s) oxidizing surface oxygen. It occurs as final product from an Figure 8: Response of a 100 s pulse of pure methanol over oxygeni n t e r m e d i a t e residing predosed electrolytic silver. SV: 82000 h~. The response function of the non-reacting system (with SiC) was smaller than the CO2 profile, long on the surface as re-adsorption of m e t h a n o l is highly unlikely 300 s after the end of the pulse in a s t r e a m i n g system. The methoxy intermediate is, however, a precursor to this state as its

',: Ii

]

•

115 population vanishes with the removal of the methoxy species responsible for gas phase formaldehyde. The copper catalyst [6,30] offers an additional way of analyzing the contribution of multiple pathways to the practical reaction by studying the regime of oscillatory behavior. For a wide range of compositions temperaturecontrolled rate oscillations [52] can be observed in the methanol partial oxidation reaction. The origin of this unstable activity is the thermally driven interchange from an inactive bare copper metal surface to an oxygen-covered state which is triggered by oxygen being dissolved in the bulk [53]. This state leads easily to a binary oxide which is only active for total 1.5 oxidation. The rate """~ T = 620 K Cu-L3 oscillations for hydrogen, _ ~ O-K o,s,,o SS'* '~t ? 1.4 formaldehyde and CO2 are presented in high temporal resolution in 1.3 j," C u L III '~~~/ Figure 9. The contribution of several . ,,, 1.2 sources to the hydrogen production can clearly be I1 ~ ~ 1.1 seen. In the beginning of each~ oscillation period "I" 1, I I I I hydrogen and CO2 evolve 0 0.1 0.2 0.3 0.4 0.5 from decomposition of CHsOH- Pressure (mbar) formate. The selective Figure 10: Integrated normalized intensities of the first resonances oxidation via oxidative of Cu L II and 0 K for copper foil treated insitu with a mixture of dehydrogenation and via methanol in oxygen gas. dehydrogenation gradually wins over the total oxidation indicating the gradual modification of the copper-oxygen chemical bonding. This interpretation of the oscillations requires again the existence of chemically inequivalent oxygen species present at the copper surface in atomic form. Using X-ray absorption spectroscopy it was possible to substantiate the dependence of the chemical bonding between oxygen and copper on the chemical potential of the gas phase. Figure 10 shows the systematic increase in the spectral weight of the ~* resonance with decreasing oxidation potential of the gas phase (modified by the addition of methanol to the oxygen atmosphere at 50 mbar). The simultaneous change of the Cu L II white line indicates the gradual transition in bonding character from strongly interacting with rehybridised states (maximum in Cu sub-oxide, higher than in Cu II oxide) to non-bonding atomic-like with a weak O 2p-Cu 4s-p interaction. The reaction conditions imply that this state is responsible for the oxidative dehydrogenation. .

.

.

.

.

=(',,

y

OK

~,

116

7. Heteropoly Acids The discussion so far has shown that several surface oxygen species coexist with a bulk-dissolved species in metallic catalysts. The HPA as molecular analogue of complex oxide catalysts possesses three types of oxygen species namely surface terminating oxygen, surface bridging oxygen and bulk bridging oxygen. The latter species are possibly similar to bulk mobile anionic species that replenish defective oxygen sites where an organic substrate has removed an atom from a surface bridge or terminating site [54,55,56]. Aspects of scientific dispute [57,58,59,60,61] are the actual structure of the HPA anion in its active state which may be an opened Keggin cage (,,lacunary form"), the amount of water present under reaction conditions and the location of the excess electronic charges during a catalytic cycle. It has to be mentioned that the structural definition of the reacting HPA is difficult [62,63,64,58] despite its ,,molecular" character. The catalytically most relevant partially cation-exchanged derivatives [65,27,66] of the free acid forms are difficult to structurally assess as the apparently simple cubic X-ray diffraction pattern needs a detailed evaluation to disclose the real structure. For example four different models have been proposed for partly cation exchanged derivatives of H4PVMollO40. Here we concentrate on the free-acid form to avoid any interference with these structural problems. Under reaction conditions for methanol selective oxidation the HPA is nominally in its anhydrous form [63,64,67] which is metastable with respect to irreversible phase separation under prolonged catalytic reaction. Pulse experiments with varying delay times are suitable to follow the reactivity pattern. A single pulse response with a weakly oxidizing feed is presented in Figure 11. The most prominent effect is the strong capacity of the HPA to store the reaction water from the oxidative dehydrogenation reaction. Under reaction conditions the catalyst contains water of insufficient abundance, however, to

117 transform the anhydrous structure into the nominal 2-hydrate form [63,64]. The selective oxidation of methanol can proceed without gas-phase oxygen as seen from the pulse profile for formaldehyde in Figure 11. The dehydration to dimethyl ether occurs only in parallel to the presence of oxygen from the gas phase indicating a change in the surface acidity when the lattice oxygen reservoir is depleted. A series of 100 consecutive pulses was insufficient to consume all lattice oxygen available to the reaction at 573 K revealing the participation of deep volumes of the catalyst which should consist of a molecular crystal. Under reaction conditions the compacted crystal structure of the anhydride seems to allow oxygen transport to the surface. A series of pure oxygen pulses shown in Figure 12 allowed the replenishment of the lattice-oxygen reservoir. In addition, the total oxidation of strongly chemisorbed intermediates was indicated by the formation of a small amount of CO2. A measurable amount of formaldehyde was found only with the very first oxygen pulse indicating t h a t methoxy represents only a small proportion of the organic surface coverage. A significant abundance of water (see Figure 12) was liberated during the oxygen regeneration indicating that the catalyst was able to store hydrogen. This shows that also on the oxide catalyst a significant activity for dehydrogenation occurs which is not detected in steady-state measurements due to the consecutive hydrogen+oxygen reaction. A series of methanol pulses showed t h a t the catalyst converted it with lattice oxygen into formaldehyde and t h a t in a slow process some methanol was stored together with the reaction water. No liberation of the stored methanol with the gradually beginning water desorption was observed. The methanol seemed to be strongly adsorbed and to be converted into reaction products. All these findings show t h a t the bulk of the HPA crystal is actively involved in the catalytic action as it can store water, polar molecules, hydrogen and oxygen. Insitu UV-VIS spectroscopy [57] is a suitable tool to follow geometric structural integrity and electronic structural modification as function of reaction parameters. Figure 13

118 s u m m a r i z e s some results. The spectra exhibit two high energy m a x i m a at 250 and 310 nm characteristic of oligo-molybdate and the intact Keggin-ion (disappears upon opening of the cage structure). In its pristine (yellow) state the catalyst shows no other absorption. Upon h e a t i n g in oxygen or u n d e r inert conditions the loss of hydration w a t e r occurs (see inset) which is accompanied w i t h a slight electronic reduction as seen from peaks at 680 nm and at 1330 nm. These w e a k structures are due to intervalence charge t r a n s f e r bands from oxov a n a d i u m structures and do not occur with vanadium-free HPA. The inset shows the unexpected correlation of reduction and dehydration t r a n s f o r m i n g the HPA into an active state with the partially reduced electronic structure which is also reflected in the broadening of the Mo-O charger transfer bands. Cooling in w a t e r vapor restores fully the initial state indicating the reversibility of the slight reduction. After 723 K in oxygen and w a t e r vapor, the reduction is irreversible as seen by the change of the 310 nm absorption indicating a partial b r e a k d o w n of the Keggin ion. U n d e r m e t h a n o l vapor and conversion to formaldehyde the s a m e spectrum as u n d e r He at elevated t e m p e r a t u r e s can be seen. The intensity of the peaks depends critically on the methanol to oxygen ratio indicating t h a t this intensity can be used to monitor the degree of 1.2 .~ ~ . ~ . ~ bulk reduction of the ~'/~ catalyst or the 1

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600

800 1000 1200 Wavelength (mn)

1400

1600

1800

involvement of lattice oxygen. Figure 14

reports such an experiment in w h i c h the methanol-tooxygen ratio was varied from 10:1 to 1:2 without a significant change in reactivity. The 680 nm

absorption scales

very well with the Figure 13: Insitu UV-VIS spectra of HPA under various conditions, modified oxygen The inset compares the water TDS (thin) with UV absorption at 1330 abundance in the gas nm (thick) revealing reduction during dehydration, phase and is fully reversible in its intensity upon re-oxidation. The time scale for reversible reaction is, however, longer (6000 s) t h a n the waiting times in Figure 14. A slow diffusion controlled process of anion migration through a lattice is a likely explanation for the response characteristics observed. In s u m m a r y , the d a t a confirm the involvement of the bulk of the HPA crystals in the catalytic activity. They support the conjecture t h a t the v a n a d i u m acts as local sink for electrons from the catalytic cycle. The slow and highly activated oxygen m i g r a t i o n plays a practical role in catalytic conversion which is very m u c h more selective with lattice oxygen t h a n with gas phase oxygen. W a t e r is

119

beneficial at least in helping to maintain the structural integrity. At high levels of hydration which are structurally relevant, water suppresses the catalytic activity in this reaction. Thermal load alone leads to a slight reduction of the catalyst which is seen as beneficial for the function as it will enhance the sticking coefficient of molecular oxygen.

7. C o n c l u s i o n s The experiments have shown that mechanistic aspects discussed in the surface-science literature of methanol oxidation are of relevance under highpressure high-temperature conditions. Surface science has provided the techniques and fundamental insight into species present on metal surfaces. The remote reaction conditions do not allow the assignment of the stable intermediates which were spectroscopically characterized as the reaction intermediates. Experiments at conditions closer to practical conversion have revealed a novel oxygen species and cast doubt on the relevance of a stable methoxy species for catalysis. Several species of atomic chemisorbed I A I B I C I B I A oxygen were 0,2identified with surface science tools 0.15 and assigned to specific functions. n" ~" 0.1 These species occur on all catalysts investigated. Their 0.05 -02/MeOH structural disposition is, however, quite 0 20 40 60 80 1 O0 120 140 different in the Time (min) materials studied. The most relevant Figure 14: Evolution of the UV absorption for reduced V centers with different insitu reaction conditions. From times A to C the oxygen to problem for surface science remains to methanol ratio was reduced produce structural details of terminating and defective oxide surfaces which would be needed for meaningful theoretical approaches to the problem of selective oxidation. The key to understand selective oxidation is to characterize the chemical bonding of various atomic-oxygen species interacting with metal ,,ions" in quite distinct ways. Several examples of characterizations were discussed in this text. The degree of d-state interaction of the oxygen 2 p states decides over the covalent or anionic bonding character which translates into a basic or oxidizing function. In the presence of hydrogen the metal-to-oxygen bonding determines further the acidity of hydroxyl groups which can open additional reaction channels by electrophilic activation of the C-O bond. In all three catalyst systems ,

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120 investigated, the bulk plays an active and catalytically relevant role by providing reservoirs of oxygen atoms which control the type and abundance of surface oxygen species.. The reaction intermediates methoxy and formate seem to occur in two forms each. One is very active and carries the reaction under steady-state conditions. The other is more a spectator species with increased binding energy to the substrate which can, however, undergo complex side reactions and can contribute to the selectivity spectrum of the overall process. The external reaction conditions affect sensitively the co-operation of the dehydrogenation and oxidative dehydrogenation pathways leaving room for the speculation to devise a technical process for undiluted formaldehyde production at reasonable conversions.

References 1 G. Reuss, W. Disteldorf, O. Grler A. Hilt, Formaldehyde, in: Encyclopedia of Industrial Chemistry, Vol. Al1., Verlag Chemie 1988, S. 619-651. 2 H. Sperber, Chemie-Ing.-Techn. 41 1969, S. 962-966. 3 M. A. Barteau R. J. Madix, The Surface Reactivity of Silver: Oxidation Reactions, in: The Chemical Physics of Solid Surf and Heterogeneous Catal. D. A. King B. P. Woodruff, Ed., Vol./4, Elsevier 1982, S. 95-142. 4 M. A. Barteau, M. Bowker R. J. Madix, Surf Sci. 94 1980, S. 303-322. 5 M. A. Barteau R. J. Madix, Surf Sci. 97 1980, S. 101-110. 6 I. E. Wachs R. J. Madix, J. Catal. 53 1978, S. 208-227. 7 L. Lefferts, J. G. van Ommen J. R. H. Ross, Appl. Catal. 23 1986, S. 385-402. 8 C. T. Campbell, Surf Sci. 157 1985, S. 43-60. 9 M. Bowker, M. A. Barteau R. J. Madix, Surf. Sci. 92 1980, S. 528-548. 10 G. R. Meima, L. M. Knijf, A. J. van Dillen, J. W. Geus, J. E. Bongaarts, F. R. van Buren K. Delcour, Catal. Today 1 1987, S. 117-131. 11 R. J. Madix M. A. Barteau, Surf Sci. 97 1980, S. 101-110. 12 B. A. Sexton, G. B. Fisher J. L. Gland, Surf Sci. 95 1980, S. 587-602. 13 T. Seiyama, Surface Reactivity of Oxide Materials in Oxidation-Reduction Environment, in: Surface and Near-Surface Chemistry of Oxide Materials, Materials Science Monographs L.-C. Dufour J. Nowotny, Ed., vol 47, Elsevier, Amsterdam 1988, S. 189-215. 14 X. Bao, J. Deng S. Dong, Surf. Sci. 163 1985, S. 444-456. 15 V. I. Bukhtiyarov, A. I. Boronin, I. P. Prosvirin V. I. Savchenko, J. Catal. 150 1994, S. 268-273. 16 V. I. Bukhtiyarov, A. I. Boronin O. A. Baschenko, Surf Rev. Lett. 1 1994 Nr. 4, S. 577-579. 17 C. Rehren, M. Muhler, X. Bao, R. SchlSgl G. Ertl, Zeitschrift f Phys. Chemie 174 1991, S. 11-52. 18 X. Bao, M. Muhler, B. Pettinger, R. SchlSgl G. Ertl, Catal. Lett. 22 1993, S. 215-225. 19 X. Bao, M. Muhler, Th. Schedel-Niedrig R. SchlSgl, Phys. Rev. B 54 1996 Nr. 3, S. 2249-2262.

121 20 C. Rehren, G. Isaak, R. SchlSgl G. Ertl, Catal. Lett. 11 1993, S. 253. 21 H. Schubert, U. Tegtmeyer, D. Herein, X. Bao, M. Muhler R. SchlSgl, Catal. Lett. 33 1995 Nr. 3-4, S. 305-319. 22 G. Rovida, F. Pratesi, M. Maglietta E. Ferroni, Surf. Sci. 43 1974, S. 230-256. 23 K. C. Prince, G. Paolucci A. M. Bradshaw, Surf. Sci. 175 1986, S. 101-122. 24 D. Herein, A. Nagy, H. Schubert, G. Weinberg, E. Kitzelmann R. SchlSgl, Z. Phys. Chem. 197 1996, S. 67-96. 25 X. Bao, B. Pettinger, G. Ertl R. SchlSgl, Ber. Bunsenges. 97 1993, S. 322-325; Phys. Chemie. 26 T. Komaya M. Misono, Chem. Lett. 1983, S. 1177-1180. 27 J.-M. Tatibouet, M. Che, E. Serwicka, J. Haber K. Brfickmann, J. Catal. 139 1993, S. 455-467. 28 C. Gleitzer, Solid State Ionics 38 1990, S. 133-141. 29 A. Ferretti L. E. Firment, Surf. Sci. 129 1983, S. 155-176. 30 A. F. Carley, A. W. Owens, M. K. Rajumon M. W. Roberts, Catal. Lett. 37 1996, S. 79-87. 31 , Surface Reactions, in: Springer Series in Surface Sciences R. J. Madix, Ed., Vol. 34, Springer 1994, S. 1-282. 32 S. T. Oyama, Heterogeneous Hydrocarbon Oxidation, in: ACS Symposium Series B. K. Warren S. T. Oyama, Ed., Vol. 638, American Chemical Society 1996. 33 A. Bielanski, A. Cichowlas, D. Kostrzewa A. Malecka, Z. Phys. Chem. NF. 167 1990, S. 93-103. 34 P. L. Gai-Boyes, Catal. Rev.-Sci. Eng. 34 1992 Nr. 1-2, S. 1-54. 35 M. Ritter, H. Over W. WeiB, Surf. Sci. 371 1997, S. 245. 36 J.M. Tatibouet, J.E. Germain, J.C. Volta, J. Catal. 82 1983, S. 240-244. 37 J.M. Tatibouet, J.E. Germain, J. Catal. 72 1981, S. 375-378. 38 R. S. Weber, J. Phys. Chem. 98 1994, S. 2999-3005. 39 S. T. Oyama, W. L. Holstein W. Zhang, Catal. Lett. 39 1996, S. 67-71. 40 F. Lazzerin, G. Liberti, G. Lanzavecchia N. Pernicone, J. Catal. 14 1969, S. 293-302. 41 M. A], J. Catal. 54 1978, S. 426-435. 42 A. Desikan, S. T. Oyama W. Zhang, J. Phys. Chem. 99 1995 Nr. 39, S. 1446814476. 43 S. Ted Oyama W. Zhang, J. Phys. Chem. 1996 100, S. 10759-10767. 44 R. Miranda, C. O. Bennett J. S. Chung, J. Chem. Soc., Faraday Trans. 1 81 1985, S. 19-36. 45 J. E. Germain, Structure-Sensitive Catalytic Reactions on Oxide Surfaces, in: Adsorption and Catalysis on Oxide Surfaces G. C. Bond M. Che, Ed., Vol. 21, Elsevier Science Publishers, Amsterdam 1985, S. 355-368. 46 X. Bao, G. Lehmpfuhl, G. Weinberg, R. SchlSgl G. Ertl, J. Chem. Soc. Faraday Trans. 88 1992 Nr. 6, S. 865-872. 47 X. Bao, J. V. Barth, G. Lehmpfuhl, R. Schuster, Y. Uchida, R. SchlSgl G. Ertl, Surf. Sci. 284 1993, S. 14-22. 48, in: Surface and Near-Surface Chemistry of Oxide Materials, Materials Science Monographs L.-C. Dufour J. Nowotny, Ed., Vol. 47, Elsevier, Amsterdam 1998.

122 49 H. Schubert, U. Tegtmeyer R. SchlSgl, Catal. Lett. 28 1994 Nr. 2-4, S. 383-395. 50 M. Bowker, S. Poulston, R. A. Bennett A. H. Jones, Catal. Lett. 43 1997, S. 267-271. 51 W. Zhang S. T. Oyama, J. Am. Chem. Soc. 1996 118, S. 7173-7177. 52 D. Herein, G. Schulz, U. Wild, R. SchlSgl H. Werner, Catal. Lett. submitted 1997. 53 M. Lenglet, K. Kartouni, J. Machefert, J. M. Claude, P. Steinmetz, E. Beauprez, J. Heinrich, N. Celati, Mater. Res. Bull. 30 1995 Nr. 4, S. 393-403. 54 Oxygen in Catalysis J. Haber A. Bielanski, Ed., Marcel Dekker, Inc., New York 1991, S. 1-467. 55 K. Brfickmann, M. Che J. Haber J. M. Tatibouet, Catal. Lett. 22 1994, S. 241255. 56 G. Ya. Popova, T. V. Andrushkevich, V. M. Bondareva I. I. Zakharov, Kinet. Catal. 35 1994, S. 80-83. 57 M. Fournier, C. Louis, M. Che, P. Chaquin D. Masure, J. Catal. 119 1989, S. 400-414. 58 M. Fournier, A. Aouissi C. Rocchiccioli-Deltcheff, J. Chem. Soc., Chem. Commun. 1994, S. 307-308. 59 B. Herzog, M. Wohlers R. SchlSgl, Microchimica Acta 14 1997, S. 703-704. 60 V. M. Bondareva, T. V. Andrushkevich, R. I. Maksimovskaya, L. M. Plyasova, A. V. Ziborov, G. S. Litvak L. G. Detusheva, Kinet. Catal. 35 1994, S. 114-119. 61 E. Cadot, C. Marchal, M. Fournier, A. T~z~ G. Herv~, Role of Vanadium in Oxidation Catalysis by Heteropolyanions, in: Polyoxometalates M. T. Pope A. Mfiller, Ed., Kluwer Academic Publisher 1994, S. 315-326. 62 H. T. Evans jr. M. T. Pope, Inorg. Chem. 23 1984, S. 501-504. 63 Th. Ilkenhans, B. Herzog, Th. Braun R. SchlSgl, J. Catal. 153 1995 Nr. 2, S. 275-292. 64 B. Herzog, T. Ilkenhans R. SchlSgl, Fresenius J. Anal. Chemie 349 1994, S. 247-249. 65 N. Essayem, G. Coudurier, M. Fournier J. C. V~drine, Catal. Lett. 34 1995, S. 223-235. 66 Y. Toyozawa, N. Yamazoe, T. Seiyama K. Eguchi, J. Catal. 83 1983, S. 32-41. 67 B. Herzog, W. Bensch, Th. Ilkenhans, R. SchlSgl N. Deutsch, Catal. Lett. 20 1993, S. 203-219.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama,A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

123

Gold as a l o w - t e m p e r a t u r e o x i d a t i o n c a t a l y s t : f a c t o r s c o n t r o l l i n g activity and selectivity M. Haruta Osaka National Research Institute, AIST Midorigaoka 1-8-31, Ikeda, Osaka 563, Japan

The catalytic activity and selectivity of gold for oxidation can be widely tuned by control of three major factors ; selection of type of metal oxide supports, size of the gold particles, and control of the contact structure of the gold particles with the supports. These factors are markedly influenced by the preparation method. Examples are presented which show surprisingly high activities for CO oxidation at temperatures as low as -77~ and excellent selectivities to propylene oxide in the partial oxidation of propylene. It is found that there is a critical particle size of gold around 1-2 nm where the catalytic nature of the supported gold changes dramatically.

1. I N T R O D U C T I O N

Gold has attracted little attention as a catalyst because of its inert character and low melting point (1063~ which causes difficulties in depositing gold on supports with high dispersion. In the past, gold catalysts were considered not to be competitive with other noble metal catalysts in terms of activity [1, 2]. However, we have found remarkably high catalytic activity in the lowtemperature oxidation of CO with some supported gold catalysts prepared by coprecipitation [3, 4]. Since this work catalysis by gold has received growing attention and is currently investigated by many research groups [5]. The catalytic properties of supported gold catalysts appear to change dramatically when different metal oxide supports or different preparation methods are used. From this standpoint, supported gold catalysts behave differently from platinum-group metal catalysts. This paper attempts to summarize major factors which determine catalytic activity and selectivity in supported gold catalysts.

124

2. EXPERIMENTAL 2.1. Preparation of highly dispersed gold catalysts The catalytic properties of Au depend markedly on the preparation method, because it can bring about a large difference in the size of Au particles and the interaction with supports. Until now six methods, including solid-, liquid-, and vapor-phase techniques, have proven effective for preparing active gold catalysts. Amorphous alloys of gold produced by arc-melting can be transformed to highly dispersed gold particles that interact strongly with the oxides formed from the alloy counter metals during catalytic reaction or through oxidizing treatment. A good example is presented by Shibata, et al. for an Au-Zr alloy used to prepare AufZrO2 [6]. Coprecipitation [4] is useful for the preparation of powder catalysts, especially in the case of Au/a-Fe203, Au/CoaO4, Au/NiO, and Au/Be(OH)2, which are extraordinarily active for the low-temperature oxidation of CO. Active AufPiO2 catalysts can also be prepared by coprecipitation, but only in the presence of Mg citrate [7]. Deposition-precipitation [8] is effective for depositing gold with high dispersion on MgO, TiO2, and A1203. This method is applicable to any form of support including beads, honeycombs, and thin films. In principle, support materials are required to have high specific surface areas, most desirably, larger than 50 m2/g. Since gold hydroxide does not deposit at low pH, this method is not applicable to acidic metal oxides having low points of zero charge, for example, SiO2. The surface reaction of [Au(PPh3)](NO3) with the OH groups of freshly prepared metal hydroxides has recently been reported by the group of Iwasawa [9]. This process is especially useful for the preparation of Au/MnOx, which is active for CO oxidation in air. Chemical vapor deposition of organic gold complexes, typically, dimethylgold (m) ~-diketone, is applicable to the widest range of metal oxides [10]. This technique can be easily used with high surface area materials. It can also be used to deposit Au as nanoparticles even on acidic metal oxides, including M CM41 [11]. Co-sputtering [12] is used for preparing clean thin films which are applicable as gas sensors. A target composed of an Au plate and a metal oxide disk is sputtered in 0.4 Pa of oxygen to deposit on a glass substrate heated at 250~ Oblique deposition on a rotating substrate produces columnarstructured porous thin films.

2.2. Catalytic activity measurements Catalytic activity measurements were carried out in a fixed bed quartz reactor of inner diameter 6 ram. The catalyst (usually 200 rag, 125-212 gm fraction) was placed between two glass wool plugs with a bed length of 10-20 ram. For the oxidation of CO and H2, a standard gas containing 1 vol% CO or H2 in air was dried in a silica gel column cooled down to 0~ or -77~ and passed through the catalyst bed at a space velocity of 20,000 h-lml/g-cat.. For the partial

125 oxidation of propylene a mixture of propylene, oxygen, hydrogen, and argon, most frequently, in a volume ratio 1:1:1:7, was passed through the catalyst bed without drying at a space velocity of 2000 h-lml/g-cat.. 2.3. Catalyst characterization Characterization of the catalysts was made using high-resolution transmission electron microscopy (TEM) (Hitachi H-9000), X-ray diffraction (XRD), extended X-ray absorption fine structure (EXAFS), X-ray photoelectron spectroscopy (XPS), temperature programmed desorption (TPD), and Fourier transform-infrared spectroscopy (FT-IR).

3. RESULTS 3.1. CO oxidation Gold supported on metal oxides and hydroxides is active for CO oxidation at temperatures below 0~ irrespective of the type of metal oxide support when gold is deposited as nanoparticles. Table I shows three groups of gold catalysts which are classified in terms of stability, activity, and requirement as the size of gold particles for activity.

Table 1 Classification of supported gold catalysts for low temperature CO oxidation. Group Stability and Support Preparation Size Activity Methods Requirement less stable but A the most active Be(OH)e, Mg(OH)e CP, DP Au clusters at -77 ~C hydroxides D<= lnm most stable and TiO2, Fe203, Co304, NiO active at -77~ semiconducting & reducible CP, DP Au particles metal oxides D<10nm less stable or A1203, SiO2, ZrO2 less active insulating & non-reducible DP, CVD, CP Au particles at -77~ metal oxides D<10nm DP: deposition precipitation, CP: coprecipitation, CVD: chemical vapor deposition B

Gold supported on Be(OH)2 [13] and Mg(OH)2 [14] (group A) exhibits surprisingly high activity even at-77~ but only when gold remains as clusters of size about 1 nm and the supports are in the form of hydroxides. The second group (group B) is represented by semiconducting and reducible metal oxide supports [4, 15]. They are thermally more stable, but have catalytic activities at -77~ a little inferior to those of group A. The third group (group C)

126 comprises gold supported on insulating and non-reducible metal oxides [10, 11]. It is often difficult to deposit gold with high dispersion on these metal oxides, especially on SiO2, however, once gold is deposited as nanoparticles with diameters smaller t h a n 10 nm catalytic activity is observed at temperatures even below-50~ It appears that the gold catalysts of this group are either less stable during exposure to air (Au/SiO2) or less active at low temperatures below -50 ~C (Au/A1203). Table 2 shows the optimum calcination conditions to prepare active AtffMg(OH)2 catalysts [14]. When gold is atomically dispersed over Mg(OH)2 having no Au-Au coordination or gold is deposited on MgO as metallic particles havingAu-Au coordination numbers close to that of bulk gold, it is inactive[16].

Table 2 Catalytic activity, XRD and EXAFS analyses of Au/Mg(OH)2 as a function of calcination temperature. Calc. Temp. Catalytic Activity XRD Au*** ~C CO(rain)* H2(~C)** Coord. No. 200 250 280 300 400

13 113 Mg(OH)2 2000< 67 Mg(OH)2 2.8 720 88 Mg(OH)2 2.4 0 200< Au/MgO 10.4 0 200< Au/MgO 12.0 * Period of 100% conversion at -70 ~C ** Temperature for 50% conversion *** Coordination number of Au-Au bonding (2.73-2.83A) Feed gas: lvol% CO or H2 in air, 2x104 h -1. ml/g-cat.

Debye function analysis (DFA) obtained through computer simulations of XRD patterns for gold species in Au/Mg(OH)2 calcined at 280~ in air gives the probable size and structure distributions of Au for a fresh sample and an aged sample which has lost activity over a period of three to four months (Figure 1) [17]. It is suggested that gold clusters of 13 and 55 atoms with icosahedral and face centered cubic cuboctahedral structures found when Mg is in hydroxide form are responsible for the low temperature CO oxidation. These results may present an interesting correlation with the reactivity changes of Au clusters with different atoms and charges in the gas phase [18]. The coordination number of icosahedral and cuboctahedral clusters of 13 atoms is five and is larger than the values obtained from EXAFS measurements for the samples calcined at 250~ and 280 ~ This discrepancy suggests that smaller gold clusters composed of less than 10 atoms may also exist in fresh active catalysts.

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Figure 1. Mass fractions of clusters corresponding to (a) the DFA of freshly prepared (1 week old) Au/Mg(OH)2 and (b) aged Au/Mg(OH)2 (3 months). The solid line gives the corresponding non-discrete distribution of mass. The influence of preparation methods on the catalytic properties of supported gold is markedly large, whereas it is marginally small on supported Pt [19]. In Table 3 are listed the mean particle diameters of Pt and Au, rate of CO oxidation at 27~ turnover frequency (TOF) based on surface exposed metal atoms calculated from the mean particle diameter and actual metal loading, and apparent activation energy. TOFs for AudiO2 prepared by depositionprecipitation are larger by about one order of magnitude than those for Pt catalysts and by about 4 order of magnitude than those for AufFiO2 prepared by photochemical deposition. The apparent activation energies also provide an interesting contrast: for all Pt catalysts and A u catalysts prepared by impregnation and photochemical deposition they are in the range of 50-60 kJ/mol, while for Au catalysts prepared by deposition-precipitation they are

128 around 20 kJ/mol.

Table 3 CO oxidation over P t f r i 0 2 and Aufrio2 Metal P r e p a r a t i o n Loading Size Method wt% nm DP 1.0 1.3 IMP 1.0 1.4 Pt FD 0.9 2.4 DP 0.7 3.1 tt 1.8 2.7 Au IMP 1.0 ? FD 1.0 4.6

Rate at 27 ~C mol.s-l.g i 1.4x10 -7 1.9x10 -7 2.4x10 -s 6.9x10 -7 5.5x10 -6 1.7x10 -1~ 1.5x10 -10

TOF at 27 ~C s -1 2.7x10 -3 3.8x10 -3 9.2x10 -3 3.4x10 -2 1.2x10 -1 .... 9.6x10 -6

Ea kJ/mol 49 60 53 19 18 58 56

DP: deposition-precipitation, IMP: impregnation, FD: photochemical deposition

3.2. Partial oxidation of propylene Table 4 shows t h at, in the presence of both oxygen and hydrogen, propylene is converted to propylene oxide(PO) with selectivities above 90% over Aufrio2 [20]. Carbon dioxide is a m a i n by-product. Over P d f r i o 2 and PtfriO2, in contrast, propane is m a i n l y produced. It should also be noted t h a t AufriO2 prepared by impre g n a t i o n is poorly active and produces mainly CO2 at higher t e m p e r a t u r e s . The oxidation of propylene in the absence of hydrogen occurs only at t e m p e r a t u r e s above 250~ to produce a l m o s t exclusively CO2, while the oxidation of hydrogen is r e t a r d e d by the presence of propylene. This implies t h a t the presence of hydrogen not only enhances the oxidation of propylene but

Table 4 Reaction of propylene with oxygen and hydrogen over Au/, Pd/, P t f r i 0 2 catalysts Catalyst Metal T Conversion, Selectivity, % PO PO STY a % yield wt% ~C C~I6 H2 PO acetone C~Is C02 % retool/h/g-cat Aufr i o 2 b 1 50 1.1 3.2 >99 1.09 0.18 Aufrio2 c 1 80 0.2 8.9 <10 >70 0.00 0.00 Pdfrio2 c 1 25 57.1 97.7 - 0.4 98 1 0.00 0.00 Ptfri02 c 1 25 12.1 86.6 2 92 6 0.00 0.00 Feed gas: propylene/~e/Ar=lO/lO/lO/70 (vol%), flow rate: 2,000 ml/hr, catalysts: 0.5g. a space time yield, b p r e p a r e d by deposition-precipitation, cprepared by impregnation.

129 also leads to the selective partial oxidation to the epoxide. Dioxygen may be transformed, probably at the perimeter interface of TiO2, through reduction with hydrogen to an active species, possibly hydroperoxo species, that can selectively oxidize propylene adsorbed on the surface of the Au particles. It is surprising that the reaction pathway switches from oxidation to hydrogenation at specific Au metal loadings. Figure 2 shows that while partial oxidation occurs at Au loadings above 0.2 wt%, below this amount the reaction switches to hydrogenation to produce propane[21]. Hydrogenation takes place with higher turnover frequencies at temperatures below 100~ and the selectivity to propane reaches 100 %. This result appears consistent with the result reported by Naito, et al. for Au/SiOe catalyst, where using very low Au loadings, it was found that the hydrogenation of olefin was accelerated by the presence of oxygen [22]. Although the number of gold particles that could be observed by TEM was not large for the catalysts with Au loadings smaller than 0.1 wt%, the main difference between the 0.4 wt% and 0.1 wt% samples appear to be whether the mean particle diameter of Au is above or below 2.0 nm (Figure 3). Taken together these results indicate that gold particles larger than 2.0 nm are primarily responsible for the formation of propylene oxide, whereas those below

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130 2.0 nm catalyze the formation of propane. It is also probable t h a t the presence of Au clusters, smaller t h a n 1.0 nm, which could not be observed by TEM, may contribute to hydrogenation. In order to improve the conversion of propylene while m a i n t a i n i n g selectivity to propylene oxide above 90%, TiO2 or Ti cations were dispersed on high-surface area supports. This i s to obtain a high density of active sites which would be, nevertheless, sufficiently separated from each other. It was expected t h a t higher conversions of propylene and the depression of successive oxidation of propylene oxide to CO2 would result. Both TiO2 deposited on SiO2 (specific surface area, 310 m2/g) and Ti-MCM supports for Au gave slightly improved conversions but with different time-on-stream behavior. TiO2 on SiO2 showed a decreasing conversion with time while Ti-MCM showed gradually increasing conversion reaching a steady s t a t e value after l h [23]. It is likely that the reaction pathways are different for Au supported on TiOJSiO2 t h a n for Au supported on Ti-MCM.

3.3. Spectroscopic

investigations

FT-IR spectroscopy clearly shows that the reactants, CO and propylene, are moderately adsorbed on the surface of Au particles. The presence of water has either no effect or an enhancing effect on the adsorption. The two absorption bands for carbonyl on the Au surface (2106-2118 cm-1 and 2090-2110 cm-1) indicate that there are the neutral and the positively charged surface of Au [24]. The Au atoms at the perimeter interface with the support (TiO2) are probably positively charged and can activate oxygen molecules. The TPD spectra in Figure 4 show that the deposition of Au on TiO2 appreciably increases the intensity of the desorption peak of oxygen. The desorption temperature at around 250~ suggests t h a t the surface oxygen species is weakly adsorbed on the surface. The desorption spectra of CO from AufriO2 show t h a t the amount of CO taken up is increased by Au deposition on TiO2 and t h a t the desorption always takes place as CO2 even at 50~ P r e t r e a t m e n t of the catalyst sample in a stream of He instead of 02 decreases the amount of CO t a k e n up and causes the desorption to take place as CO in the lower temperature region. These results indicate t h a t the uptake of CO involves oxygen adsorbed on the surface of Aufrio2.

4. DISCUSSION The experimental results presented above clearly show t h a t there are three important controlling factors which determine the catalytic properties of gold supported on metal oxides for oxidation. The first is the selection of suitable m e t a l oxide supports. In the complete oxidation (combustion) of CO and nitrogen-containing organic compounds like trimethylamine gold catalysts are

131

Introd. CO pulse

(c)

m I

(d)

CO

(e)

0 Q 0 I=

02 p r e - t r e a t . (b)

x4 I

0

.,

100

I

I

200

300

400

Temperature, ~

Figure 4. TPD for Au/TiO2 (Au loading: 3.3 wt%, DAu: 3.5nm) and TiO2. Pretreatment in O2: (a) Au/TiO2 and (b) TiO2. Pretreatment in 02 or He followed by exposure to CO pulses up to full uptake: (c) Au/TiO2 in 02, (d) Au/TiO2 in He, and (e) TiO2 in 02. Heating rate: 10 ~C/min, rate of flow: 30 ml/min, sample weight: 200mg.

more active than Pt-group metal catalysts. For the combustion of trimethylamine, iron-based metal oxide supports gave the most active gold catalysts. Since iron oxides adsorb trimethylamine strongly, this suggests that metal oxides having a stronger affinity for the reactants provide higher catalytic activity. For CO oxidation, when gold is deposited as nanoparticles, almost all metal oxides can provide activity at temperatures below 0 ~C. In a previous paper [4] we reported that only select metal oxides, the oxides of 3d transition metals of group VIII (Fe, Co, Ni), gave catalytic activity at -77 ~C. However, this was only for the case where gold catalysts were prepared by coprecipitation. Since then, we have succeeded in depositing gold on SiO2 as nanoparticles and have found similar catalytic activities for CO oxidation at temperatures below 0~ Because a simple mechanical mixture of gold nanoparticles with diameter of 5

132 nm in colloidal dispersion and TiO2powder exhibits poor activity [25] and Au deposited on SiO2 exhibits lower activity when calcined at 300 ~C t h a n at 400 ~C [11], a strong interaction between gold particles and the metal oxide support appears to be necessary in the genesis of the low-temperature activity for CO oxidation. A second factor important for activity is the minimization of the size of gold particles. In general, the diameter of gold particles should be smaller than 10 nm for CO oxidation over Au supported on metal oxides. It is very interesting t h a t for CO oxidation over Au/Be(OH)2 and Au/Mg(OH)2 extraordinarily high catalytic activity is observed only when gold clusters of 13 to 55 atoms with icosahedral and fcc cuboctahedral structures are deposited on these metal hydroxides. This requirement as the size and structure of gold needs further investigation. The reaction of propylene with oxygen and hydrogen over Aufrio2 also indicate the existence of a critical size for gold, in this case, at around 2 nm. When gold particles are larger than 2 nm propylene oxide is selectively formed, while, when they are smaller than 2 nm only propane is formed. The results also indicate that hydrogen molecules can be dissociated on the surface of gold in the presence of oxygen. Since propane was not obtained from propylene and hydrogen in the absence of oxygen, oxygen might modify such small gold particles rendering them electron deficient so as to alow them behave like Pd and Pt. The third factor necessary to obtain high activity is to control the interaction of the gold particles with the metal oxide supports. The remarkably large influence of preparation methods on the catalytic properties of supported gold is observed for the oxidation of CO and H2. Over unsupported gold powder and gold deposited on SiO2 by impregnation, the oxidation of hydrogen takes place at lower temperatures than the oxidation of CO, whereas the opposite is the case over gold deposited on metal oxides (including SiO2) by deposition-precipitation and chemical vapor deposition. The turnover frequency of Aufrio2 prepared by deposition-precipitation for CO oxidation is larger by 4 orders of magnitude than that of Aufrio2 prepared by photo-deposition. These differences may arise from different contact structures at the periphery of the gold particles.. While poorly active gold catalysts are composed of spherical gold particles weakly interacting with the metal oxide supports, highly active gold catalysts are composed of hemispherical gold particles attached to the supports on their flat planes. This contact structure presents a longer distance around the perimeter interface of the gold particles. It is likely that the perimeter junction contains the sites for the activation of oxygen for CO oxidation and for the epoxidation of propylene.

133 5. CONCLUSIONS Gold supported on metal oxides and hydroxides has been prepared by several different methods and tested for CO oxidation and partial oxidation of propylene. The conclusions obtained are as follows: 1) The method of preparation has a considerable effect on the catalytic properties of supported gold. Coprecipitation, deposition-precipitation, chemical vapor deposition methods are especially effective for depositing gold as nanoparticles with diameters smaller than 5 nm and with strong interaction with the supports. 2) Over highly dispersed gold catalysts, CO oxidation can take place even at -77 ~C, and propylene oxide can be selectively produced at temperatures around 100 ~C. 3) For CO oxidation over gold supported on metal oxides, the catalytic activity is almost insensitive as to the type of metal oxide support but strongly depends on the strength of interaction with the supports. For the reaction of propylene with oxygen and hydrogen, only titanium-based oxide supports lead to the selective production of propylene oxide. 4) The size effect is dramatic in CO oxidation over Au/Be(OH)2 and Au/Mg(OH)2 and in the reaction of propylene with oxygen and hydrogen over Aufrio2 catalysts. In the latter case, a critical diameter of gold particles is 2 nm; above this value propylene oxide is selectively obtained, while below it propane is favored. 5) The catalytic nature of gold can be widely tuned by the three factors; type of support, size of the gold particles, and the structure of the contact interface between gold and the support.

REFERENCES

1. I.E. Wachs, Gold Bull., 16 (1983) 98. 2. J. Schwank, Gold Bull., 16 (1983) 103, and 18 (1985) 2. 3. M. Haruta, T. Kobayashi, H. Sano, and N. Yamada, Chem. Lett., (1987) 405. 4. M. Haruta, N. Yamada, T. Kobayashi, and S. Iijima, J. Catal., 115 (1989) 301. 5. M. Haruta, Catal. Today, 36 (1996) 153. 6. M. Shibata, N. Kawata, T. Masumoto, and H. Kimura, Chem. Lett., (1985) 1605. 7. M. Haruta, T. Kobayashi, S. Tsubota, and Y. Nakahara, Jap. Pat. 1778730 (1993). 8. S. Tsubota, D. A. H. Cunningham, Y. Bando, and M. Haruta, Preparation of Catalysts VI, G. Poncelet, et al. (Eds.) Elsevier, Amsterdam, pp.227-

134 235, 1995. 9. Y. Yuan, K. Asakura, H. Wan, K. Tsai, and Y. Iwasawa, Chem. Lett., (1996) 755 and Catal. Lett., 42 (1996) 15. 10. M. Okumura, K. Tanaka, A. Ueda, and M. Haruta, Solid State Ionics, 95 (1997) 143. 11. M. Okumura, S. Nakamura, S. Tsubota, T. Nakamura, M. Azuma, and M. Haruta, Catal. Lett., submitted. 12. T. Kobayashi, M. Haruta, S. Tsubota, and H. Sano, Sensors and Actuators, B1 (1990) 222. 13. M. Haruta, T. Kobayashi, S. Tsubota, and Y. Nakahara, Chem. Express, 3 (1988) 159. 14. S. Tsubota, M. Haruta, T. Kobayashi, A. Ueda, and Y. Nakahara, Preparation of Catalysts VI, G. Poncelet, et al. (Eds.), Elsevier, Amsterdam, pp. 695-704, 1991. 15. M. Haruta, S. Tsubota, T. Kobayashi, H. Kageyama, M. J. Genet, and B. Delmon, J. Catal., 144 (1993) 175. 16. D.A.H. Cunningham, W. Vogel, H. Kageyama, S. Tsubota, and M. Haruta, submitted to J. Catal.. 17. W. Vogel, D.A.H. Cunningham, K. Tanaka, and M. Haruta, Catal. Lett., 40 (1996) 175. 18. D.M. Cox, R. Brickman, K. Creegan, and A. Kaldor, Atoms, Molecules and Clusters, 19 (1991) 353. 19. G.R. Bamwenda, S. Tsubota, T. Nakamura, and M. Haruta, Catal. Lett., 44 (1997) 83. 20. T. Hayashi, K. Tanaka, and M. Haruta, Symp. Heterogeneous Hydrocarbon Oxidation, 211th ACS Meeting, New Orleans, March 1996, pp. 71-74. 21. K. Tanaka, T. Hayashi, and M. Haruta, Interf. Sci. Mater. Interconnection, Proc. JIMIS-8, Jpn. Inst. Metals, pp.547-550, 1996. 22. S. Naito and M. Tanimoto, J. Chem. Soc. Chem. Commun., (1988) 832. 23. Y.A. Kalvachev, T. Hayashi, S. Tsubota, and M. Haruta, Proc. 3rd World Congr. Oxid. Catal., September 21-26, 1997, San Diego. 24. F. Boccuzzi, A. Chiorino, S. Tsubota, and M. Haruta, J. Phys. Chem., 100 (1996) 3625. 25. S. Tsubota, T. Nakamura, and M. Haruta, Catal. Lett., submitted.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

135

The Selective Epoxidation of Non-Allylic Olefins Over Supported Silver Catalysts John R. Monnier Chemicals Research Division, Research Laboratories, Eastman Chemical Company, P.O. Box 1972, Kingsport, TN, 37662, USA.

ABSTRACT The epoxidation of non-allylic, or kinetically-hindered, olefins can be carried out using supported silver catalysts. While epoxidation does occur for unpromoted catalysts, the strength of olefin epoxide adsorption leads to low activity and selectivity, as well as irreversible catalyst fouling. The additon of certain alkali metal salts, such as CsCI, lowers the desorption energy of the olefin epoxide, permitting dramatic increases in activity, selectivity, and catalyst lifetime. In the case of butadiene, the addition of an optimum level of CsCI increases activity and selectivity from approximately 1% butadiene conversion and 50% selectivity for epoxybutene to 15% conversion and 95% selectivity, respectively. Epoxidation of butadiene occurs by addition of dissociatively-adsorbed oxygen to one of the localized C=C bonds to form epoxybutene. The addition of oxygen across the terminal carbon atoms does not occur to any measurable extent. The direct participation of molecular oxygen can be ruled out based both on selectivity arguments as well as the kinetic model for the reaction. The kinetics imply a dual site mechanism. One site, which is unpromoted, serves as the site for butadiene adsorption, while the second site, which is promoted, functions as the site for dissociative oxygen adsorption and epoxybutene formation. Epoxybutene and derivatives represent the beginning of several new families of chemicals that were either not available, or were too expensive, to be considered for large-scale, or even fine chemical, production. More than one hundred chemicals have been prepared so far; several of these are in commercial production at the semiworks scale. 1. INTRODUCTION The gas phase epoxidation of ethylene using molecular oxygen to produce ethylene oxide is one of the most successful examples of heterogeneous catalysis to date. In 1995, over 7.62 billion lb. of ethylene oxide were produced in the US alone(I). In fact, ethylene oxide accounts for 40-50% of the total value of organic chemicals produced by heterogeneously-catalyzed oxidation(2). Thus, while the importance of olefin epoxidation is apparent, efforts to selectively epoxidize higher olefins to their corresponding epoxides using molecular oxygen have been unsuccessful. Many explanations have been proposed for the lack of success in this important area, although the most obvious reason is the reactivity of allylic C-H bonds, which are present in the higher olefins that have been tested. The bond dissociation energy of the allylic C-H bond in propylene is 77 kcal/mole, while the vinylic C-H bond in ethylene is 112 kcal/mole (3). Thus, abstraction of one of the

136

allylic C-H bonds in propylene becomes energetically favorable relative to electrophilic addition of oxygen across the C=C double bond. Once abstraction of hydrogen occurs, formation of an epoxide is precluded. Because of the higher bond strengths of the C-H bonds in ethylene, oxygen addition is the energeticallypreferred route. Previous attempts at higher olefin epoxidation have largely ignored the kinetic reactivity of allylic C-H bonds, but focused on modifying the desirable and unique catalytic properties of silver (the ability to activate oxygen so that it electrophilically adds across a C=C double bond) by promotion or selection of reaction conditions to hopefully suppress allylic hydrogen abstraction while not affecting the addition of oxygen to the C=C double bond. In almost all cases, the only effect has been to poison the silver surface with no substantial improvement in selectivity to the olefin epoxide. In this study we show that it is possible to selectively epoxidize higher olefins to their corresponding epoxides if the olefins do not contain reactive allylic hydrogen atoms. Many olefins, including styrene, substituted butadienes, and norbornene have been selectively epoxidized. Because of the usefulness of 3,4-epoxy-l-butene as a new chemical intermediate, most of the data and discussion will involve the selective epoxidation of butadiene. The catalyst composition and overall kinetics of the reaction will be discussed in some detail. 2. EXPERIMENTAL METHODS

Catalyst preparation has been discussed earlier (4,5) in greater detail. The methods of preparation are typical of those used for supported silver catalysts. In the cases where AgNO3 was used as the silver source, water was used as the solvent. The promoter salts in the appropriate amounts were added to the AgNO3 solution before impregnation and reduction of the AgNO3. The volume of the aqueous salt solution was typically in slight excess of the volume required for incipient wetness. If necessary, excess liquid was decanted before tumble drying at approximately 60-70~ The dried precursor was reduced in a flowing H2/N2 stream at temperatures up to 400~ The various combinations of reduction conditions (gas composition, space velocity, temperature ranges, etc.) have been previously discussed (6). Unless otherwise stated, the silver weight Ioadings for all the catalysts ranged from 12-15%. Promoter Ioadings were varied from a few ppmw to several thousand ppmw. In addition to AgNO3 as the silver source, in some cases silver oxalate, Ag2C204, was used as the silver source. In these cases, a solution of ethylenediamine and water was used to dissolve Ag2C204. Promoter salts were added to the Ag2C204 solution before impregnation and calcination. These catalysts were activated by thermal reduction in flowing air at temperatures up to 300~ The supports used for the various compositions were low surface area, fused alpha-AI203 carriers typically used for supported silver catalysts. Surface areas ranged from 0.2-1.0 m/gm with pore volumes between 0.4-0.6 cc/gm; median pore diameters ranged from 1-10 microns. Most catalysts were prepared on shaped carriers, which were sieved to give a narrow range of granules before evaluation. All catalysts were evaluated at steady-state conditions in a one atmosphere flow reactor using gas flows of He, C4H6, and 02, which were delivered by mass flow controllers. For the kinetic dependency studies, CO2 was also added by a mass flow controller, while different levels of epoxybutene and H20 were introduced from a helium-swept vapor-liquid saturator. In a similar manner, furan, epoxybutene, 2,5-dihydrofuran and crotonaldehyde were added to the feed stream

137

during investigation of the reaction mechanism. Analyses of both the feed and product gas streams were made using in-line gas sampling. Comparison of the feed and product streams permitted accurate mass balances at all reaction conditions. The gas chromatograph employed both thermal conductivity and flame ionization detectors hooked in series to give quantitative analyses of all gaseous products. 3. RESULTS AND DISCUSSION Table 1 Comparison of olefin epoxidation reactions over unpromoted 5% Ag/A120 3. Reaction conditions: T = 250~

catalyst weight = 2~

grams, and flow rate = 50 SCCM of

He/olefin/O 2 = 3/1/1.

Reaction

% Select

CH2=CH2+ 02 EO CO2

53 47

CH2=CHCH3+ 02 CO2 PO CH2=CHCHO

92 0 8

CH2=CHCH2CH3+ 02 CO2 BO CH2=CHCH=CH2

91 0 9

CH3CH=CHCH3+ 02 CO2 2-BO Other

98 0 2

CH2=CHCH=CH2+ 02 CO2 EpBT M Oxirane

25 75

% Conv. 12

3~

2.8

The data in Table 1 summarize catalytic activities for epoxidation of a variety of olefins over an unpromoted 5%Ag/AI203 catalyst. These data illustrate the preferential reactivity at the allylic position relative to addition of oxygen across the C=C bond. While the selectivity to ethylene oxide is typical for an unpromoted catalyst, the selectivities to propylene oxide and butylene oxides are non-existent for propylene, 1-butene, and 2-butene, respectively. In addition to small amounts of the selective allylic oxidation products (acrolein in the case of propylene and butadiene in the case of 1-butene), the only products are those of combustion. However, the results for butadiene reveal it is possible to epoxidize this non-allylic olefin at moderate selectivity and activity. What is not obvious from Table 1 is the short-lived nature of this activity. After 2-3 hours of reaction time, activity and selectivity typically decreased to approximately <1% conversion of C4H6 and approximately 50-75% selectivity to epoxybutene. A typical chromatogram of the activity of an

138

unpromoted Ag catalyst is shown in Figure 1. In addition to the epoxybutene, other oxygenated products are formed. Crotonaldehyde, or 2-butenal, is the expected product from the acid-catalyzed isomerization of epoxybutene. The existence of furan and 2,5-dihydrofuran suggests that 1,4-electrophilic addition of oxygen across the ends of butadiene may be one of the reaction pathways. Allylic oxidation of 2,5dihydrofuran is the likely source of the furan. The existence of acrolein, or propenal, indicates that hydrogenolysis of some C4-intermediate has also occurred. In fact, formaldehyde (detected by the TC detector hooked in series below the FI detector) was usually formed in a 1:1 molar ratio along with acrolein. The mechanistic scheme summarized in Figure 2 provides three different reaction pathways to account for all the products. In this mechanistic scheme, 1,2electrophilic addition to either of the localized C=C bonds forms epoxybutene, while 1,4-electrophilic addition across the ends of butadiene forms 2,5-dihydrofuran. This is the pathway favored by Madix (7), who analyzed the thermal desorption products of butadiene chemisorbed on oxygen-precovered Ag(110) surfaces. Madix concluded that oxygen preferentially adds to the ends of adsorbed butadiene to form 2,5-dihydrofuran, which then undergoes oxidation to form furan. In agreement with the results of Madix, Jorgensen (8), using extended Huckel calculations for butadiene interaction on oxygen-precovered Ag(110) surfaces, calculated that addition of oxygen at the 1,4-position is energetically favored over addition at the 1,2-position. The mechanism of Madix and calculations of Jorgensen do not provide any reasonable explanation for the formation of the epoxybutene seen in Figure 1. Thermodynamically, the formation of epoxybutene from 2,5-dihydrofuran is unfavorable by approximately 19.5 kcai/mole. The third mechanistic pathway is direct hydrogenolysis of an unspecified C4H6-O2 type of intermediate to form 1:1 amounts of acrolein and formaldehyde.

~

~O

1,4-electrophilic~-~

fast

o

~ §

2 electiop "ic l addition~ ~3, r- ~ Ag / '

C-C scission HCHO

H ~"""~i"~Oerization H~-+ ~isom ~Combustion O2~" CO2/H20 -.O

Figure 1. Gas chromatogram for unpromoted silver Figure 2. Mechanistic scheme for butadiene epoxidation. The catalyst. See Table 1 for reaction conditions~ arrow widths approximate relative rates of reaction.

139

In order to determine the relative roles of each of these pathways during butadiene epoxidation, different reaction products were added to the feed stream during the epoxidation reaction. These results, summarized in Table 2, indicate that epoxybutene undergoes conversion to form virtually all of the non-selective products seen in Figure 1. In fact, the relative amounts of the conversion products resulting from epoxybutene addition are in reasonable agreement with those displayed in the gas chromatogram of Figure 1. The epoxybutene addition data also revealed only 55% accountability of the epoxybutene which was added to the feed stream. Further, during the period of time that epoxybutene was being added the conversion of butadiene was greatly suppressed, indicating some type of kinetic inhibition by epoxybutene. After removal of epoxybutene from the feed, much of the activity was slowly restored. These results suggested that epoxybutene was the preferred product and that the strong adsorption of epoxybutene resulted in (1), rearrangement to 2,5-dihydrofuran (with subsequent oxidation to furan), (2), isomerization to crotonaldehyde, (3), hydrogenolysis to acrolein, (4), irreversible adsorption leading to fouling of the catalyst, thus giving poor mass accountability, and (5), strong, reversible adsorption resulting in a kinetic inhibition effect. The failure of 2,5-dihydrofuran to produce any epoxybutene also indicated that 1,4addition of butadiene to oxygen was at best only a minor reaction pathway. The data did imply, however, that lowering the desorption energy of epoxybutene from the Ag surface would be critical in improving the overall performance for butadiene epoxidation. Table 2 Decomposition products from the addition of intermediates to feedstream during reaction condition at 250~

Decomposition products are expressed

as percentage of additive introduced into feed. Feed composition is He/C4H6/O2/additive = 3:1:1:0.003. Feed Additive Furan

Crotonaldehyde

2,5-DHF

Epoxybutene

Furan

N/A

0

75

4

Acrolein

0

0

0

4

2,5-DHF

0

0

N/A

0

Epoxybutene

0

0

0

N/A

Crotonaldehyde

0

N/A

6

25

CO2

0

4

16

14

% Conv. of

0

4

24

46

100

100

97

55

Additive % Accountability of Additive

140

GC chromatogram for unpromoted catalyst C4H 6 Conv: 0.8% EpB TM Oxirane Select: 45% (CO2 = 4%) Other = 51%

GC chromatogram for CsCl-promoted catalyst C4H 6 Conv: 13% EpB TM Oxirane Select: 92% (CO2 = 8%)

1 ~__ _11 1

Figure 3. Comparison of gas chromatograms for unpromoted and CsCl-promoted, Ag/AI203 catalysts at similar reaction conditions.

The gas chromatograms in Figure 3 show the effects of CsCI promotion on a Ag catalyst. Both the selectivity and activity are dramatically increased by CsCI promotion. The only detected reaction products were epoxybutene, CO2, and H20. Further, no decline in activity was detected over an eight hour reaction period. This type of promoter effect is not usually seen for most catalyst promoters. Typically, selectivity is increased at the expense of activity. In the case of butadiene epoxidation, however, both selectivity and activity are strongly dependent upon the strength of adsorption of the epoxide. These results are consistent with the hypothesis that the CsCI promoter lowers the desorption energy of epoxybutene, which increases both the selectivity to epoxybutene and the turnover rate for epoxybutene formation. Silver catalysts promoted by Cs salts are also claimed to increase the selectivity to ethylene oxide (9,10). Arguably, the Cs salt promoters have the same effect for ethylene oxide formation as for epoxybutene formation, but because the rate determining step is not linked with desorption of the epoxide, the observed effect of Cs salt promoters is limited to higher selectivity, since lower desorption energy of ethylene oxide decreases the rate of ethylene oxide combustion without affecting the turnover rate. The curves in Figure 4 show the effects of promoter loading vs. activity (% C4H6 conversion) for three different families of CsCI-promoted, silver catalysts. Each curve is for a different support or method of catalyst preparation. These curve shapes are the typical "volcano" curves often seen for promoted catalysts. The optimum level of promoter represents the balance between under-promotion (not all sites are promoted) and over-promotion (surface is poisoned by excess promoter concentration). In order to assess whether a catalyst optimized for epoxybutene formation was active and selective for ethylene oxide formation as well as whether a state of the art ethylene oxide catalyst (11) was active and selective for epoxybutene production, the two different catalysts were evaluated and the data summarized in Table 3. The ethylene oxide catalyst showed excellent performance (even at one atmosphere pressure) for the formation of ethylene oxide, yet was virtually inactive

141

20 Cs added as CsC1 15 -

-

,o'"

-

:

~o,/~u U

'~

/ /!

% C4H 6 Conversion 10

-

.o

i

.. /

\

='

.;

,, |

1

I

PPM Cs Per Gram of Figure 4. The effect of Cs loading on catalytic activities for three different families of catalysts. The selectivities to EpB TM oxirane are 93 - 95% at maximum activities.

Table 3 Comparison of catalysts optimized for ethylene and budadiene epoxidation. EO Catalyst CH2=CH2 + 02 Conversion EO Selectivity

Unpromoted

Promoted

N/A N/A

11.3 88.4

N/A N/A

0.1 75

CH2=CH-CH=CH 2 + 02 Conversion EpB TM Oxirane Seiectivity EpB TM Oxirane Catalyst CH2=CH2 + 02 Conversion EO Selectivity

Unpromoted

Promoted

12 47

0.3 90

2.8 75

21 96

CH2=CH-CH=CH 2 + 02 Conversion EpB TM Oxirane Selectivity

142

Table 4

Epoxidation of other olefins using CsCl-promoted, Ag/AI20 3 catalysts. Reaction

+02

,0oc

r~'-

Conversion

(%)

95

19

85

21

225~ r-~'~ CO2,H20

0

100

250~

95

1.5

92

43

36

4

~N~ +02 245~

+02

Molar

Select. (%)

CH3 ~;~

,•

+02

,~0

+02 225~~.= O ~ ~ ' +02

210~

O~

for the formation of epoxybutene. Likewise, the catalyst promoted for optimum epoxy butene formation gave excellent yields to epoxybutene, but was almost inactive for ethylene oxide formation. This direct comparison of these two catalysts indicate the compositions are quite different, not altogether surprising since the rate determining steps for the two reactions are different; these results clearly indicate the impact of differences in kinetics on promoter requirements for otherwise similar reactions. As stated earlier, it should be possible to epoxidize non-allylic olefins other than butadiene, or even olefins with allylic hydrogen atoms, as long as the allylic hydrogens are kinetically non-reactive. The olefins in Table 4 indicate this is indeed the case. In all cases, the catalysts were promoted with CsCI; the unpromoted catalysts were either inactive or exhibited very low and transient activities. The data for the epoxidation of styrene and 4-vinylpyridine are discussed in greater detail in an earlier patent (12). The epoxidation of styrene over silver surfaces has also been observed by Blum (13) and Hawker, Lambert et al (14), although the catalysts evaluated in Table 4 are much more active and selective than those described by Blum. The transient temperature programmed reaction spectroscopy (TPRS)

143

results described by Hawker for oxygen precovered Ag(111) surfaces also support the formation of styrene oxide; styrene adsorbed on oxygen-precovered Ag(111) at 110~ resulted in the appearance of styrene oxide at a temperature of approximately 500~ Modification of the Ag(111) surface with 0.33ML of CI atoms increased the selectivity from approximately 60% to 93% styrene oxide. It is not possible to assess long term operation from the TPRS results, since there was only one reaction turnover. From our earlier work with epoxidation of butadiene, it is not possible to assess longer term activity based only on initial rates, especially when the formation of olefin epoxide is desorption-limited. Interestingly, Table 4 indicates that the presence of the para -CH3 group in 4-vinyltoluene results in combustion of this otefin to CO2/H20; the GC column used in these experiments was not capable of detecting any small amounts of other oxidation side products. Clearly, there was no formation of the epoxide product. The -CH3 group is both allylic and benzylic to the aromatic ring and is highly reactive towards C-H bond breaking during oxidation. Thus, as in the case of propylene and 1-butene, the existence of any reactive, allylic C-H bond on a substrate olefin will result in extensive combustion in the presence of 02 and a Ag surface. Table 4 also indicates that norbornene can be easily and selectively epoxidized to norbornene oxide over a CsCI-promoted, Ag/AI203 catalyst, in agreement with the transient TPRS data from Madix (15), who observed that norbornene oxide was formed at 310~ when coadsorbed norbornene and atomic oxygen were heated from 120 to 700~ on an unpromoted Ag(110) surface. On the other hand, Cant et al (16) reported that during continuous oxidation of norbornene by molecular 02 over an unpromoted Ag sponge catalyst in a single-pass flow reactor only benzene was formed as the partial oxidation product. No norbornene oxide was detected; Cant also concluded that norbornene oxide was not an intermediate in the formation of benzene. Our data, taken from continuous flow experiments, suggest that unpromoted silver catalysts do not produce any measurable norbornene oxide (most likely because the norbornene oxide initially formed does not desorb from the silver surface), but that CsCI-promoted, silver catalysts are active and selective for norbornene oxide. Comparison of these different studies indicate the difficulties in extrapolating from transient ultra-high vacuum studies to continuous flow experiments, especially when the kinetics of the catalytic reaction may involve product desorption as the slow step. What is interesting is that norbornene does contain a bridgehead C-H bond that is allylic to the C=C bond. However, either due to the puckered nature of the norbornene molecule, which projects the allylic C-H bond of adsorbed norbornene away from the Ag-O surface to give lower reactivity or due to the high strain energy of the resultant pi-allylic structure if C-H bond breaking does occur, the selectivity to norbornene oxide indicates that bridgehead C-H bond breaking does not take place. In contrast, epoxidation of bicyclo[2,2,2]oct-2-ene is only 36% selective to the desired epoxide. The addition of the an additional-CH2-group in the bridging position results in a dramatic decrease in selectivity to the epoxide. The extra -CH2group decreases the puckered geometry of bicyclo[2,2,2]oct-2-ene in comparison to norbornene, which results in the bridgehead C-H bond for bicyclo[2,2,2]oct-2-ene being projected 5-10 ~ more over the Ag-O surface than in the case of norbornene. It is not possible to state whether the more favorable C-H bond angle or the lower strain energy of the allylic structure (after C-H bond rupture) is responsible for the lower selectivity of bicyclo[2,2,2]oct-2-ene. However, comparison of these two bicyclic olefins indicate that relatively small changes in structure can have dramatic effects on selectivities to corresponding epoxides. Prevention of allylic C-H bond breaking is critical for epoxide formation.

144

Table 5 Reaction conditions during kinetic experiments. 1. Catalyst A. CsCl-promoted, Ag/AI20 3 B. 1.00 grams sieved to 0.0788-0.125" diameter 2.

Reaction Conditions A. Temperature: I90~ B. Flow Rate: 300 SCCM to give GHSV = 30,000 hr- 1. (1; = 0.12 sec) C. V= 8.6 cm/sec D. DCE = 1.0 ppm E. For 0 2 dependency, C4H 6 = 0.17 atm E For C4H 6 dependency, 0 2 = 0.17 atm G. For EpB TM oxirane and CO 2 dependencies, C4H 6 and 0 2 = 0.17 atm H. For Arrhenius plots, feed = diluent/C4H6/O 2 = 4:1 :l

3. Reactor Alumium-clad reactor to minimize thermal gradients

As stated earlier, the product distribution during butadiene epoxidation over an unpromoted catalyst indicated that epoxybutene was strongly bound to the Ag surface and that the CsCI promoter lowered the desorption energy of epoxybutene. These observations should be reflected in the steady-state kinetics of the reaction. The data summarized in Table 5 list the steady-state reaction conditions used to determine the reaction orders for the reactants C4H6 and 02 as well the reaction products epoxybutene, CO2, and H20. In all these experiments differential conversions of C4H6 and 02 were maintained and the data fitted to the typical power rate law expression for epoxybutene formation

eepox.vbutene - keca,6 02 eco2X pEps en20 z Application of the power rate law method is straightforward for those reactions in which the reaction products do not inhibit the rate of product formation. In the case of butadiene epoxidation, the partial pressures of epoxybutene, CO2, and H20 inhibit the rate of additional epoxybutene formation. In these experiments, the effects of CO2 and H20 on the rate of epoxybutene formation can be neglected since the molar selectivity to epoxybutene was typically 98-99%, which gave very low and relatively constant amouts of CO2 and H20 in the gas stream. However, the epoxybutene formed in the reactor required the normalization of the rate of epoxybutene formation to account for the inhibition effect of epoxybutene. Finally, one ppm of 1,2-dichloroethane (DCE) was added to the feed stream to maintain constant activity for the 3-4 week period of time over which the kinetic experiments were conducted. The kinetic plots, which are summarized in Figure 5, reveal many different and interesting dependencies. The 02 dependency is first order over the entire 02 pressure range that was investigated, while butadiene is zero order until very low partial pressures of butadiene are encountered, at which point there is a transition to first order. The pressure dependencies for both epoxybutene and CO2 exhibit transitions from fractional negative orders (-0.4 for epoxybutene and -0.5 for CO2) to full negative first reaction orders. The reaction order for H20 is-0.15 order over the full range of H20 partial pressures. These kinetic dependencies provide very clear strategies from both catalyst design and process control standpoints to maximize the formation of epoxybutene.

145 Ln Rate EpB T M Oxirane Formation 0.50 m

0.06

9

~I~

1

~

9

D

9

H20

-0.38 -0.82 -1.26 -1.70

EpB T M Oxirane

=2.14 --2.58 --

CO 2

-3.02 -3.46 --3.90 -9.20

I

-8.33

I

-7.46

I

-6.59

I

-5.72

I

-4.85

I

-3.98

I

-3.11

I

-2.24

I

-1.37

-0.50

Ln Pressure

Figure 5. Kinetic map for epoxybutene formation using conditions described in Table 5. The combination of (1), transition in negative reaction orders from fractional negative to full negative first order for both epoxybutene and CO2 pressures, (2), a positive first order in 02 pressure, and (3), a simultaneous zero order in butadiene pressure is best explained by assuming a dual-site mechanism for epoxybutene formation. One site is associated with unpromoted Ag, while the second site is associated with a Cs-promoted Ag site. The unpromoted site serves as the site for butadiene adsorption, while the promoted site functions as the site for dissociative oxygen adsorption which is incorporated into adsorbed C4H6 to form epoxybutene. Thus, the sites for butadiene adsorption are different from the sites for 02 adsorption and each site is 1/2 order in 02 pressure. Our results agree with those of Lambert (14) and Madix(15) in the type of active oxygen for selective epoxidation. If we apply the "6/7" rule (see Sachtler (17) for explanation) typically cited as evidence for the role of molecular 02 in selective epoxidation of ethylene for the case of butadiene epoxidation, we would not expect selectivity for epoxybutene to exceed "11/12", or 91.7%. In fact, selectivities of 93-96% are typically seen at all reaction conditions. Selectivities of 97-98% are observed at differential conditions and lower reaction temperatures. Therefore, based only upon the observed selectivities to epoxybutene, dissociatively-adsorbed oxygen is clearly the active oxygen in butadiene epoxidation. Further, the kinetic model, which has been derived from the kinetic plots in Figure 5 has been used to very satisfactorily fit a wide variety of reaction data from several different reactor formats, assumes dissociatively-adsorbed oxygen at both promoted and unpromoted Ag sites. The oxygen incorporated into epoxybutene is dissociatively-adsorbed oxygen, not molecular oxygen. Finally, consistent with the earlier statement that promoters lower the desorption energy of epoxybutene from the Ag surface, the Arrhenius plots for both

146

Ln RateEpB Formation 4.00 2.40 0.80 -0.80 -2.40

Promoted

m

- .

.

m

-4.00i -5.40 -7.20 --

'~~

Unpromoted

E[app] = 40.7kcal/mole

-8.80t -10.40 I I I ~ t -12.00 1.90 1.94 1.99 2.03 2.08 2.12 2.17 2.21 2.26 2.30 2.35 1000/'1"(~ Figure 6. Arrhenius plots for EpB TM oxirane formation. Conversions were differential in both cases and feed compositions were diluent/C4H6/O 2 =4/1/1.

CsCI-promoted and unpromoted Ag catalysts are shown in Figure 6. The apparent activation energy for epoxybutene formation is >14 kcal/mole lower than that for the unpromoted Ag catalyst. Epoxybutene is a very versatile and reactive chemical intermediate that has not been prepared before in large, or even reasonable, quantities. Currently, it is available only from chemical supply houses for prices > $10/gram. The molecule is highly functional, with each carbon atom chemically distinct. It even contains an asymmetric carbon atom at the C3-position. Epoxybutene has two separate and reactive functionalities, a C=C bond and an epoxide group. Each of these groups has rich chemistry, both from a polymer viewpoint and a chemical feedstock perspective. The histogram plot in Figure 7 indicates many different structural isomers are thermodynamically accessible from epoxybutene. In fact, the derivative tree in Figure 8 outlines some of the more than 100 compounds that have been prepared from epoxybutene. Key to this tree of chemical compounds is the thermodynamically favored transition of the family of C4H60 structural isomers from epoxybutene to 2,5-dihydrofuran (by 19 kcal/mole) to 2,3-dihydrofuran (by 7 kcal/mole). Even the formation of cyclopropylcarboxyaldehyde (CPCA) becomes thermodynamically favorable enough at elevated temperature, so that high conversions and high selectivities are observed at temperatures above 350~ In fact, this transition from epoxybutene to CPCA is being commercially practiced to provide a rich variety of chemicals to be used in the pharmaceutical and agricultural sectors (18,19). The various classes of products outlined in Figure 8 include addition of oxygen-containing and nitrogen-containing nucleophiles to epoxybutene to yield an almost endless variety of hydroxy ethers and amino alcohols, respectively, having an extremely wide range of chemical properties as well as boiling points and solvating properties. Addition of H20 to epoxybutene gives both 1,4- and 1,2butenediols, providing a novel means of formation for some interesting glycols.

147

Kcal/Mole 40 30

29.7

20

,04//~o 4.75

-20

t

.,3., ~.,.o-----r~-1s.1 1

I~-cH~

-40

/~,/CHO

1

-21.5

-25.6

Compound

Figure 7. Thermodynamic enthalpies of formation of epoxybutene and related isomers. Enthalpies calculated from CHETAH (Chemical Thermodynamic and Energy Release Evaluation Program) at 25~ and ideal gas conditions.

HO~

O

R

t

OR

OH

t,/

CH ~:~OH

t

X

t

''

RHN~

O

H

1 Figure 8. Abbreviated EpB TM oxirane derivative tree. Compounds in boxes are currently being produced or have been demonstrated in pilot plant operation.

148

Addition reactions to the C=C bond include hydrogenation to give epoxybutane, providing an indirect, yet efficient way to prepare this important epoxide (when coupled with hydrogenation) using molecular 02. Halogenation with CI2 and Br2 form the corresponding dihaloepoxybutanes, which are possible components in flame retardents. Another interesting addition reaction is the selective addition of CO2 to form the highly versatile monomer, vinyl ethylene carbonate, which is similar to ethylene carbonate. From 2,5-dihydrofuran, the most obvious derivative is tetrahydrofuran (THF), formed by the hydrogenation of the C=C bond. The C=C bond is also reactive for hydroformylation and olefin metathesis to add additonal functionality and structure. As stated earlier, the next intermediate, 2,3-dihydrofuran (2,3-DHF) serves as the gateway to CPCA family of chemicals. Hydration of 2,3-DHF produces 2-hydroxytetrahydrofuran, which can be readily converted to gamma-butyrolactone, pyrrolidinone (and N-substituted pyrrolidinones). Finally hydrogenation of hydroxytetrahydrofuran yields 1,4-butanediol in high yields. Yet another demonstration of the versatility of epoxybutene comes from the asymmetric center, which has been converted into a number of four-carbon chiral synthons, such as 3-butene-l,2-diol and various derivatives, in >99% enantiomeric purity (20,21). 4. CONCLUSIONS

The epoxidation of non-allylic olefins, or olefins containly kinetically-hindered allylic olefins, using promoted silver catalysts has been demonstrated. The epoxidation of butadiene to form epoxybutene marks the first example of an olefin other than ethylene to be selectively epoxidized at steady state and commerciallyrelevant conditions using gas phase oxygen and heterogeneous silver catalysts. Epoxidation of higher, non-allylic olefins does occur without the use of an alkali metal salt promoter. However, the olefin epoxide is strongly adsorbed to the silver surface and undergoes a number of side reactions as well as surface fouling. The addition of a promoter, such as CsCI, apparently lowers the desorption energy of the olefin epoxide from the surface, permitting both selectivity and activity to dramatically increase. In the case of butadiene, the addition of an optimum level of CsCI promoter to the silver catalyst increases selectivity and activity from about 50% selectivity and 1% conversion to approximately 95% selectivity and 15% conversion, respectively. Catalyst lifetime also increases from less than 3-4 hours to commercially-relevant periods of time. Epoxidation of butadiene occurs by electrophilic addition of dissociativelyadsorbed oxygen to one of the localized C=C bonds to form the epoxide. The addition of oxygen across the terminal carbon atoms to form 2,5-dihydrofuran does not occur to any measurable extent. When it is observed for unpromoted catalysts, 2,5-dihydrofuran is formed from the isomerization of strongly-bound epoxybutene. The direct participation of molecular oxygen addition to the C=C bond can be ruled out based both on selectivity arguments as well as the kinetic model for the reaction. The kinetics imply a dual site mechanism for butadiene epoxidation. One site, which is unpromoted, functions as the butadiene adsorption site, while the second site, which is promoted, serves as the site for dissociative oxygen adsorption and epoxybutene formation. Under normal reaction conditions, the reaction is zero-order in butadiene pressure and first-order in oxygen pressure (each site is 1/2 order in oxygen pressure). Because of the kinetic inhibition effect of epoxybutene, the reaction at high conversion is negative first order in epoxybutene.

149

Finally, epoxybutene and derivatives represent the beginning of several new families of chemicals that were previously either not available or simply too expensive to be considered for large-scale industrial, or even fine chemical, application. More than one hundred chemicals have been synthesized from epoxybutene, and many more are currently being synthesized and evaluated for a wide variety of applications.

ACKNOWLEDGEMENTS

The author acknowledges the efforts of Peter Muehlbauer and George Oltean for assistance in catalyst preparation and reaction kinetics, David Hitch in reaction engineering, and Jerome Stavinoha and Windell Watkins for many meaningful discussions regarding development of this technology. Steve Godleski and Steve Falling, among many, were instrumental in much of the derivative work involving epoxybutene as a new organic intermediate. REFERENCES ,

2. 3. .

5. 6. 7. ,

9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21.

Chem. & Eng. News, April 8, 1996, p. 1. R. K. Grasselli and J. D. Burrington, Adv. Catal. 30, 133 (1981). J. A. Dean, "Lange's Handbook of Chemistry," p. 4.25, McGraw-Hill, Inc., New York, 1992. J. R. Monnier and P. J. Muehlbauer, U.S. Patent No. 4,897,498 (1990). J. R. Monnier and P. J. Muehlbauer, U.S. Patent No. 4,990,773 (1990). J. R. Monnier and P. J. Muehlbauer, U.S. Patent No. 5,081,096 (1992). J. T. Roberts, A. J. Capote, and R. J. Madix, J. Am. Chem. Soc. 113,9848 (1991). B. Schiott and K. A. Jorgenson, J. Phys. Chem. 97, 10738 (1993). A. M. Lauritzen, U.S. Patent No. 4,833,261 (1989). M. M. Bhasin, P. C. EIIgen, and C. D. Hendrix, U.S. Patent No. 4,916,243 (1990). Commercial ethylene oxide catalyst graciously supplied for purposes of comparison. J. R. Monnier and P.J. Muehlbauer, U.S. Patent No. 5,145,968 (1992). P. R. Blum, U.S. Patent No. 4,894,467 (1990). S. Hawker, C. Mukoid, J. S. Badyal, and R.M. Lambert, Surf. Sci. 219, L615 (1989). J. T. Roberts and R. J. Madix, J. Am. Chem. Soc. 110, 8540 (1988). N. W. Cant, E. M. Kennedy, and N. J. Ossipoff, Catal. Lett. 9, 133 (1991). W. M. H. Sachtler, C. Backx, and R. A. Van Santen, Catal. Rev.-Sci. Eng. 23, 127 (1981). Chem. & Eng. News, August 21, 1995, p. 7. D. Denton, S. N. Falling, J. R. Monnier, J. L. Stavinoha, Jr., and W. C. Watkins, Chimica Oggi, May 1996, p. 17. N. W. Boaz and R. L. Zimmerman, Tetrahedron: Asymmetry, 5, 153 (1994). N. W. Boaz, U.S. Patent No. 5,445,963 (1995).

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

151

R e d o x M o l e c u l a r Sieves as Heterogeneous Catalysts for Liquid Phase Oxidations R.A. Sheldon Laboratory for Organic Chemistry and Catalysis, Delft University of Technology, Julianalaan 136, 2628 B L Delft, The Netherlands 1. INTRODUCTION Catalytic oxidation is widely used for the conversion of petroleum-derived hydrocarbons to commodity chemicals [1 ]. Moreover, in fine chemicals manufacture there is increasing environmental pressure to replace traditional stoichiometric oxidations with inorganic reagents such as dichromate and permanganate, with cleaner, catalytic alternatives which do not generate excessive amounts of inorganic salts as byproducts [2]. Catalytic oxidations in the liquid phase generally employ soluble metal salts or complexes as the catalyst. However, solid catalysts have several advantages compared to their homogeneous counterparts, such as ease of recovery and recycling and amenability to continuous processing. Moreover, siteisolation of discrete redox metal centres in inorganic matrices can lead to oxidation catalysts with unique activities and selectivities by virtue of the fact that oligomeriTation of active oxometal (M = O) species to inactive ~t-oxo complexes is circumvented. One approach to designing stable solid catalysts with unique activities is to incorporate redox metal ions or complexes into the framework or cavities of zeolites and related molecular sieves. So-called redox molecular sieves [3, 4], unlike conventional supported catalysts, possess a regular microenvironment with homogeneous internal structures consisting of uniform, well-defined cavities and channels. Confinement of the redox active site in these channels and/or cavities can endow the catalyst with unique activity as a result of strong electrostatic interactions between acidic and basic sites on the internal surface and the substrate or reaction intermediate analogous to interactions with acidic carboxyl and basic amino groups of amino acid residues in the active site of (redox) enzymes. Indeed, the analogy with enzymes can be taken even further: f'me-tuning of the size and hydrophobicity of the redox cavity (see later) can provide unique, tailor-made catalysts by influencing which molecules have ready access to the active site, on the basis of their size and/or hydrophobic/hydrophilic character. Hence, application of the terms 'mineral enzymes' and zeozymes to such catalysts is appropriate [5]. Up until the late seventies attempts to develop redox molecular sieves were mainly limited to the ion-exchange approach (see later). This situation changed dramatically with the discovery, by Enichem scientists in 1983 [6, 7], of the unique activity of titanium silicalite-1 (TS-1) as a catalyst for oxidations with 30% aqueous hydrogen peroxide. Following the success ofTS-1, interest in the development, and application in organic synthesis, of redox molecular sieves has increased exponentially and has been the subject of several recent reviews [8-11 ]. It has even provoked a revival of interest in another approach to producing redox molecular sieves: the so-called ship-in-a-bottle method [ 12-15]. Why use a redox molecular sieve? Although the major motivation stems from the expectation of producing a catalyst with unique activity we note that, in many cases, a stable,

152 recyclable solid catalyst exhibiting the same activity/selectivity as its homogeneous counterpart would be a useful objective. Thus, even trace amounts of 'heavy metals' in aqueous effluent are undesirable which means that catalysts should be recyclable. A good case in point is chromium: the 'chromofobia' which is currently in vogue imposes an essentially zero emission constraint on this metal. 2. OXIDATION MECHANISMS A conditio sine qua non for understanding oxidations of organic substrates with 02, H202 and RO2H catalyzed by redox molecular sieves is a thorough appreciation of the elementary mechanisms of oxidations in the liquid phase [ 1]. One can safely assume that the elementary steps involved in liquid phase oxidations (see Figure 1) do not change fundamentally when the metal catalyst is located in a molecular sieve; only that the relative rates of these steps may change substantially.

Free r a d i c a l c h a i n a u t o x i d a t i o n initiation

RH

diff.controiled

R. + 0 2 RO 2 9+ RH ROaH + M n§ RO2H + M(n-1)+ ROe + RH RO 9+ RO2H 2 RO 2 9 \ 2 CHO 2 9 /

rate iim. slow

(1)

RO 2

(2)

ROaH + Re

(3)

~- RO 2, + M(n+)+ + H +

fast fast fast

termination

R,

=

(4)

RO, + Mn'IOH

(5)

ROH + R,

(6)

ROH + RO 2 9

(7)

2 ROe + 0 2

(8)

~HOH + ~ C = O + ' 0 2 / /

(9)

Catalytic oxygen transfer RO2H + S

M

~--

ROH + SO

(10)

M+SO

(11)

2M--O

(12)

Mars-van Krevelen mechanism M~O+S 2M+O

2

Figure 1. Oxidation mechanisms

153 One aspect which sets oxidation apart from other reactions, e.g. hydrogenation and carbonylation is the fact that there is almost always a reaction (free radical chain autoxidation) in the absence of the catalyst (Reactions 1-3). Moreover, (transition) metal ions which readily undergo a reversible one-electron valence change, e.g. manganese, cobalt, iron, chromium, and copper, catalyze this process by generating alkoxy and alkylperoxy radicals from RO2H (Reactions 4-6). From the viewpoint of selectivity this ubiquitous competing free radical chain autoxidation is often, but not always, something to be avoided: In principle, it can be circumvented by employing H202 or RO2H as the terminal oxidant in an oxygen tranfer process (Reaction 10). Particularly in the context of fine chemicals such reagents can be economically viable. However, as noted above, variable valence metals can also be expected to catalyze the homolytic decomposition of RO2H and H202 via the so-called Haber-Weiss mechanism (Reactions 4 and 5; R = alkyl or H). Moreover, subsequent free radical chain decomposition of RO2H via reactions 7 and 8 leads to the formation of dioxygen which returns the reaction to an autoxidation manifold (Reactions 1-3). In the case of secondary alkylperoxy radicals, termination via the Russell mechanism [16] can even lead to the formation of singlet dioxygen (Reaction 9). Obviously competition between reactions 6 and 7 will depend on the relative concentrations of substrate (RH) and oxidant (R'O2H). In many cases both homolytic and heterolytic pathways afford the same products, e.g. alcohols and ketones from hydrocarbons, which means that results have to be interpreted with care. Certain elementary tests for homolytic pathways need to be performed, e.g. inhibition by a radical scavenger such as Ionol indicates a free radical chain mechanism and loss of yield on flushing with an inert gas suggests the intermediacy of dioxygen in reactions with H20~ or RO2H. More sophisticated 'reality tests' can also be performed to demonstrate the intermediacy of alkoxy radicals in oxidations with RO,H [17]. The 'holy grail' of oxidation chemistry is the design of catalytic systems capable of mediating oxygen transfer from dioxygen, without the need for a sacrificial reductant, i.e. a Mars-van Krevelen mechanism [18] in the liquid phase. Indeed, the confinement of substrate molecules in the micropores of molecular sieves might be expected to create quasi gas phase conditions conductive to such a mechanism at the expense of free radical chain autoxidation (we note, however, that a mechanism involving two metal centres as shown in eq. 12 would be difficult to achieve in a molecular sieve). The active oxidant in catalytic oxygen transfer processes may be an oxometal (M = O) or a peroxymetal (MOOR) species (Figure 2) [19]. It will be readily appreciated that catalytic oxygen transfer may be considered as a special case of the Mars-van Krevelen mechanism. Peroxometal mechanisms are favourable when the metal in its highest oxidation is both a Lewis acid and a weak oxidant;e.g, early transition metal ions with do configurations (Mo v~, W vl and Tirv). The oxidation state of the metal ion does not change during the catalytic cycle; catalysis being due to the Lewis acid character of the metal ion. Strong (one-electron) oxidants, exemplified by later and/or first row transition elements such as Crv~, Mn m, Co tu and Fe I", favour oxometal pathways and/or Haber-Weiss type homolytic decomposition of RO2H. Vanadium(V), being a strong Lewis acid and a reasonably strong (one-electron) oxidant exhibits all three types of activity.

154

( RO,,H

ROM n*

....=

-ROH

v

M.*

=

MoVI

b)

M(n*2)+

-

Zr w

(a)1

\O/

I

[

Ti w

S

M" * OR

RO

V v

WVX

ROMn* + SO

0 =

V v ' C r vl , F e v , R u vx

M .§

Co m

C r vI

Mn m

Vv

Fe m'

E~

1.82

1.48

1.51

1.0

0.771

I I

I

1

TiIV

MoVX

WVl

0.08

0.2

0.03

Figure 2. Peroxometal (a) vs oxometal Co) pathways in oxygen transfer Hence, the oxometal pathway is more complicated than one might assume from Figure 2. Conversion of peroxometal to oxometal species can involve either homolytic or heterolytic cleavage of the O-O bond (Figure 3). If'leak~e' of RO. (or HO- in the case of H202) occurs, the reaction is transferred to an autoxidation manifold (Figure l). Competition between homolytic and heterolytic pathways for oxometal formation will be influenced by many factors, e.g. ligands surrounding the metal, solvent, etc. In short, catalytic oxidations with 02, H202 and RO2H are, from a mechanistic viewpoint, intricately interwoven, including various homolytic and heterolytic pathways. Examples of oxidative transformations involving heterolytic, pemxometal pathways are olefin epoxidation, sulfoxidation and oxidations of various nitrogen compounds. In contrast, allylic and benzylic oxidations and (cyclo)alkane oxidations are typical of oxometal and/or free radical autoxidation pathways, which are difficult to distinguish. Alcohol oxidations may involve peroxometal or oxometal pathways. There are few cases, e.g. stereoselective olefin epoxidation with TiW/RO2H (R ffi H or alkyl) must involve a heterolytic, peroxometal pathway, which are unambiguous.

155

M.~...01~0 R hemolysis= I!(,.~)+.O 1 diffusion....=RO, =

recombination Mn+(~ b,,'f~'ORheterolysis /?~") ~,,~-o/ = 0 OR Figure 3. Hemolysis vs heterolysis ofperoxomelal complexes

3. REDOX MOLECULAR SIEVES: GENERAL CONSIDEI~TIONS 3.1. Structures and composition of molecular sieves Zeolitcs and zeotypes are crystalline oxides comprising comer slmfing TO4 tetrahedra (T = Si, AI, P, etc.) and consisting of a regular pore system with diameters of molecular dimensions, hence the term molecular sieve. Zeolites refers to altmfinosilicates (T = Si and AI) and zeotypes to molecular sieves having analogous sm~tures but with a different elemental composition, e.g. aluminophospha~s (AIP0's; T = AI and P) and siticoaluminophosphates (SAPO's; T = Si, AI and P). Molecular sieves are categorized on the basis of their pore diameters into small pore (< 4 A~, medium pore (4-6 A), large pore (6-8 A), extralarge pore (8-14 A) and mesoporous (15-100 A). The pore system may be one, two or three dimensional. This can be important in the context of catalysis as a few obstructions in a one dimensional pore system would seriously impede access to a large proportion ofthe catalytic sites while in two or three dimensional pore systems alternative diffusion paths can be found. A molecular sieve having a particular topology is described by a mnemonic three letter code [201. The AFI structure (7.3 ~), for example, is one dimensional while the FAU structure (7.4 A) consists of three orthogonal channel sytems (7.4 A) intersecting in larger cavities (13 A), so-called supercages, and molecules can travel in all three directions. Other selected examples are given in Table 1.

156 Table 1 Pore dimensions and dimensionalities of molecular sieves Pore size Small Medium

Structure type LTA

Trival name

Pore Ringdiameter [/~] size

Zeolite A

4.1

8

3 3 3

MFI MEL AEL

ZSM- 5, TS-I ZSM-11, TS-2 AIPO4-11

5.3 x 5.5 5.3 x 5.4 3.9 x 6.3

10 10 10

MOR BEA FAU

6.5 x 7.0 7.6 x 6.4 7.4

12 12 12

7.4 x 6.5

12

2

LTL AFI

Mordenite Zeolite beta Faujasite Zeolite X or Y Hexagonal faujasite Linde type L AIPO4-5

7.1 7.3

12 12

1 1

VFI CLO

VPI-5 Cloverite

12.1 13.2

MCM-41

ca. 40

Large EMT

i

Extra large Mesoporous

Dimensionality

'

18 20

1

I I

1

3 1

The framework of molecular sieves is not completely rigid and incoming molecules are able to induce slight structural changes. Hence, ca. 10% should be added to the pore diameters given in Table 1 to obtain the limiting sizes of molecules having access to the pores. Their well-defined pore systems combined with their capacity for at least small substrate-induced structural changes enable molecular sieves to recognize, discriminate and organize molecules with a precision of < 1 A [21 ]. This capability to organize and discriminate molecules with high precision endows them with shape selective properties [22], analogous to enzymes. Hence, also by analogy with enzymes one would expect the highest activity to be observed with the best fit, i.e. when the dimensions of the substrate are comparable to those of the micropores. The recently discovered mesoporous (alumino)silicates, e.g. MCM-41, consist of a regular array of uniform one dimensional pores with diameters in the range 15-100 A and have properties intermediate between those of amorphous SiO2 and A1203 and microporous sieves [23]. This has considerably extended the size of molecules that can be adsorbed: the immobilization of enzymes in MCM-41 has even been achieved [24]. 3.2 A c i d i t y a n d h y d r o p h o b i e i t y

In zeolites the different valencies of Si (tetravalent) and AI (tdvalent) produce an overall negative charge for each A1 atom which is balanced by an alkali (alkaline earth) cation. Exchange of these cations by protons affords strong Brvnsted acids, comparable in strength to sulfuric acid. The acid strength increases with decreasing Al content, while the number of acid sites decreases. Substitution of tfivalent aluminium in the zeolite framework by tetravalent

157 species, e.g. Si or Ti produces (metallo)silicalites with an electrically neutral, hydrophobic framework. Likewise, substitution of silicon by phosphorus produces the electrically neutral, hydrophilic aluminophosphates (AIPO's) or acidic SAPO' s.

3.3 Types of redox molecular sieve We can distinguish three types ofredox molecular sieve on the basis of the method of synthesis: (a) ion exchange, Co) framework substitution and (c) encapsulation. As noted above, the negative charges of zeolite and SAPO frameworks are compensated by exchangeable cations, so that redox cations can be introduced by direct ion exchange. However, the high mobility of the exchangeable cations results in a marked propensity for leaching. Moreover, ion exchange is not applicable to neutral molecular sieves such as AIPO's and silicalites. Framework substitution of A1, Si or P by a redox metal ion leads, in general to more stable redox molecular sieves (Figure 4). So-caUed isomorphous substitution, in which the metal ion is coordinated tetrahedraUy by four oxygen atoms should be possible when the r~io~/roxys~ ratio is between 0.225 and 0.414 [25]. It should be noted, however, that the oxidation state of the metal and, hence, structure and charge of the framework, may change substantially when an as-synthesized material is calcined. For example, Cr-substituted sieves generally contain Crm in the as-synthesirzd material but on calcination this is transformed to Cr vl. Since the latter contains two extra-framework Cr = 0 bonds it can only be anchored to a surface defect site rather than isomorphously substitutexl. By the same token, as-synthesized CrmAPO contains a neutral frwnework (Crm replaces AIm) while in the calcined CrVIAPOthe fraraework contains one Brznsted acid site per Crv~. Hydrophobicity

I

205 !

I AI.

I I

ALPOs

I

M-APO

T

I !

SAPOs

Redox MetaI,M

! s ~)2

I [

Y

Fe

Ti

Zr

III

IV

IV

III

IV

IV

If

ZEOLITES

l

Ti-beta

V

IV I IV IV I

V

Cr

9

ii T

TI-ZSM-5

Sn I I

']

m

i

J IM-SILICALITES I

M-ZEOL

-TAPSO -CrAPSO

-VAPO -CrAPO -CoAPO

As-synthesized Calcined

~

Mn

Co

III

II

II

VI

III

III

TS-1 VS-1 CrS-1

I

Figure 4. Types of redox molecular sieves and oxidation states of the metal ions

158 Another approach to creating novel redox active molecular sieves involves the encapsulation of transition metal complexes in intrazeolitc space: the so-called ship-in-abottle method [ 12-15]. Encapsulated metal complexes should, in principle, ideally combine the advantages of homogeneous and heterogeneous catalysis. Molecular sieves containing supercages, e.g. FAU and EMT (hexagonal faujasite) are ideal for encapsulation as substrate molecules have ready access via the micropores (7.4 A) to the metal complex which is trapped in the supercages (13 A) An advantage of encapsulation is that it allows for the synthesis of redox molecular sieves containing elements that are too large to be incorporated by isomorphous substitution (see earlier). 3.4 Choice of solvent Redox molecular sieves have one important advantage over other heterogeneous catalysts: it is possible to influence which substrate molecules approach the active site by a suitable choice of molecular sieve and solvent. The molecular sieve can be viewed as a second solvent which extracts the substrate molecules from the bulk solvent. Which molecules are selectively extracted depends on the size and hydrophobicity of the micropores and of the substrate. Highly siliceous molecular sieves, such as silicalite-1, are hydrophobic and will selectively adsorb apolar hydrocarbon substrates. This is especially important in hydrocarbon oxidations where the primary products - alcohols, aldehydes and ketones - are polar molecules and are more susceptible to oxidation than the hydrocarbon substrates. Hence, by using a redox molecular sieve it is, in principle, possible to obtain much higher selectivities to primary oxidation products than with analogous homogeneous catalysts. The selective oxidation of alkanes with H202 in methanol solvent, in the presence of TS-1 is, presumably, a manifestation of this effect. Catalytic oxidation with H202 in homogeneous solution are generally strongly inhibited by the water present in the H202 or formed in the reaction. Hence, another important advantage of hydrophobic redox molecular sieves, such as TS-1, is that they are not deactivated by water owing to the preferential adsorption of the hydrocarbon substrate and H202. In general, a hydrophilic solvent, e.g. acetone or methanol, is used to create a single liquid phase with aq. H202 although it has been claimed that this is not necessary [26]. Hydrophilicity of redox molecular sieves increases with increasing aluminium content. Hence, high-alumina zeolites, AIPO's and SAPO's are strongly hydrophilic and selectively adsorb hydrophilic substrates. In this case a hydrophobic solvent should be used to facilitate the adsorption of substrates. Furthermore, it should be noted that the incorporation of AI in silicalites or Si in AIPO's generates Brensted acid sites which may catalyze undesirable sidereactions (see later).

4. SYNTHESIS AND CHARACTERIZATION 4.1 Framework-substituted molecular sieves The so-called hydrothermal synthesis of molecular sieves involves allowing an aqueous gel, containing a source of the framework building elements (AI, Si, P) and a structuredirecting agent (template; usually an amine or a tetraalkylammonium salt) to crystallize in an autoclave, under autogenous pressure, at temperatures ranging from 80 to 200 ~ [27]. Crystallization times can vary from several hours to weeks. Redox molecular sieves are similarly prepared by adding a source of the redox metal ion to the synthesis gel. The as-

159 synthesized material is calcined at ca. 500 ~ to remove the template. As noted above, this can lead to oxidation of the redox metal ion to a higher valence state. In the synthesis of silica-based materials a mineralizer (OH, F) is required to regulate the dissolution and condensation process, i.e. synthesis is generally carried out at high pH. In contrast, (redox) aluminophosphates are crystallized from gels prepared by mixing an alumina slurry with a solution of the redox metal ion in aq. H3PO4 and the template, i.e. synthesis occurs at low pH. Titanium-substituted silica-based molecular sieves, in particular TS-1 (MFI), have been the most intensively studied [6, 7, 9]. This generally involves controlled hydrolysis of a mixture of Si(OEt)4 and Ti(OEt)4 in the presence of the template, the tetrapropylammonium cation in the case of TS-1. Many workers have experienced problems in TS-1 synthesis and the various pitfalls haven been reviewed [9]. Small amounts ofNa § or K § orginating from commercial samples of the template suffice to prevent the substitution of Ti into the framework. Similarly, the presence of F leads to the formation of octahedral extra framework titanium. Following the success of TS-1 a variety of Ti-substituted molecular sieves were prepared by hydrothermal synthesis (Table 2) [28-32]. Furthermore, various redox metals, e.g. V, Cr, Mn, Fe, Co, Cu, Zr, and Sn, have been reportedly incorporated into silicalites, zeolites, A1PO' s and SAPO's [8-11 ] and the list is still increasing. Alternatively, framework substitution can be achieved by post-synthesis modification of molecular sieves, e.g. via direct substitution of A1 in zeolites by treatment with TiCI4 in the vapour phase [34] or by dealumination followed by reoccupation of the vacant silanol nests. Boron-containing molecular sieves are more amenable to post-synthesis modification than the isomorphous zeolites since boron is readily extracted from the framework under mild conditions [35]. Synthesis of framework-substituted molecular sieves via post-synthesis modification has the advantage that it is applicable to commercially available molecular sieves which have already been optimized for use as catalysts.

Table 2 Titanium-substituted molecular sieves Material

Template

Pore size (A)

Ref.

Ti-silicalite- 1 (TS- 1)

Pr4NOH

5.3 x 5.5

6,7

Ti-silicalite-2 (TS-2)

Bu4NOH

5.3 x 5.4

28

Ti-ZSM-48

H2N(CH2)sNH2

5.4 x5.1

29

Ti-beta

Et4NOH

7.6 x 6.4

30

Ti-MOR

none

7.0x6.7

31

Ti-APSO-5

C6HIINH2

7.3 x 7.3

32

Ti-MCM-41

CI6H33(CH3)3NOH

ca. 40

33

Ti-HMS

Cl2H25NH2

ca. 40

33

160 A veritable arsenal of techniques has been mobilized to provide information regarding the structure of redox molecular sieves [9-14]. X-Ray powder diffraction (XRD) provides an immediate check for crystallinity and structural type. X-Ray absorption fine structure spectroscopy (EXAFS) and X-ray absorption near edge spectroscopy (XANES) give further insights into coordination geometry and bond lengths. Infrared and Raman spectroscopy have been used to identify characteristic features, e.g. the 960 cm -~ bond attributed to the Si-OTi stretching vibration in TS-1 [9]. Diffuse reflectance UV-Vis spectroscopy (DREAS) and EPR provide useful information regarding the oxidation state of the metal. Other techniques that are regularly applied are MAS-NMR, X-ray photoelectron spectroscopy (XPS) and scanning and transmission electron microscopy (SEM and TEM). Finally, BET surface area measurements and adsorption experiments are indispensible for checking the structural integrity of the molecular sieve. 4.2. Molecular sieve-encapsulated metal complexes Three different approaches are used to achieve encapsulation: a) Intrazeolite complexation (flexible ligand method) b) Intrazeolite ligand synthesis c) Metal complexes as templates for zeolite synthesis In the first method the metal complex is assembled in the zeolite cavities by allowing the metal-exchanged zeolite to react with ligands that are small enough to access the micropores. The metal complex, once formed, is too large to diffuse out. For example, bis- or trisbipyridyl complexes of FeII, Ru n, Mn H, Co t~and Cun have been encapsulated in zeolite Y (FAU) [ 12-15, 36], Metal-Salen and related Schiff's base complexes have been similarly encapsulated in faujasites [12-15, 37, 38]. However, in this case there is virtually no difference in kinetic diameter between the complex and the free ligand and metal-Salen complexes are readily leached by protic solvents, such as ethanol [ 12]. In the second method the ligand itself, constructed by intrazeolite synthesis, is too large to exit the supercages via the micropores. Most examples of this type pertain to FAUencapsulated metallophthalocyanines, first reported by Romanovsky and coworkers in 1977 [39]. They are prepared by first introducing the metal into the zeolite and then adding 1,2dicyanobenzene, which reacts at elevated temperatures to form the metallophthalocyanine in the supercages. Different methods have been used to introduce the metal ion: via ion exchange or as a metal carbonyl or metallocene [12]. In the former two cases the phthalocyanine ligands are largely metallated but many uncomplexed metal ions are also present. In the rnetallocene method, in contrast, there are no uncomplexed metal ions present but a large proportion of the encapsulated phthalocyanine ligands are metal-free. By using substituted 1,2-dicyanobenzenes encapsulated analogs of substituted metallophthalocyanines can be prepared [ 12]. In the template method the zeolite is allowed to crystallize around the metal complex which is assumed to act as a structure directing agent, i.e. the bottle is built around the ship. This allows for the encapsulation of well-defined complexes without contamination by the free ligand or uncomplexed metal ions (see above). The method is restricted to metal complexes that are stable under the relatively harsh conditions of temperature and pH involved in hydrothermal synthesis. Balkus and coworkers [14, 40, 41] used this approach for the encapsulation of metallophthalocyanines in faujasite. However, in order to fit into the faujasite supercages the phthalocyanine ligands are strongly deformed and Jacobs has

161 expressed some doubt [ 12] regarding the structure-directing capability of a template that requires initial deformation. The characterization of zeolite-encapsulated complexes is by no means simple and the same battery of techniques (see earlier) has been brought to bear [14] as with frameworksubstituted sieves. 5. CATALYTIC OXIDATIONS- FRAMEWORK-SUBSTITUTED MOLECULAR SIEVES 5.1. Ti, Zr, Sn and V Titanium(IV) silicalite (TS- 1), the first example of a framework-substituted redox molecular sieve, catalyzes a variety of synthetically useful oxidations with 30% aqueous hydrogen peroxide under mild conditions (see Figure 5). Examples include phenol hydroxylation [42], olefin epoxidation [43], cyclohexanone ammoximation with N113/I-I202 [44], secondary alcohols to ketones [42], primary amines to oximes [45], secondary amines to hydroxylamines [46], sulfides to sulfoxides [47] and alkane oxygenation [48]. The remarkable reactivity of TS-1 is believed to be largely due to site-isolation of tetrahedral titanium(IV) in a hydrophobic environment. The latter ensures that hydrophobie substrates will be adsorbed from a reaction medium containing large amounts of water.

OH

O

NOH

11

1 : 1 o:p

~'

R~/ "- C H 2

L~

o NH a

R2CHOH + ~.

[ :

/R,S

I~-1

+

~ J

R2NH ~ , " =_ R2NOH

~RR'CHNH, RR'CHOH

RzSO

RR'C-'-NOH RR'C~O

Figure 5. TS-1 catalyzed oxidations with H20~

162

The TS- I catalyzedhydroxylationof phenol to a I:I mixture of catcchol and hydroquinonc has been commercialized by Enichem. Similarly,the arnmoximation of cyclohcxanone is being developed commercially as a low-saltalternativeto the conventional process for the production of cyclohexanone oxime, the raw materialfor nylon-6. The reaction involves initialTS-I catalyzedoxidationof NHa by H20~ to give NH2OH. The factthat bulky kctones such as cyclododecanone undergo ammoximation is consistentwith subsequent reaction of N H ~ O H with the ketone substratetaking place outsidethe molecular sieve.The method has been used [49] for the conversion of p-hydroxyacctophenonc to the corresponding oxime which is the precursorof the analgesicparacetamol (Reaction 13). NHCOCH

NOH

1

NH3/H202

(13)

TS-I OH

OH

OH

TS-1 is a particularly active catalyst for olefin epoxidation [43], even unreactive olefins such as allyl chloride being smoothly epoxidized at temperatures close to ambient. Relative reactivities of olefin substrates are completely different to those observed in analogous homogeneous systems. Owing to the steric restrictions of the micropores of TS-1 only straight-chain olefins are smoothly epoxidized. Cyclohexene is completely unreactive (see Table 3). Similarly, in contrast to homogeneous titanium catalysts, TS-1 shows no enhanced reactivity towards allylic alcohols indicating that there is no coordination through the OH group. Table 3 TS-1 catalyzed epoxidations with H202a Olefin

T (~

conv. (%)

Epoxide sel. (%)'

H202

Propene

40

72

90

94

l-Pentene

25

60

94

91

1-Hexene

25

70

88

90

Cyclohexene

25

90

10

n.d.

Allyl chloride

45

30

98

92

Allyl alcohol

45

35

81

72

.

_

t (min)

.

.

.

.

.

_

"MeOH solvem; olefin~202 molar ratio = 5; data taken from M.G. Clerici and P. Ingallina, J. Catal., 140 (1993) 71.

163 The solvent of choice is methanol which gives higher rates than aprotic solvents [43]. This is attributed to the formation of a titanium(IV)-hydroperoxo comples (I) in which coordination of the alcohol promotes oxygen transfer to the olefin (Figure 6). Coordination of the alcohol becomes increasingly difficult with i n ~ g steric bulk, consistent with the observed decrease in rate methanol > ethanol > tert-butanol.

sio

\

SiO --Ti--OSi / sio

SiO

R

SiO __XTiJ~H SiO

/C~~C~'~H

H202

SiO HOSi

\ ~- S i O - - T i

OH

ROH

sio/ \o /

c--c

sio \

/o\ c--c

=- SiO - - / T i - SiO

OR

Figure 6. Mechanism of TS-I catalyzed epoxidation Similar heterolytic mechanisms can be envisaged for other nucleophilic substrates, e.g. ammonia, amines, sulfides, phenols, alcohols. With alkanes or aromatic hydrocarbons, on the other hand, homolytic mechanisms, with possible involvement of HO. radicals, would seem more likely. A titanium(IV)-silicalite-2 (TS-2) catalyst having the MEL topology gives similar reactions to TS-I [50]. However, the scope of TS-1 and TS-2 catalyzed oxidations is limited to the relatively small molecules which can access the micropores (5.6 x 5.3 A and 5.3 x 5.4 A, respectively). This stimulated several groups to investigate the incorporation of titanium into larger pore sieves. Thus, Corma and coworkers [30] s u ~ e d in incorporating titanium in zeolite beta. The resulting titanium~silicoaluminate, Ti-BEA, catalyzed the oxidation of larger substrates such as cyclohexene and cyclododecane [51]. However, owing to the Br~rnsted acidity of Ti-BEA, the major products of olefm oxidation were glycols and glycol monomethyl ethers resulting fTom ring opening of the epoxide by H20 or MeOH, respectively. We subsequently showed [52] that ring opening could be suppressed by simply neutralizing the Brvnsted acid sites by ion exchange with alkali metal ions (see Table 4). Recently, aluminium-free titanium-substituted beta was synthesized, using a different template, and shown to catalyze the epoxidation of olefms with H20,, albeit with some ringopening [53].Ti-BEA also catalyzes epoxidations with TBHP [54], in contrast to TS-1 which cannot accommodate the transition state for epoxidation with the bulky TBHP. Ti-APSO-5, which also contains Brznsted acid sites afforded the diol as the main product in the oxidation of cyclohexene with H202 while with TBHP the epoxide was formed in 79% selectivity [55].

164

Table 4 Titanium-catalyzed epoxidations Olefin

Catalyst

Oxidant

, Conv. ]

Res

Selectivity (%)

(%) /

m

1=Hexenea

m

H202

TS-I Ti-beta

Cyclohexene

i

TS- 1 Ti-beta

,,....

98 80

H202 i

.,

,

epoxide

glycol(ether)

96 .12

4 88

nn

,,

:

l

m

,

,,=,l,

l

<5

100

-

H202

80

-

100

.

,

el

,,,

51

H202 ~

,=

51

|

i

1-Octene"

Li-TS-I Ti-beta Li-Ti-beta Na-Ti-beta

M H202 H202 H202 m

1-Octene b

Ti-beta Li-Ti-beta

Cyclohexene ~

TAPSO-5 TAPSO-5

m

,,

TBHP TBI-IP m

"MeOH; 40-60~

b

85 48 31 22

H202

,

98 O 87 84 9

m

47 38 |

0 97 5 6 ,,,,,, ,,

38 100

39 0

0 79

79 12

i !

,

'54 i

,|

H202 TBHP

,|

!

55

CF3CH2OH; 90~ c Me:CO/70 o.

In some cases the presence of both redox and Breasted acid sites can be an asset as it can provide the possibility of perfom~g bifunctional catalysis. For example, Ti-BEA and TiMCM-41 (see later) catalyzed the epoxidation of linalool with TBHP, with in-situ acidcatalyzed rean~gement to a mixture of cyclic ethers (Reaction 14) [56].

OH TBHP

OH v

CHaCN;80~

-I-

H

Catalyst

Furan/pyran

Ti-BEA Ti-MCM-41

1.62

0.89

(14)

165 The titanium-substituted mesoporous zeolite, Ti-MCM-41 has recently been synthesized and shown to catalyze oxidations of bulky substrates with H~O2or TBHP [57], e.g. reactions 15 and 16.

TBHP

Ti-MCM-41 CHzCI z " 4 0 ~ "5h

_

(15)

Conv. 30% Sol.

OH

90%

O

30% H20~, Bu~/But

(16)

Ti-MCM-41 o Conv. 20% Sel.

98%

I n short, the development of T S - 1 and related titanium-substituted molecular sieves has opened up a whole new area of selective oxidation chemistry with H202 and TBHP. For Some bulk chemicals, e.g. propylene oxide, the relatively high price of H=Ozcould result in unfavourable economics. This led Enichem workers [58] to investigate the possibility of insitu generation of H202. They showed that TS-1 was an effective catalyst for the epoxidation of propylene with H~O~ generated in-situ by autoxidation of a dihydroanthraquinone (Reaction 17). The antiuaquinone coproduct is hydrogenated in a separate step (Reaction 18) resulting in an overall two-step process for the epoxidation ofpropylene with H2 and 05.

OH

§247

Ts.,

+

+ H~.O

(17)

OH

OH

O + H2

O

PdlC

(18) OH

The success of titanium-substituted molecular sieves stimulated the investigation of

166 other metal-substituted silicalites, e.g. zirconium [59] and tin silicalites [60] with the MFI structure have been synthesized. They were shown to catalyze the hydroxylation of phenol with H202, albeit with lower activities and selectivities than TS-1. The incorporation of vanadium in silicalites and aluminophosphates has also been widely investigated [9-12]. The as-synthesized materials contain vanadium(~) which is converted to the pentavalent state on calcination. The resulting materials appear to contain vanadium(V) attached to the framework at defect sites [9 ]. V-MFI and V-MEL silicalites have been shown to catalyze the hydroxylation of aromatics and alkanes with H202. However, H202 efficiencies are generally low and, owing to the relatively high redox potential of the V v / v iv couple, homolytic mechanisms, involving intermediatehydroxyl radicals,probably predominate, as is observed in homogeneous vanadium systems [61].V-APO-5 was found to catalyze epoxidations and benzylic oxidationswith T B H P [62].However, more recent studies have revealed thatcatalysisis probably due to leached vanadium [63]. 5.2. Co, Cr, Mn and Fe Typical one=electron oxidants such as Mnm, Com, and Fem do not generally react via peroxometal pathways and one would not expect the same type of reactivity as that observed with titanium. ,4 priori one might expect free radical chain autoxidation pathways analogous to homogeneously catalyzed oxidations with these metals [1]. This appears to be borne out in practice, e.g. CoAPO-5 catalyzes the autoxidation of cyclohexane in acetic acid [64]. More careful examination revealed that cobalt is leached by the acetic acid and that the observed catalysis was homogeneous [65]. In contrast, it was claimed [65] that, in the absence of acid, the integrity of CoAPO and CrAPO (see later) is preserved in the autoxidafion of cyclohexane at 130 ~ However, it is worth noting that adipic acid is formed in these reactions which can also be expected to leach the cobalt (or chromium) at conversions higher than a few percent. Chromium-substituted molecular sieves contain Crv~after calcination, which might be expected to give reactions typical of oxometal pathways with ROzH, as is observed in solution [66]. Indeed, this proved to be the case: we found that CrAPO-5 effectively catalyzes benzylic [67], allylic [68] and alkane [67] oxidations, oxidation of secondary alcohols to ketones [69] and the selective decomposition of secondary alkyl hydroperoxides to the corresponding ketones [70]. The stoichiometric oxidants are either TBHP or 02 (Figure 7). Similar reactions are observed with CrAPO-11 and chromium silicalites with MFI or MEL structure [71].

R~,C--O

R2CHO2H R2C--O O~=R,CHOH

ArCOR ArCH,R~Oz

Figure 7. Selective oxidations catalyzed by CrAPO-5

167 For example, a series of alkylaromatics afforded the corresponding aralkyl ketones in high selectivities with TBHP in the presence of CrAPO-5 (3 m %) at 80 ~ (Table 5). TBHP could be replaced by Oz but this required neutralization of Brensted acid sites on the CrAPO5, by ion-exchange, in order to avoid acid-catalyzed decomposition of the benzylic hydroperoxide to the corresponding phenol, which inhibits the autoxidation. The addition of a small amount of TBHP to initiate the reaction also had a beneficial effect.

Table 5 CrAPO-5 catalyzed oxidations of alkylaromatics with TBHP [67] Conv. (%)a

ArCOR sel. (%)a

Ethylbenzene

70

90

p-Ethyltoluene

68

97

n-Propylbezene

59

93

n-Butylbenzene

59

92

Diphenylmethane

50

94

3-Ethylpyridine

43

80

ArCH2R

"Based on substrate.

The oxidation of hydrocarbons and alcohols with TBHP is presumed to involve oxidation of the substrate by oxochromium(VI) followed by reoxidation of Cr TM by TBHP, i.e. an oxometal pathway. When 02 is the oxidant the substrate undergoes initial chromiumcatalyzed autoxidation to the corresponding hydroperoxide. The latter undergoes catalytic decomposition and/or functions as the oxidant. The observation that the bulky triphenylmethyl hydroperoxide, which cannot be accommodated in the micropores, gave no reaction was construed as evidence for the reactions taking place in the micropores. This subsequently proved to be a misinterpretation of the results (see later). Manganese- and iron-substituted aluminophosphates have also been synthesized but they have proven to be particularly unreactive oxidation catalysts [72]. This lack of redox activity may be a result of the stable environment of isomorphously substituted Mn m and Fe tII. Interest in iron-substituted molecular sieves is derived from he fact that many redox enzymes contain iron in their active site. Hence, with this element one is obliged to adopt the alternative approach of encapsulation. 6. CATALYTIC OXIDATIONS: ENCAPSULATED COMPLEXES Pioneering studies of zeolite-encapsulated iron phthalocyanine (FePc) complexes were performed by Herron [73] who coined the term ship-in-a-bottle complex. He studied the 'oxidation of alkanes with iodosylbenzene catalyzed by FePc encapsulated in zeolites Na-X

168 and Na-Y. Although rather interesting (shape) selectivities were observed, reminiscent of heme-containing enzymes, activities were poor (ca. 6 turnovers) which was attributed to pore blockage by iodoxybenzene. After laying dormant for a number of years, Jacobs and coworkers re-initiated these studies using TBHP as the oxidant [11-13]. They found that FePc-Y, containing 1 FePc per 77 supercages, gave turnovers as high as 6000, compared to 25 with homogeneous FePc, in the oxidation of n-octane with TBHP [74]. It was important that the catalyst be prepared by the metallocene method (see earlier) in order to avoid the presence of uncomplexed iron. The catalyst eventually deactivates, presumably owing to ligand degradation (the TBHP should also be added slowly to minimize degradation). More recently, this was extended to the use of EMT [75] the hexagonal polymorph of faujasite and the 18membered ring aluminophosphate VPI-5 [74, 76] as the host. A further increase in catalytic activity/stability was obtained by substitution of the phthalocyanine periphery with electronwithdrawing nitro groups [77]. Balkus and coworkers [ 14, 78] synthesized a Na-X encapsulated ruthenium perfluorophthalocyanine complex (Ru-Fl6Pc-X) and showed it to be an active catalyst for the room temperature oxidation of cyclohexane with TBHP (Table 6). The catalyst showed no sign of deactivation, in contrast to the iron analogue.

Table 6 Ru phthalocyanine-catalyzed oxidation of cyclohexane with TBHP [78] Catalyst

Time (h)

Conv. (%)

Selectivity (%) ketone

alcohol

TBHP efficiency (%)

TOF (h-!)

RuPc

5

47

72

27

30

7.5

RuFi6Pc

24

83

78

22

48

15

RuFI6Pc-X

192

86

98

97

125

Another recent development is the use of faujasite-encapsulated manganese(II) bipyridyl complexes as catalysts for the oxidation of cyclohexene to adipic acid using 30% aqueous H202 [79]. The reaction proceeds via initial epoxidation followed by ring opening to the diol and oxidative cleavage. The turnover frequency was 15 h ~ and H202 efficiency 47-62% at cyclohexene conversions of 9-17%. Other recent variations on this theme include the encapsulation of VO(bipy)22+ [37], VO Salen [38] and iron(II) N,N-bis-(2-pyridinecarboxamide) [80] complexes in zeolite Y. The vanadium complexes were used as catalysts for the oxidation of cyclohexene or cyclohexane with TBHP. Oxidation of the latter, to a mixture of cyclohexanol, cyclohexanone and cyclohexyl hydroperoxide, presumably involves hydroxyl radicals by analogy with homogeneous systems [61 ]. Moreover, the question of leaching still has to be addressed in these systems. The encapsulated iron complex catalyzed the oxidation of cyclohexane with H202, probably via a homolytic Haber-Weiss-type mechanism.

169 More recently, these systems have been elevated to a new level of biomimetic sophistication, namely by embedding the zeolite-encapsulated complex in a polydimethylsiloxane membrane their performance could be improved even further [81 ]. The hydrophobic membrane mimics the phospholipid membrane in which cytochrome P-450 resides and acts as an interface between the two immiscible phases (cyclohexane and aqueous 70% TBHP). Based on these recent developments ship-in-a-bottle catalysts appear to hold much promise, extending even to the design of chiral 'ships' for enantioselective oxidation. 7. THE QUESTION OF LEACHING An important question regarding the observed catalytic activity of divers redox molecular sieves is the following: is the catalyst truly heterogeneous or is the observed activity due to leached metal ions? A more than cursary examination of the literature reveals that in most cases an adequate verification of true heterogeneity has not been performed. One point which has become clear from our own work is that the standard recycling of the catalyst without significant loss of activity can by no means be construed as evidence of heterogeneity (see later). Another important question to ask is: is it important if trace amounts of metal ion are leached from the catalyst? Obviously if virtuaUy all of the observed catalysis is due to leached metal ion then it would be simpler to carry out the reaction with a homogeneous catalyst. On the other hand if, as is the case with titanium(IV), the metal ion shows little activity in homogeneous solution, and is not considered environmentally hazardous, then leaching of trace amounts may be less of a problem. We recently performed a detailed investigation of the stability of chromium-substituted molecular sieves [82]. As a model reaction we chose the allylic oxidation of a-pinene with TBHP (Reaction 19) which gives verbenone in high selectivity.

I TBHPICrCatalyst PhCI, 80 ~ "-'-

(19)

Experiments were performed in which the catalyst (CrAPO;5, CrAPO-11 and CrS-1) was filtered, at the reaction temperature (80 *C), after 30 minutes, which corresponded to 30-40% conversion. Filtration at the reaction temperature is important in order to avoid possible readsorption of leached metal ion on cooling. When the filtrate was allowed to react further at 80 ~ the reaction continued at roughly the same rate. Further insights into the mechanism of catalysis were obtained from the allylic oxidation of valencene (I) which is too bulky to enter

170 the pores of CrAPO-5. When TBHP was the oxidant smooth oxidation was observed. In contrast, when the bulky triphenylmethyl hydroperoxide (H) was used, essentially no reaction was observed.

OH

I

II

These results can be rationalized by assuming that chromium is leached by reaction with RO2H, which occurs with TBHP but is impossible with the bulky (II) which cannot access the micropores. This was confirmed by performing an experiment in which CrAPO-5 was stirred with a solution of TBHP. After filtration the filtrate exhibited the same activity as was observed with the CrAPO-5 present. When the CrAPO-5 was initially stirred with the substrate, filtered and TBHP added to the filtrate no catalysis was observed. Having established that the observed catalysis was due to leached chromium we quantitatively determined the amount of chromium in the filtrates [83] which was found to vary with the catalyst used. The lowest amount of leaching was observed with CrAPO-5 and corresponded to 0.3% of the chromum present in the catalyst, i.e. 0.3% of 0.88%. In the above mentioned oxidations this corresponds to a substrate/catalyst ratio (S/C) in solution of 17000. Hence, we performed an oxidation of g-pinene with TBHP and homogeneous chromium(VI) (pyridiniumdichromate) at S/C = 17000 and observed essentially the same reaction rate as with CrAPO-5. An important lesson learnt from this work is that with some metals, e.g. chromium, minute amounts of leached metal ions, in our ease 0.3% of the available chromium, can account for the observed catalysis. This means that the catalyg could, for example, be recycled ten or even a hunderd times while still observing the same activity. Hence, we conclude that, in the absence of unambiguous evidence to the contrary, many literature claims for heterogeneous catalysis by redox molecular sieves are, to say the least, questionable. Indeed, solubilization by reaction wih RO2H appears to be widespread phenomenon. For example, van Hooff and coworkers [63] observed this with VAPO catalysts and Schuchardt and coworkers similarly observed leaching with V, Cr, Mn, Fe and Co-substituted MCM-41

[841.

There seems little doubt that the remarkable activity of TS-1 is heterogeneous in nature. We note, however, that small amounts of leaching to give catalytically inactive species cannot be excluded, especially in the presence of NH3 or carboxylic acids. Similarly, the enhanced reactivity and stability of zeolite-encapsulated metallophthalocyanines, compared to homogeneous counterparts suggests that the catalysis is truly heterogeneous, although it would be comforting to see unambiguous verification. The same cannot be said, however, of

171 other encapsulated complexes, e.g. we have recently shown that vanadium-Schiff's base complexes encapsulated in zeolite Y undergo facile leaching by reaction with TBHP [85]. 8. FUTURE PROSPECTS The discovery, in the mid-eighties, of the remarkable activity of TS-1 as a catalyst for selective oxidations with aqueous H202 fostered the expectation that this is merely the progenitor of a whole family of redox molecular sieve catalysts with unique activities. However, the initial euphoria has slowly been tempered by the realization that framework substitution/attachment of redox metal ions in molecular sieves does not, in many cases, lead to a stable heterogeneous catalyst. Nevertheless, we expect that the considerable research effort in this area, and the related zeolite-encapsulated complexes, will lead to the development of synthetically useful systems. In this context the development of chiral ship-ina-bottle type catalysts for intrazeolitic asymmetric oxidation is an important goal. Such an achievement would certainly justify the appellation 'mineral enzyme'. 9. REFERENCES

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

177

Synergistic Effects in M u l t i c o m p o n e n t Catalysts for Selective O x i d a t i o n P. Courtine and E. Bordes, D6partement de G6nie Chimique, Universit6 B.P. 20529, 60205 Compibgne Cedex, France.

de Technologie

de Compi~gne,

Examples of synergistic effects are now very numerous in catalysis. We shall restrict ourselves to metallic oxide-type catalysts for selective (amm)oxidation and oxidative dehydrogenation of hydrocarbons, and to supported metals, in the case of the three-way catalysts for abatement of automotive pollutants. A complementary example can be found with Ziegler-Natta polymerization of ethylene on transition metal chlorides [1]. To our opinion, an actual synergistic effect can be claimed only when the following conditions are filled: (i), when the catalytic system is, thermodynamically speaking, biphasic (or multiphasic), (ii), when the catalytic properties are drastically enhanced for a particular composition, while they are (comparatively) poor for each single component. Therefore, neither promotors in solid solution in the main phase nor solid solutions themselves are directly concerned. Multicomponent catalysts, as the well known multimetallic molybdates used in ammoxidation of propene to acrylonitrile [2, 3], and supported oxide-type catalysts [4-10], provide the most numerous cases to be considered. Supported monolayer catalysts now widely used in selective oxidation can be considered as the limit of a two-phase system. One important fact is that synergistic effects are also known in solid state chemistry, in the absence of reactive atmosphere. For a two-phases AOx/BOy sytem, a physical property of one component BOy, e.g., the temperature at which an allotropic transition proceeds, can be modified by the presence of AOx. In the examples of catalytic synergy we have studied in the past, we have also observed such a "solid state synergy" in inert atmosphere, which cannot be explained by anything else than cooperative transformation due to coherent interfaces. Therefore, the question to answer is: what is the common explanation to these catalytic and non catalytic phenomena, observed as well with oxide/oxide (or chloride/chloride) as with metal/oxide? The oldest and first example in selective oxidation is V205-TiO2(anatase), catalyst of o-xylene oxidation to phthalic anhydride. Nearly in the same time, three teams in Poland, Great-Britain and France [4-10] observed the catalytic synergy and studied also the

178 non-catalytic phenomenon by thermal analysis and X-ray diffraction (XRD). Vrjux and Courtine proposed an explanation by using the concept of coherent interfaces [9], already known to apply in other fields like semiconductivity. This concept, proposed earlier by Ubbelohde [ 11] to account for hysteresis and other non-equilibrium phenomena in other kinds of systems, had soon after been used in the field of reactivity of solids when progress in electron microscopy allowed to evidence extended defects and microdomains formed during the reduction of titanium, vanadium and niobium oxides [ 12]. Since that time, we have made several experiments to check the validity of the model [13-23], which has been extended to other systems than those based on ReO3-type oxides [1, 24-28]. The present paper aims at considering briefly the methodology we used and the main examples we studied. We shall also consider the remote control model proposed by Delmon et al. [29, 30] to interpret the same kind of synergistic effects for similar systems, in order to see whether or not the two models are conciliable and on what points of view. 1. THE CONCEPT OF COHERENT INTERFACES It is well-known in thermodynamics that, when two phases are in physical contact, the surface and interface energies barriers are so high that each phase behaves as if it was alone and keeps its own properties. However, when the crystal structures are closely related, the crystallographic misfit can be very low, and the interface is said to be coherent [ 11, 12]. The atoms close to the boundaries are strained, so that their potential reactivity is greater than when they are inside the pure phases, and the diffusion of various specie (atoms, ions, electrons) across the interface is favored because the interfacial energy barrier is strongly lowered. At short range order (say, 1-10 nm), that is when considering the crystal framework, microdomains of one phase inside (and/or at the surface of) the other can be formed. This is a way to account for non-stoichiometry and solid-solid cooperative reactions, but also for the generation of new active centres in catalytic reactions [9, 15, 28, 31-33]. The turn-over frequency or the specific activity is therefore expected to increase. Microdomains can be formed during the transient state and be kinetically, but not thermodynamically, stable during the steady state. However, other conditions are needed to observe selectivity enhancement, because the additional active sites must be also selective. Several examples of the formation of coherent microdomains have been given usingin situ High Resolution Transmission Electron Microscopy (HRTEM), in the case of V205 or MoO3 reduction as well as during catalytic reactions by Gai [34]. A low crystallographic misfit (few percents) is to be considered as a pre-requisite condition to observe synergy. Finally, it is well conceivable that, since synergistic effects depend strongly on the way the interfaces are "prepared", factors such as the method of preparation and activation of the catalyst (supported or

179 multicomponent), the operating catalytic conditions, the steady or transient state, etc., are preponderant. 2. E X A M P L E S

With an AOx/BOy powder system, and apart the formation of a third phase ABOz in conditions where normally it should not exist, either only one phase reacts and the other acts as an external boundary surface, or both are transformed. Obviously, the observed situation depends on the operating conditions and on the chemical reactivity (and physical properties) of the phase(s), but also on surface characteristics. 2.1. One phase reacts while in contact with the other.

In the first case, one (or more) crystallite face exposed by AOx in a convenient AOx/MOy system may share coherently a face of MOy, provided its surface plasticity is large enough, that is, if the constituents are labile enough to accommodate the misfit. The model of coherent interfaces was applied to "bulk" V205 supported on TiO2 (anatase) to account for the synergistic effect evidenced in the seventies [4-10]. Since that date, the demonstration has been extended to other supports structurally related to anatase [14, 19]. The degree of "plasticity" is related to the Tamman temperature which amounts ca. 450 K for V205, well below the usual temperature of reaction. Several names have been given to this phenomenon, from "premelting" by Ubbelohde [8], to sintering, contamination, wetting and thermal spreading. The latter terms were introduced mostly for monolayer supported catalysts [35, 36], for which the concept of coherence does not seem to be appropriate. However, if the structure of the monolayer does not depend on the specific oxide support MOy (M = Ti, Nb, Zr, A1, Si), its reactivity does since the ease of reduction of VOx depends on the strength of the V-O-support bonds [37, 38]. We have been able to oxidize ethane to acetic acid using VPOx/TiO2 and VOx/TiO2 catalysts at low temperature (200-250~

[21, 22], but when

zirconia or silica were used, no acetic acid was found. The majority of examples of catalytic synergy in selective oxidation and related reactions belongs to this first category and involves ReO3-type oxides and related oxysalts, including rutile structures and compounds containing antimony (VSbO4, Sb204). V205 and vanadiumcontaining catalyts are supported on titania or other structurally related supports, or mixed with Mo and/0r Nb as in the more recent examples of oxidation of ethane to acetic acid [21] or propane dehydrogenation [39]. When TiO2 anatase was used as a support for other oxides (MOO3) or oxysalts like, e.g., CoMoO4, the partial conversion of TiO2 anatase to rutile was observed in non catalytic conditions and n-butane was oxidized to maleic anhydride on the three-phases CoMoO4/MoO3/TiO2 system [20].

180 Metallic molybdates MMoO4 (M= Mn, Fe, Co, Ni, Cd, etc.) are more active and/or selective when MoO3 is present in such an amount so as to be detected by XRD. Several examples have been given by Ozkan et al. [40-42] for oxidative dehydrogenation and oxidation reactions. We studied the system N i M o O ~ o O 3 which is one of the catalysts for propane to propene. The best stoichiometry is Mo/Ni = 1.26/1, and, to observe the highest conversion, the catalyst must be prepared by a special method which leads to a maximum of contacts between NiMoO4 and MoO3 during the calcination [23]. Structural relationships between bismuth molybdates have soon be suspected [38], and complex formulae involving several metallic molybdates have allowed to increase the yield in acrylonitrile or acrolein obtained by (amm)oxidation of propylene [44]. The only way to account for this necessary complexity was to consider that coherent interfaces exist between the main phases because they have close structural relationships [15, 31-33]. The system has been recently investigated by Ponceblanc et al., who measured the change of electronic conductivity according to the number of phases simultaneously present [45]. In other examples, extensively studied by Delmon et al., Sb204 was used with MoO3 for isobutene oxidation to methacrolein [29, 30], and for the dehydration of N-ethyl formamide [46, 47]. Antimony is one of the elements frequently found in selective oxidation catalysts, as in the pionneering work on uranium antimony oxides for ammoxidation of propene [48], and more recently in ammoxidation of propane on V-Sb-A1 system [49]. Courtine et al., have shown that some effects, like the drastic increase of activity in Ziegler-Natta polymerization of ethylene when the catalyst is supported by MgC12 [ 1, 24], or the stabilization of platinum catalysts for abatement of automotive pollutants when the alumina is doped with rare earth oxides [25-27], could also be interpreted by using the same concept. In the first example, the active solid solution TiC13-A1C13 is supported by MgC12. These second generation of catalysts are prepared by ball-milling of the two components in inert atmosphere, and their activity has been increased by (at least) a 103 factor. The optimum depends on the time of ball-milling, and is related to a definite degree of disorder in the stacking planes of TiC13-A1C13crystallites. Both parameters and framework of the hexagonal platelets of MgC12 and of TiC13-A1C13 are indeed very close [1, 24]. Obviously, the mechanical energy brought during ball-milling favors the interactions between the platelets, and the resulting supported catalyst could be said "contaminated" by the support. Ball-milling is now a method to increase the activity of VPO catalysts for n-butane oxidation to maleic anhydride [50], as well as of perovskites for NOx abatement [51]. In the second example, it has been shown that the thermal stabilization of the ~5-alumina used as the washcoat on cordierite is achieved by means of the nucleation of a cubic LnA103 (Ln = La, Pr,'Nd, Ce) perovskite structure. Microdomains of LnA103 have been evidenced by HRTEM. Moreover,

181 the trend to sintering shown usually by platinum crystallites would be lowered because the matching between the [011] zone axis of Pt and the [110] of 8-alumina is very good [25-27]. 2.2. The two phases are transformed (cooperative reactions) The second case, is, to our opinion, as important as the first, because it is a way to demonstrate that the catalytic atmosphere is not necessary for the existence of interfacial coherence [9, 15, 17]. Heating the systems under nitrogen leads generally to transformations which would not proceed for components alone. The most important experiment concerned the transformation of anatase into ruffle in the presence of V205 by heating under nitrogen, V205 being simultaneously reduced. Then the idea was that, if the framework of TiO2 (anatase) was responsible for reduction of V205, similar frameworks like those of MNbO4 (M= A1, Ga) as well as of TiO2(B), would lead to the same effect [14, 19]. The reverse (rutilization of anatase in the presence of MOO3, VOPO4, etc.) was also observed [17]. Similarly, the transition c~ ~ ~-CoMoO4 is delayed in the presence of MoO3/TiO2 (Mo/Co = 1.9:1), the result being an increase of the activity/selectivity of the mixed system for butanemaleic anhydride. On the contrary, the transition a ~ 13-NiMoO4 is advanced in the presence of MoO3 (Mo/Ni = 1.26:1) in catalysts prepared by the fight way for propane-propylene, which is correlated to a stronger stabilization of MoO3 since it does not sublimate easily [23]. The formation of a third phase from the two first precursors can also be responsible for the stabilization of a support, as in the case of the three-way catalysts Pt/LnA103/8-A1203 just evoked. Other examples can be found in the ternary V-Nb-Mo-O system for ethane oxidation to acetic acid: one composition (Mo/V/Nb= 0.73/0.18/0.09) gives the best catalytic results, which has been interpreted as due to the formation of definite microdomains of (V, Nb) solid solutions of Mo18052 and Mo5014-type, embedded in a matrix of MoO3 [21, 22]. In such a way, the active sites would be more isolated than in pure MOO3, and consecutive oxidation of acetic acid or C2H4 to CO2 could be avoided. According to the considered system, the coherence of interfaces can lead to two opposite effects: the surface lattice of the catalyst being maintained in a metastable state, the result is an increase of activity (and/or selectivity) compared to the unsupported catalyst, which can be accompanied by a stabilization on time (and on stream). A last example is indeed provided by the case of H4PMollVO40 supported on SiO2 doped with potassium for the oxidative dehydrogenation of isobutyric acid to methacrylic acid [52, 53]. The easy adaptation of the amount of water and of protons between the Keggin units according to the degree of hydration and to the kind of reaction results in a remarkable "elasticity" of the heteropolyacid lattice allowing its anchoring on suitable supports. The formation of layer(s) of KxHl-xPMo12040type on the surface of K-SiO2 (cristoballite) would be responsible for the anchoring of the active H4PMol 1VO40 phase of cubic symmetry.

182 3. CONCLUSION In all the cases studied above, the common point is the existence of coherent interfaces due to the closely related structures of the two (or more) components. Very recently, a computer modelling study of the interfaces between V205 and TiO2 has given support to this model by considering the morphology of the two components, and showed also the differences in the case of a monolayer of V205/TiO2 [54, 55]. The atoms on the boundaries are strained, so that their activity is greater than when they are inside the pure phases, and the lowering of surface energy allows various specie to cross the interface. The coherence does not depend on the reactants surrounding the crystallites, since cooperative reactions can proceed in non reactive atmosphere. Microdomains of one phase inside the other can be formed during the transient state and be kinetically, but not thermodynamically, stable during the steady state. These microdomains are also a way to account for the necessity of site isolation proposed by Callahan and Grasselli to avoid total oxidation [56]. Finally, in the case of mixed oxides as, e.g., Bi-Fe-Mo-O [31, 32] or Mo-V-Nb-O [21], in which a second redox couple (Fe3+/Fe 2+ or V5+/V 4+ respectively) is supposed to modify the Mo6+/Mo5+potential as compared to that in the single phase (bismuth molybdate or MoO3 respectively), the lowering of the surface energies which allows the transfer of electrons through the interface is the only way to account for the mutual and beneficial influence of the two redox couples. Therefore, the generation of new active centres which accelerate the reaction, as proposed by Delmon [29, 30], can be understood as the result of the coherence of interfaces. The same idea is valid for oxygen ions which are supposed to be able to cross the interface if an oxygen vacancy is available on the other side. Though hydrogen, a far smaller atom, is well known to spill over metals and oxides, any anionic oxygen vacancy present on the surface of a metallic oxide will act as a trap for 0 2- (or other ionic forms). When applying the Crystallochemical Model of Active Sites (CMAS) to vanadyl pyrophosphate [57], thermodynamic considerations showed that hydrogen could indeed spill over the (100) (VO)2P207 surface (2 to 3 sites) but that for oxygen, a kind of hopping model (known for electrons), or afortiori oxygen spill over, was unlikely.

Acknowledgements:

Part of this work has been made in the frame of the EC-HCM

programme EBRXCHCT-93-0261.

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184 33. R.G. Teller, M.R. Antonio, J.F. Brazdil and R.K. Grasselli, J. Solid State Chem., 64 (1986) 249. 34. P. Gai-Boyes, Catal. Rev.-Sci. Eng., 34 (1992) 1. 35. J. Haber, New Developments in Selective Oxidation by Heterogeneous Catalysis, P. Ruiz and B. Delmon (eds.), Stud. Surf. Sci. Catal., 72 (1992) 279. 36. H. Kn6zinger and E. Taglauer, Catalysis, 10 (1993) 1. 37. I.E. Wachs, R.V. Saleh, S.S. Chan and C.C. Cherzick, Appl. Catal. 15 (1985) 339. 38. G. Deo and I.E. Wachs, J. Catal. 146 (1994) 323. 39. R.H.H. Smits, K. Seshan, J.R.H. Ross, L.C.A. van den Otelaar, J.H. Helwegen, M.R. Anantharaman and H.H. Brongersma, J. Catal., 157 (1995) 584. 40. U. Ozkan and G.L. Schrader, J. Catal., 95 (1985) 120, 137 and 147; ibid Appl. Catal., 23 (1986) 327. 41. U. Ozkan, R.C. Gill and M. Smith, J. Catal., 116 (1989) 171. 42. U.S. Ozkan, M.R. Smith and S.A. Drsicoll, J. Catal., 123 (1990) 173. 43. J.L. Callahan, R.K. Grasselli, E.C. Milberger and H.A. Strecker, Ind. Eng. Chem. Prod.Res. Develop., 9 (1970) 134. 44. W. Gerhartz ed., Ullmann's Encyclopedia of Industrial Chemistry, 5th ed., vol A (1985) 153. 45. H. Ponceblanc, G. Coudurier, J.M. Herrmann and J. Vedrine, J. Catal., 142 (1993) 373. 46. P. Ruiz, B. Zhou, M. Rely, T. Machej, B. Doumain, F. Aoun and B. Delmon, Catal. Today, 1 (1987) 181. 47. B. Zhou, E. Sham, T. Machej, P. Bertrand, P. Ruiz and B. Delmon, J. Catal. 132 (1991) 157, 183 and 200. 48. R.K. Grasselli, J.D. Burrington and J.F. Brazdil, J. Chem. Soc., Farad. Disc., 72 (1982) 203. 49. G. Centi, E. Foresti and F. Guarnieri, New Developments in Selective Oxidation II, V. Cortbs-Corberan and S. Vic Bellon (eds), Stud. Surf. Sci. Catal., 82 (1994) 281. 50. V.A. Zazhigalov, J. Haber, J. Stoch, L.V. Bogutskaya and I.V. Bacherikova, Appl. Catal. A, 135 (1996) 155. 51. T. Hibino, K. Suzuki, K. Ushiki, Y. Kuwahara, M. Mizuno, Appl. Catal. A, 145 (1996) 297. 52. C. Desquilles, M.J. Bartoli, E. Bordes, G. Hecquet and P. Courtine, Erd61 Erdgas Kohle, 3 (1993) 130. 53. Richard-Tessier, E. Bordes, P. Courtine, l lth Int. Congress on Catalysis, Baltimore, July 1-7, 1996, Preprint p. 247. 54. D.C. Sayle, D.H. Gray, A.L. Rohl, C.R.A. Catlow, J.H. Harding, M.A. Perrin and P. Nortier, J. Mater. Chem., 6 (1996) 653. 55. D.C. Sayle, C.R.A. Catlow, M.A. Perrin and P. Nortier, J. Phys. Chem., 100 (1996) 8940. 56. J.L. Callahan and R.K. Grasselli, AIChE. J., 9 (1963) 755. 57. J. Zio~kowski, E. Bordes and P. Courtine, J. Catal., 122 (1990) 126; ibid., Stud. Surf. Sci. Catal., 55 ( 1991) 625.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

Synergetic oxides

effects promoted

by

in operandi s u r f a c e

185

reconstructions

of

Eric M. Gaigneauxa, c, J. N a u d b, P. Ruiz a and B. Delmon a a Unit6 de Catalyse et Chimie des Mat6riaux Divis6s, Universit6 catholique de Louvain - Place Croix du Sud 2/17, B-1348 Louvain-la-Neuve, Belgium. b Laboratoire de G6ologie et de Min6ralogie, Universit6 catholique de L o u v a i n Place L. Pasteur 3, B-1348 Louvain-la-Neuve, Belgium. c "Aspirant" fellow for the Fonds National pour la Recherche Scientifique (F.N.R.S.) of Belgium. Synergetic effects are observed in the selective oxidation of isobutene to methacrolein w h e n catalysts are prepared by mechanically mixing MoO3 with BiPO4. Both higher conversion and selectivity to methacrolein are observed. Catalyst characterization was carried out before a n d after catalytic tests. Scanning electron microscopy revealed that, when mixed with BiPO4, the (010) faces of the MoO3 crystals, known to be non selective, reconstructed to more active and selective (100) faces d u r i n g the test. No effect was observed with p u r e MoO3 in the same conditions. In agreement with previous results, there was no indication of mutual contamination. The results lead to conclude that the two effects, namely the increase in activity and selectivity as well as the reconstruction, are linked and are both due to a remote control mechanism. This remote control is due to the action of spillover oxygen "Oso" generated by BiPO4 (the "donor") on MoO3 (behavin@ as an Oso "acceptor"). This confirms the interpretation given to identical results with mixtures of MoO3 and ~-Sb204 (which, as BiPO4, is k n o w n to be an Oso donor) and reinforces the conclusion that the origin of the synergy is not mutual contamination. As different donors have the same effects, the conclusion is that the synergy is due to spillover oxygen and not to the chemical nature of the phase mixed with MOO3. 1. I N T R O D U C T I O N Synergetic effects between oxide phases are very frequent in the selective oxidation of hydrocarbons. Multiphasic catalysts have performances superior to the sum of those of the constituting phases used alone. The main phenomena observed are an increase of the conversion of the hydrocarbon, an I m p r o v e m e n t of the selectivity for the desired partially oxygenated products and an increase of the lifetime and resistance to deactivation in severe operating conditions. Different hypotheses could in principle explain this cooperation (or these synergetic effects): formation of more active mixed phases or solid solutions, m u t u a l surface contamination, support effects, existence of structurally coherent phase boundaries (epitaxy), etc. However, only the remote control mechanism ("RCM") is relevant when the synergetic effects are obtained with physical mixtures of simple crystalline phases with no chemical reaction nor contamination of the phases even after long reaction times [1,2]. In the particular case of oxidation, this "RCM" theory considers that one of the phases (called the "donor") is able to activate molecular oxygen into a

186 highly active mobile species, called "spillover oxygen" ("Oso"). This phase migrates onto the surface of the other phase (the "acceptor")with which it reacts. This reaction brings about the creation of new selective sites a n d / o r the regeneration of sites deactivated during the normal redox cycles of the oxidation mechanism (these sites are usually slightly reduced). The remote control mechanism (rather than other possible hypotheses) has been shown to explain the synergetic effects observed between various phases (transition metal oxides, mixed metal oxides, etc) in various oxidation reactions (true selective oxidation, oxidative dehydrogenation, oxygenassisted dehydration) on various hydrocarbons (light alkanes, olefi-ns, amides, etc) in a wide range of operating temperatures (from 423 K to 723 K) [1,3-6]. The theory has also been strongly s u p p o r t e d b v many more fundamental arguments includin~ proofs of the migration of labellec~ 18Oso and accurate kinetic modelling [7-10]. ~' In the past years, many independent investigators have shown strong evidences of the existence of oxygen spillover and of the positive role played by this species in improving the performance of catalysts. In the latter case, these authors often alluded to a "spillover effect" rather than a "remote control". These two different terms nevertheless reflect the same concept [11-15]. The general purpose of the present work is to investigate new experimental lines to further clarify the mechanisms underlying the effects of the RCM and to better understand the role of Oso at the surface of the "acceptor phases" to improve their performances. In this context, the synergetic effects between MoO3 and (zSb204, prepared separately and physically mixed afterwards, have been extensively investigated in the selective oxidation ofisobutene to methacrolein. These effects corresponded to an increase of the isobutene conversion and of the selectivity for methacrolein at the expense of the total oxidation products. Strong arguments showed that the cooperation was caused by a remote control through the migration of Oso, with oc-Sb204 as "donorphase" and MoO3 as "acceptor phase" [16]. A special series of experiments was performed with MoO3 samples mixed with oc-Sb204, which were mainly developing the (010) basal face. In agreement with the wellknown "structural specificities" of MoO3 (the (010) crystalIographic faces of MoO3 are non selective), these special MoO3 were weakly active and totally non selective when used alone [17-20]. But when mixed with oc-Sb204, the synergetic effects observed both for the conversion and the selectivity to methacrolein, were the most dramatic ever detected. These suggested that selective sites had really been created in. operandi (during the reaction) at the surface of MOO3. Scanning electron mzcroscopy investigation of the MoO3 crystallites after the reaction in the presence of 0c-Sb204 revealed a reconstruction. The edges between the (010) faces and the (100) lateral faces, which were sharp in the fresh samples, had developed at the micrometric scale a facetted structure composed of a succession of steps with the vertical walls oriented as (100) [21,22]. The phenomenon was also observed at the nanometric scale using atomic force microscopy on macroscopic MoO3 monocrystals [23]. The reconstruction was not observed in either sets of experiments when MoO3 was used in the absence of 0c-Sb204. This is because Oso, produced by 0c-Sb204, reacts with the surface of MoO3 and favors the coordinations of surface Mo atoms typical of the (100) faces. In the in operandi conditions, this leads progressively to the reconstruction of the crystallites from (010) faces to (100) steps. The enhanced performance of the mixtures supposes the creation of new selective sites that are attributed to the creation of more of these (100) steps which are known to be more active and more selective than (010) faces (another aspect of the "structural specificities" mentioned in the literature). As this only occurred in the presence of ~Sb204, the picture was consistent with the occurrence of the synergetic effects for the mixtures and fitted perfectly with the atomic scale model proposed in the RCM for the creation of selechve sites under the influence of Oso [3,4]. The objective of the present communication is to further support this picture. If the model proposed above for the role of Oso in the creation of new selective sites

187 at the surface, of MoO3 is.correct, it can be speculated, that the same effects, namely synergetlc effects and simultaneous reconstructxons of the crystals, should be observed whatever the origin of the Oso, namely whatever the "Oso donor phase" with which MoO3 is mixed. Following this line, a series of experiments parallel to the ones summarized above was performed with another "donor phase", switching from the MoO3 + 0~-Sb204 system to the MoO3 + BiPO4 system. BiPO4 has been shown to act as an "Oso donor phase" in many other oxidation reactions over multiphasic catalysts [16,24-27]. Moreover, extensive experiments have discarded the possibility of mutual contamination between MoO3 andBiPO4 [25]. The oxidation of ~sobutene was carried out in the same conditions as with c~-Sb204 and, similarly, the catalysts were characterized by scanning electron microscopy, X-ray diffraction and specific area measurements before and after the catalytic test.

2. EXPERIMENT 2.1. Catalyst preparation (010) oriented molybdenum trioxide. MoO3 crystallites developing preferentially the (010) faces were synthesized by recrystallisation of an isotropic molybdenum trioxide commercial powder (BDH Chemicals, 99.5+%) in a flow of pure 0 2 at 873 K during 12 hours. The position of the X-ray diffraction (XRD) peaks of the obtained yellowish solid fitted perfectly those of the JCPDS molybdite phase standard [28]. The anisotropy of the crystallites was determined by scanning electron microscopy (SEM) and checked by XRD, according to the method proposed by Ozkan et al. (Y. IhO0 / ~ IOkO = 0.005, where ~ IhO0 and ~ IOkO are the sums of the intensities of the peaks corresponding respectively to the reflections of the (h00) and the (0k0) series of crystallographic planes) ~18]. The SBET area was 0.81 m2.g -1. Bismuth plwsphate. A 0.05 M aqueous solution of Bi(NO3)3.5 H 2 0 (previously complexed with mannitol, mannitol / Bi3+ = 3 / 1 molar) was precipitated at room temperature using a 0.05 M aqueous solution of (NH4)2HPO4. The quantity of h y dro g eno..phosphate, used. was. determined considerin g a com p lete stoichiometric reaction with bismuth ions (1 B13+ for 1 PO43-). After washing with distilled water and subsequent lyophilisation, BiPO4 was obtained as a fine colloidal powder. The sample was thereafter calcined at 773 K in air during 20 hours. XRD pattern of the obtained solid corresponded to the JCPDS bismuth phosphate standard [28]. SBET area was 6.39 m2.g -1. Mechanical mixture. MoO3 (2.184 g i;,e. 1.77 m 2) were mixed in 250 ml of npentane with BiPO4 (0.316 g i.e. 2.01 m z) and physically interdispersed using ultrasounds during 10 minutes. No further mechanical agitation was made in order not to damage the MoO3 crystallites. Pentane was then removed by vacuum evaporation at room temperature before the mixture was dried at 353 K during 20 hours. The success of the mterdispersion of the two phases was checked by SEM. To rigorously compare the activity otthe mixture with that of the pure oxides, each pure oxide was submitted to exactly the same "mixing" procedure before being catalytically tested. Figure 1 shows SEM micrographs of the pure oriented MoO3 and the pure BiPO4 h a v i n g undergone the "mechanical mixture" procedure, and the mechanical mixture of these two. 2.2. Catalytic activity measurement The catalytic activity measurements were performed in a fixed bed reactor at 693 K. The gas tlow composition was isobutene 7 0 2 / He = 1 / 2 / 7 (vol.) with a total flow of 30 ml.min -1. The masses of catalysts used were so that the area developed by each phase (when present) in the reactor was identical when tested alone and in2the m!xture : namely, 150 mg of BiPO4 (i.e. 0.716 m2), 1039 mg of MoO3 d.e.U./3u m ) for the ~sts with the pure phases and 1189 mg ofdnechanical mixture (i.e. 150 mg or 0.716 m z of BiPO4 mixed with 1039 mg or 0.730 m z of MOO3). In order

188

Figure 1. SEM micrographs of the pure oriented MoO3 (left) having undergone the "mechanical mixture" procedure, and the mechanical mixture of the t w o - BiPO4 corresponds to the small white rod-shaped particles (right). not to perturb, the orientation, of the MoO3 crystals, the catalysts were used as powders without being pressed into pellets. The volume of the catalytic bed was kept constant for all the tests by diluting the catalysts in small glass balls previously checked to be inactive. The heating of the reactor was realized in the same flow as during the reaction at a carefully controlled rate of 7.5 K.min -1. The activity was measured during 3 hours, after which the catalysts were cooled down in the reactant stream at 7.5 K.min -1. Untransformed isobutene and the products of the reaction were analyzed at the reactor oulet by on-line chromatography. The catalytic activity was expressed in terms of conversion of isobutene (%C), yields (%Y) and selectivities (%S) in the different products, calculated as shown in equations la, b and c. Conversion of isobutene : %C = moles of isobutene transf~

of isobutene injected * 100

(la)

Yield in the product X (containing c atoms of carbon) : (lb) %Yx = (moles of X formed * ~ / m o l e of isobutene injected * 4) * 100 Selectivity for the product X" (lc) %SX = %Yx/~%C* 100 The synergetic effects between the oxide phases in the mechanical mixture were evaluated b y comparing the observed activities with theoretical values calculated assuming that no cooperation occurred, namely that the two phases in the mixture were behaving as if they were alone in the reactor. Equations 2a, b, c show how these theoretical values (noted with superscript ' th") were estimated.

189 Theoretical conversion of isobutene 9 %C th = %C obtained with BiPO4 + %C obtained with MoO3

(2a)

Theoretical yield in the product X" (2b) %yth,~ = %Yx obtained with BiPO4 + %Yx obtained with MoO3 Theoretical selectivity for the product X" (2c) %S~ = %Y ~ / / %cth 2.3. Characterization All the catalysts were characterized by X-ray diffraction (XRD), scanning electron microscopy (SEM) and specific area measurement (SBET) before and after the catalytic tests. XRD was achieved in the continuous symmetric analysis mode on a Kristalloflex Siemens D5000 diffractometer using the Kc~l radiation of Cu (~=1.5406A) for 20 angles going from 10 d e g to 80 deg. The scan rate was 0.4 deg.min -I (step size = 0.04 deg, step time = 6s). An additional high resolution XRD analysis from 10 to 60 de;~ (stev size = 0.016 de~, step time = 20 s, scan rate : 0.048 deg.min -1) was performec~ on the used mixture ~'or detecting an eventual very small amount of a contamination phase formed during catalysis. SEM was performed on a Hitachi S-570 microscope using a 15kV accelerating voltage. During the analysis, attention was focused on the detection of any modification in size, morphology and orientation of the crystals for both oxides during the catalytic work. All the SEM micrographs shown are representative of the whole of the samples, checking for the occurrence of the presented features in m a n y different places throughout the samples. Specific area measurement was made on a Micromeritics ASAP 2000 device. The analysis was based on the adsorption and desorption isotherms of Kr at the liquid nitrogen temperature. Specific areas were calculated according to the B.E.T. equation. Theoretical SBET values were calculated both for the fresh and the used mechanical mixtures on the basis of the specific areas developed by the pure oxides submitted to the same treatment and of the massic composition o f the mixture. 3. RESULTS

3.1. Catalytic activity measurement Table 1 summarizes the catalytic activity measurements for both pure phases and the mechanical mixture. The pure MoO3 presented a low conversion of isobutene, with a weak selectivity for methacrolein. The activity of the pure BiPO4 was similar to that of the molybdenum trioxide, but with a slightly higher conversion and selectivity for methacrolein. Important synergetic effects were detected for the mechanical mixture: (i) both the conversion ofisobutene and the yield and selectivity for methacrolein were higher than the theoretically calculated values (assuming the absence of cooperation between the phases, see equations 2a, b and c), (ii)the selectivity for CO2 was lower than the values expected if the phases had been behaving completely individually, (iii) propenol (which was not formed on any of the pure phases) was produced in small amounts on the mechanical mixture.

190 Table 1. Observed and theoretical values of the conversion of isobutene (%C), yields in methacrolein and propenol (%Ymeth and %Yprop), and selectivities in methacrolein and propenoI(%Smeth and %Sprop). Thedretical values (figures in parenthesis) have been calculated according tdeqhations 2a, b and c (assuming the absence of cooperation). Catalyst Pure MoO3 Mech. Mixt.

%C 5.13 25.87 (20.12) 14.99

Pure BiPO4

%Ymeth 0.34 6.01 (2.12) 1.78

%Yprop 0 0.41 (0) 0

%YcO2 3.94 10.71 (12.75) 8.81

%Smeth 6.59 23.24 (10.51) 11.85

%Sprop 0 1.59 (0) 0

%Sco2 76.86 41.39 (63.37) 58.75

3.2. Characterization results

3.2.1. X-ray diffraction The X-ray diffraction patterns obtained for the fresh pure phases submitted to the "mechanical mixture" procedure did not present any modification with respect to those obtained with the freshly synthesized samples. Similarly, for the fresh mechanical mixture, all the reflections corresponded perfectly, for the position and for the intensity ratios, to those of the constituting pure phases. No additional peaks were detected. After the catalytic tests, the comparison of the patterns of both used pure phases with the corres.,ponding fresh ones did not. reveal any modification. A similar perfect superposmon of the patterns was obtained for the fresh and used mixtures: (i) the MoO3 reflections presented unmodified ~ IhO0 / ~ IOkO ratios, (ii) all the detected reflections were assigned to either MoO3 or BiPO4 (this even if considering shoulders of intense peaks and small peaks with intensity in the range of the noise level- noise level = 0.2% of the intensity of the main peak) [28]. For the all the catalysts, no sign of amorphisation was observed. Figure 2 shows the XRD patterns obtained with the fresh and the used mechanical mixture.

ca B _

_

.

J _ _

J

ol.-.I

'

0

'

'

I

20

'

'

'

I

'

'

'

I

40 60 2 T H E T A A n g l e s (deg)

'

'

|

I

80

Figure 2. XRD patterns (lower resolution analysis) of the mechanical mixture MoO3 + BiPO4 before (A) and after (B) the catalytic reaction. For further exploring the possible occurence of a crystallographic modification in the mechanical mixture during the reaction, in particular the possible formation of mixed Bi-Mo-O or Bi-Mo-P-O phases, a systematical critical investigation was carried out on the high resolution XRD pattern (noise level = 0.1% of the intensity of the

191 main peak) of the used mixture. First, table 2 shows that it was possible to assign all the peaks detected (even the less intense ones) to either MoO3 or BiPO4. Table 2. XRD p e a k s d e t e c t e d for the u s e d m e c h a n i c a l m i x t u r e (high r e s o l u t i o n analysis 9scan rate = 0.048 d e g . m i n -1) - letters in r e g a r d of each p e a k r e p r e s e n t the p h a s e to which it is assigned" B = BiPO4, M = MOO3.

ilW.+~mIT~ 6.9334 5.1961 4.7701 4.6616 4.1630 4.0870 3.8519 3.8093 3.5072 3.4673 3.2812 3.2627

36.1 0.4 0.6 1.5 2.0 0.9 1.4 19.4 4.7 100 4.1 25.2

M B B B B B M M B B+M B M

3.1426 3.0688 3.0087 2.9644 2.9378 2.8656 2.7023 2.6534 2.5987 2.5293 2.4422 2.4355

0.8 4.4 2.5 0.9 0.7 2.4 3.0 5.5 1.5 3.0 0.9 1.0

B B M B B B M M B+M M B B

~lIIV:,/FII~ 2.3284 5.5 2.3113 78.8 2.2712 6.3 2.1742 1.4 2.1508 1.2 2.1314 1.8 2.1148 1.5 1.9812 3.2 1.9593 6.6 1.9312 0.7 1.8851 0.9 1.8661 0.9

B+M M M B B B+M B M B+M B B B

~lIIV'.ll/IVII~ 1.8493 3.6 1.8221 15 1.7957 0.9 1.7702 1.8 1.7597 1.2 1.7564 1.6 1.7343 3.9 1.6938 1.6 1.6636 2.1 1.6313 2.0 1.5974 5.0 1.5650 5.7

B+M M B M M B M M M M M B

Second, the s t a n d a r d peak files of all the phases involving Bi, Mo and O or Bi, Mo, P and O (which w e call "mixed phases") available in a regularly u p d a t e d JCPDS data b a n k w e r e c o m p a r e d with the m e a s u r e d p a t t e r n of the u s e d m l x t u r e [28]. Table 3 lists the s t a n d a r d reflection m a i n lines of these s t a n d a r d s that w e r e missing on the p a t t e r n of the tested mixture. A peak was considered as missing w h e n (i) if present, it w o u l d not h a v e o v e r l a p p e d with a line of MoO3 or of BiPO4, (ii) w h e n for the considered d-spacing, the m e a s u r e d pattern only presented a b a c k g r o u n d signal, (iii) a n d w h e n the closest m e a s u r e d peak p r e s e n t e d a shift of at least 0.2 deg f r o m the considered s t a n d a r d line. A c c o u n t taken of these requirements, no s t a n d a r d line of the m i x e d phases (as d e f i n e d above) was detectable. The conclusion is that, on the basis of the actual JCPDS files, the two sets of information, (i) the one from the high resolution XRD analysis, and (ii) the perfect s u p e r p o s i t i o n of the l o w e r resolution p a t t e r n s for the fresh and the used mixture, allow to discard the eventuality that a m u t u a l crystalline contamination b e t w e e n MoO3 and BiPO4 has f o r m e d in operandi in the mixture [28]. 3.2.2. Specific area m e a s u r e m e n t s Table 4 s h o w s the o b s e r v e d and theoretical SBET values for the catalysts before and after the catalytic work. The p u r e phases having u n d e r g o n e the "mixture" p r o c e d u r e p r e s e n t e d a lower SBET than the freshly p r e p a r e d ones. This was very likely d u e to a flocculation of the grains d u r i n g the low t e m p e r a t u r e drying after the p e n t a n e evaporation. O n the other hand, for all the catalysts, a slight increase of the SBET w a s o b s e r v e d after the catalytic reaction, p r o b a b l y c o r r e s p o n d i n g to deflocculation and p e r h a p s an attrition p h e n o m e n o n . For the mixture, h o w e v e r , the difference b e t w e e n the o b s e r v e d and the theoretical values always r e m a i n e d similar (in the range of the precision of the analysis).

192 Table 3. d-spacing (~) of the missing non-overlapping standard reflection lines for the Bi-Mo-O and Bi-Mo-O-P phases satisfying the detection criteria (see text). Phase Bi2MoO6 Bi4MoO9 Bi2MoO6 Bi2MoO6 Bi2Mo3012 q-Bi2MoO6 ~/'-Bi2MoO6 Bi20MoO33 Bi2xMo1-xO3 Bi2Mo3012 Bi6Mo2015 Bi2MoO9 ~/'-Bi2MoO6 [3-Bi2Mo209 3Bi203.2Mo03 7Bi203.MoO3 Bi0.27Mo205 Bi4MoO9 Bi6MoO12 Bi38Mo7078 Bi12Mo0.12018.4+x Bi7.9Mo0.1012.15 Bi3.64Mo0.3606.55 Bi9PMo12052

JCPDS N ~ 07-0401 12-0149 18-0243 21-0102 21-0103 22-0112 22-0113 23-1031 23-1032 23-1033 26-0216 31-0196 33-0208 33-0209 34-1250 34-1270 35-1491 36-0115 36-0116 38-0249 43-0196 43-0443 43-0446 30-0193

d-spacing (/k) 2.73, 2.68, 2.47,1.92,1.65,1.37,1.26, 1.25,1.22,1.21.... 2.83, 1.70, 1.41 4.30, 2.79, 2.00 8.09, 2.75, 2.74,1.94,1.65 7.89, 6.29, 4.90, 4..57, 3.62, 3.59, 3.27, 3.19, 2.80, 2.00.... 5.62, 2.79, 2.06 3.20, 2.80, 2.48,1.65 2.88, 2.75, 1.61 2.90, 2.79, 2.07, 1.68, 1.64 4.92, 3.60, 3.18, 2.76, 1.65, 1.54 6.04, 5.91, 5.71, 2.90, 2.68, 2.39, 2.07 2.83, 2.72 5.66, 2.80, 2.00 6.65, 5.95, 3.21, 3.20, 2.81 2.82, 2.73, 2.01,1.70 2.75 3.33, 2.04 2.90, 2.83, 2.77, 2.03, 1.70 2.80, 1.70 2.82, 2.80,1.70 2.74, 2.19,1.71 3.19, 2.74 2.82,1.70 4.86, 4.55, 3.59r 3.18, 2.88r 2.77~2.49, 2.00, 1.99r....

Table 4. SBET values (m2.g -1) of the catalysts before and after catalytic reaction. Values in parenthesis are the theoretical values calculated on the basis on the composition of the mixture and the observed values of the c o r r e s p o n d i n g constituting pure phases. Catalysts Pure MoO3 Mechanical Mixture Pure BiPO4

Before 0.7026 1.2699 4.7747

(1.2141)

After 0.7241 1.4491 5.2239

(1.2919)

3.2.3. Scanning electron microscopy The comparison of the SEM micrographs (magnification up to 30,000) of the pure MoO3 a n d the p u r e BiPO4 before and after the catalytic reaction did not show any modification o f the samples. Size and m o r p h o l o g y of the crystallites were unchanged. For the pure MOO3, in particular, the edges between the (010) and the (100) faces of the crystallites were totally identical in the fresh and the used samples, presenting a sharp intersection. On the other hand, for the mechanical mixture, important modifications occurred during catalytic reaction. While the BiPO4 crystallites remained unchanged, the MoO3 crystallites exhibited a morphology reflecting reconstruction. The edges between the (010) and the (100) faces acquired a facetted structure. Instead of a single intersection border (as in the fresh sample), these edges were composed o f a succession of small parallel steps oriented in the [001] direction of the crystals, namely with the vertical walls indexed as (100) faces. These special features were never observed in the fresh sample nor in MoO3 catalytically tested in the absence of

193 BiPO4. Figure 3 shows close views of the edges of a MoO3 crystallites before and after reconstruction in the presence of BiPO4, namely before and after catalytic test in mixture with BiPO4.

Figure 3. SEM micrographs of the edges between (010) and (100) faces of MoO3 crystallites in the mixture with BiPO4, before catalysis (left) and after catalysis (right). Arrows indicate the in operandi reconstructed features. 4. DISCUSSION The results of the catalytic tests clearly indicate the existence of a cooperation between MoO3 and BiPO4. The fact that synergetic effects were observed both for the conversion of isobutene and for the selectivity to methacrolein suggests that there has been creation of new selective sites during the reaction. It could first be argued that this creation of new selective sites is due to the formation of a more active mixed phase through a reaction between MoO3 and BiPO4. This hypothesis can be discarded, as high resolution XRD gives no indication of the presence of any of the 23 bismuth molybdate phases or of Bi9PMo12052 (the only Bi-P-Mo-O mixed phase reported in JCPDS files), even in very small amounts, in the used mechanical mixture [28]. The SEM investigations d i d not suggest the formation of crystallites or domains that could not be assigned to MoO3 or BiPO4 in the mixture after the test. In the same line, it could be also argued that the higher activity obtained with the mixture could be due to a sort of contamination that could not be detected by XRD, namely amorphous phase, solid solution or surface contamination (formation of active monolayer). Authors have reported the existence of such phenomena when starting from MoO3 and Bi203 (either from mixtures of these two phases, or from one phase i m p r e g n a t e d with the other) [29,30]. To our knowledge, no such possibilities have ever been reported w h e n starting from MoO3 and BiPO4. Moreover, exhaustive characterizations of the surface (including ISS) of mechanical mixtures of MoO3 with BiPO4 used in a similar reaction, namely the oxygen-assisted dehydration of N-ethyI-formamide, have definitely shown the absence of any mutual contamination [25].

194 Another possible . .explanation . of .the synert~eetic effects is alp hysical change of. the catalyst particles during the preparation of mxxture, or cturmg the catalytic reaction [31]. This could be a decrease of the size of the crystallites or a deflocculation phenomenon of the aggregates of crystallites, both corresponding to an increase of the amount of catalytic sites exposed. The tendency of the catalysts to present higher SBET values after the tests could be an argument in favor of thx's hypothesis. However, even if such phenomena could not be discarded, the observed SBET values for the mixture and those calculated on the basis of the massic composition of the mixture and of the SBET values measured for the pure oxides submitted to the same treaments actually remained similar. This implies that, even if some attrition or deflocculation occurred, the intensities of these phenomena were identical for the mixture and for the pure phases reacted alone. Each phase in the mixture should then behave as if it were aIone in the reactor without being influenced by the other one. This is not the case. Consequently, neither of these two phenomena could explain the enhanced performances of the mixture with respect to the pure phases (especially in selectivity, which is increased by a factor > 2). On the other hand, there is a conspicuous reconstruction of the MoO3 crystallites, namely the formation of more (100) steps at the edges with the (010) faces. The (100) crystallographic faces are selective for the partial oxidation [17-20]. It can then be concluded that the creation of selective sxtes as deduced from the catalytic activity measurements corresponds to the creation of more (100) selective faces by reconstruction of non selective (010) faces. Since the phenomenon only occurred when MoO3 was tested in the presence of BiPO4, this accounts perfectly for the observed synergetic effects. Recently, Smith and Rohrer indicated that some modifications of the catalytic performances of MoO3 could be correlated with the appearance of (100) steps on the (010) faces. But the difference with thepresent study is that they artificially triggered the reconstruction by drastically reducing MoO3 crystals with hydrogen instead of observing the phenomenon in a real catalytic process [32]. At this stage of the investigation, the types of newly formed sites responsible for the propenol formation are still not unc~erstood. Nevertheless, the correlation between reconstruction and synergetic effects observed in the system MoO3 + BiPO4 is perfectly identical to that observed in the system MoO3 + 0~-Sb204 [21,22]. This constitutes a further argument for discarding the contamination to explain the synergy. It is very unlikely that two different systems composed of different elements and phases, which would have brought about contaminations with different compositions, would present similar catalytic performances and similar morphologicalfeatures. As both BiPO4 and 0~-Sb204 are "spillover oxygen donors", it is much more logical to consider that the reconstruction was triggered by the reaction of Oso (flowing from BiPO4 or from,,~-Sb204) with the (010) faces of MOO3. Oso would favor the 'selective coordination of Mo atoms or groups of atoms with different coordinations typical of the (100) faces. This probably concerns groups of 4 Mo atoms permitting the concerted elimination o f 2 hydrogen atoms and the insertion of one oxygen atom into the hydrocarbon molecule. The formation of this selective structure would take place at the expense of the non selective one typical of (010) (namely an arrangement of Mo atoms, all identical in coordination, exhibiting one Mo=O bond and linked together by bridging oxygens). The following mechanism by which Oso influences the surface of MoO3 crystallites (reacts with it) for triggering the observed reconstruction can be proposed. It has been shown that the selective sites for partial oxidation at the surface of MoO3 were constituted of pairs of fully oxidized MOO6- octaedra resenting the "corner-sharing" structure (2 octaedra linked by 1 corner). After aving transferred an oxygen atom to the hydrocarbon, the resulting reduced structure reorganizes in an "edge-sharing" structure (2 octaedra linked by 1 edge). When the catalytic reaction is carried out in the absence of Oso, the concentration of

~

195 edge-sharing octaedra, pairs becomes progressively, higher and these edge-sharin, g structures start forming aggregates, leading to the nucleation of non selective shearstructures tfiat are difficult to reverse. The role of Oso is then, by reoxidizing the edge-sharing structures, to stop the growth of the rows of edge-sharing octaedra and to maintain these rows below the critical nucleus size corresponding to the formation of the shear-structures. In other words, Oso will continuously maintain a rapid dynamic swing between corner- and edge-sharing octaedra, so favorin, g frequent and rapid catalytic cycles...However,. at each. cy.cle, the reduced ed g...e-sharlng structure can eventually be reoxldlzed in a configuration different to the lmtlal one (a different crystallographic orientation). Hence, by increasing the frequency of the cycles, Oso also increases the probability to trigger a modification of the orientation of the pairs of octaedra. At the micrometric scale, this phenomenon leads, cycle after cycle, to the facetting of the large basal (010) faces of MoO3 crystals to (100) lateral steps. This modification of the surface must thus be considered as a dynamic phenomenon, namely a progressive process occurring during the succession of the reductionreoxidation cycles that are typical of the oxidation mechanism in steady state conditions [21-23]. 5. CONCLUSION When MoO3 was catalytically reacted in the presence of a spillover oxygen donor phase, namely BiPO4, the (010) faces of the crystallites got reconstructed to (100) steps. This corresponds to the creation of more selective sites (on (100) faces) at the expense of non selective ones (of (010) faces). This does not occur when MoO3 is tested alone. The phenomenon then accounts perfectly for the important synergetic effects observed between the two phases in the oxidation of isobutene to methacrolein. These results are completely identical to the ones observed when MoO3 was used in mixture with 0~-Sb204, another oxygen spillover donor. The observation of the identical phenomenon in two catalytic systems of different compositions, but each one presenting MoO3 (Oso acceptor) mixed with an Oso donor, strongly supports the relevance of this hypothesis. Even if corresponding to a different approach, the mechanism proposed here is intimately connected with the theoretical model corresponding to the remote control concept explaining how Oso maintains (or restores) the activity of MoO3 by keeping its surface covered by more selective corner-sharing pairs of octaedra at the expense of non selective edgesharing ones [3,4]. ACKNOWLEDGMENTS

Authors aregrateful to the Fonds National de la Recherche Scientifique for the fellowship awarded to Eric M. Gaigneaux. REFERENCES

1. 2. 3. 4. 5. 6.

L.T. Weng and B. Delmon. Appl. Catal. A : General, 81 (1992) 141. S. Breiter, M. Estenfelder, H.G. Lintz, A. Tenten and H. Hibst. Appl. Catal. A : General, 134 (1996) 81. B. Delmon. Surface Review and Letters, 2 N ~ 1 (1995) 25. B. Delmon. Heterogenous Chemistry Letters, 1 (1994) 219. E.M. Gaigneaux, D. Herla, P. Tsiakaras, U. Roland, P. Ruiz and B. Delmon. in : Heterogeneous Hydrocarbon Oxidation, ACS Symposium Series, Vol. 638, Eds. B.K. Warren and S.T. Oyama (USA, Washington, 1996) p. 330. E.M. Gaigneaux, P. Tsiakaras, D. Herla, L. Ghenne, P. Ruiz and B. Delmon. American Chemical Society Annual Meeting, Symposium on "Catalysis and

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Photocatalysis on Metal Oxides", Chicago, IL, USA, August 21-25, 1995. (accepted for publication in Catalysis Today, Eds. U.S. Ozkan, under press). L.T. Weng, P. Ruiz, B. Delmon and D. Duprez. J. Mol. Catal., 52 (1989) 349. D. Martin, P. Kaur, D. Duprez, E. Gaigneaux, P. Ruiz and B. Delmon. Catal. Today, Vol. 32 N ~ 1-4 (1996) 329. a) G. Mestl, P. Ruiz, B. Delmon and H. Kn6zinger. J. Phys. Chem., 98 (1994) 11269. b) G. Mestl, P. Ruiz, B. Delmon and H. Kn6zinger. J. Phys. Chem., 98 (1994) 11276. c) G. Mestl, P. Ruiz, B. Delmon and H. Kn6zinger. J. Phys. Chem., 98 (1994) 11283. D. Vande Putte, S. Hoornaerts, F.C. Thyrion, P. Ruiz, B. Delmon. Catal. Today, Vol 32 N ~ 1-4 (1996) 255. T. Otsubo, H. Miura, Y. Morikawa and T. Shirakaki. J. CataI., 36 (1975) 240. W.C. Conner, G.M. Pajonk and S.J. Teichner. Adv. Catal., 34 (1986) 1. G.E. Batley and A. Ekstr6m. J. Catal., 34 (1974) 36. D. Maret, G.M. Pajonk and S.J. Teichner. in Catalysis on the Energy Scene, Elsevier Science Edition (Amsterdam, The Netherlands)(1984), 347. in New Aspects of Spillover Effects in Catalysis (T. Inui et al, Editors), Elsevier Science Publishers (Amsterdam, The Netherlands) (1993) : (a) K. Fujimoto. p9, (b) S.J. Teichner. p27, (c) G.M. Pajonk. p85, (d) Y. Moro-Oka. p95. L.T. Weng, S.Y. Ma, P. Ruiz and B. Delmon. J. Mol. Catal., 61 (1990) 99. J.M. Tatibouet, J.E. Germain. J. Catal., 72 (1981) 375. R.A. Hernandez and U.S. Ozkan. Ind. Eng. Chem. Res., 29 (1990) 1454. K. Briickman, R. Grabowski, J. Haber, A. Mazurkiewicz, J. Sloczyniski and T. Wiltowski. J. Catal., 104 (1987) 71. a) B. Mingot, N. Floquet, O. Bertrand, M. Treilleux, J.J. Heizmann, J. Massardier and M. Abon. J. Catal., 118 (1989) 424. b) M. Abon, J. Massardier, B. Mingot, J.C. Volta, N. Floquet and O. Bertrand. J. Catal., 134 (1992) 542. E.M. Gaigneaux, P. Ruiz and B. Delmon. Catal. Today, Vol. 32 N ~ 1-4 (1996) 37. E.M. Gaigneaux, P. Ruiz and B. Delmon. 11th International Congress on Catalysis. Baltimore, MA, USA. June 30 - July 5 (1996). E.M. Gaigneaux, P. Ruiz, E.E. Wolf and B. Delmon. 6th Iketani Conference. International Symposium on Surface Nano-Control of Environmental Catalysts and Related Materials. Tokyo, Japan. Nov. 25- 27 (1996).(sumitted for publication in Appl. Surf. Sc.). L.T. Weng, P. Ruiz and B. Delmon. in New Developments in Selective Oxidation by Heterogeneous Catalysis (Eds. P. Ruiz and B. Delmon), Studies in Surface Science andCatalysis, Vol. 72 (1992) 399. a) J.M.D. Tascon, P. Grange and B. Delmon. J. Catal., 97 (1986) 287. b) J.M.D. Tascon, P. Bertrand, M. Genet and B. Delmon. J. Catal., 97 (1986) 300. c) J.M.D. Tascon, M.M. Mestdagh and B. Delmon. J. Catal., 97 (1986) 312. F.Y. Qiu, L.T. Weng, P. Ruiz and B. Delmon. Appl. Catal., 47 (1989) 115. F.Y. Qiu, L.T. Weng, E. Sham, P. Ruiz and B. Delmon. Appl. Catal., 51 (1989) 235. @ 1996 JCPDS- International Centre for Diffraction Data. K. Brtickman, J. Haber and T. Wiltowski. J. Catal., 106 (1987) 188. N. Arora, G. Deo, I.E. Wachs and A. M. Hirt. J. Catal., 159 (1996) 1. G. Mestl, B. Herzog, R. SchI6gI and H. Kn6zinger. Langmuir, 11 (1995) 3027. R.L. Smith and G.S. Rohrer. J. Catal., 163 (1996) 12.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

197

F u r t h e r S t u d y o n the S y n e r g e t i c E f f e c t s b e t w e e n M o O 3 a n d S n O 2 E.M. Gaigneaux 1, S.R.G. Carraz~n, L. Ghenne, A. Moulard, U. Roland 2, P. Ruiz and B. Delmon Unit~ de Catalyse et Chimie des MatEriaux Divis6s, Universitd Catholique de Louvain, Place Croix du Sud 2/17, B-1348 Louvain-la-Neuve, Belgium 1 Aspirant of Fonds National de la Recherche Scientifique, Belgium 2Instituto de Qui'mica, Facultad de Quimica, Universidad de Salamanca, Spain

In an earlier investigation, synergistic catalytic effects were observed between SnO 2 and MoO 3 in the dehydration - dehydrogenation of 2-butanol to butene and methyl-ethyl ketone at low temperature (190~

Synergy was explained on the basis of a remote control

mechanism involving the migration of spillover oxygen, "Oso".

SnO 2 was deemed the

"donor of Oso" and MoO 3 the "acceptor of Oso". A main concern voiced about this explanation was the probability of mutual contamination of the two phases during catalytic reaction which could account for the observed enhanced catalytic performance. The current study deals with experiments designed to investigate specially prepared compositions with "mutual contamination" between SnO 2 and MoO 3. The conclusion is, that the contamination, in any form, if formed in the mixture, cannot account for the synergistic effects. Additional experiments conducted with the same catalysts in the selective oxidation of isobutene at high temperatures (380 - 420~ also exhibit a catalytic synergism between SnO 2 and MoO 3. In this case, SnO 2 becomes highly active, which triggers its continuous reduction during the reaction. To restore its high oxidation level, SnO 2 pumps lattice oxygen from MoO 3. Consequently, MoO 3 becomes reduced. A further supply of Oso can prevent this reduction phenomenon. Finally, experiments using TPD of NH 3 show that Oso increases the acidity of MoO 3, consistent with an increase of butene production from 2-butanol when mixtures of SnO 2 and MoO 3 are used as the catalyst.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

199

T h e N a t u r e of the A c t i v e / S e l e c t i v e Phase in V P O C a t a l y s t s and the Kinetics of n - B u t a n e O x i d a t i o n D. Dowell, and J.T. Gleaves, Washington University Department of Chemical Engineering, One Brookings Drive, Campus Box 1198 St. Louis, Missouri 63130 Y. Schuurman, Centre National De Recherche Scientifique, Institut de Recherche sur la Catalyse, Delegation Regionale Rhone-Alpes, Secteur Vallee du Rhone, 2 avenue Albert Einstein, BP 1335, 69626 Villeurbanne Cedex, France The reaction of n-butane with "oxygen-treated" (VO)2P207 based catalysts has been investigated using high speed transient response techniques and Raman spectroscopy. Results indicate that two types of active oxygen are present on the VPO surface after oxidation. One type is associated with the production of CO2 and the other with the production of MA and CO2. Results also indicate that significant amounts of active oxygen can be stored in the VPO lattice, but Raman spectroscopic data indicates that the stored oxygen is not associated with a crystalline VOPO4 phase. 1. Introduction

The unique properties of VPO catalysts have motivated a large number of research studies, and there is strong interest in gaining a fundamental understanding of how the VPO system functions. Although the VPO has been investigated extensively a number of important features are still not well understood. This paper presents new results from high speed transient response experiments and Raman spectroscopic studies that provide insights into one of the key unresolved issues, namely, the nature of the active-selective phase. It is well established [1-3] that vanadyl pyrophosphate (VO)2P207 is an essential component of the most selective VPO catalysts. For example, structural and chemical characterization studies of "reactor equilibrated" VPO catalysts indicate that the predominate crb'stalline phase is vanadyl pyrophosphate (VO)2P207 [1-3], that the bulk P/V ratio is close to 1.0, and that the average vanadium oxidation state is close to +4.0 [35]. A number of studies [2,5] have indicated that alkane oxidation primarily involves oxygen adspecies adsorbed at vanadium surface sites, and relatively little bulk lattice oxygen. It is also generally agreed that surface vS+ species play a role in the n-butane reaction [4-10]. Recent studies using Raman spectroscopy, 31p_NMR, and other physical characterization techniques indicate that, in addition to surface species, V 5+ phases may be involved in catalyst operation [4-8,11,12,13]. Consistent with this idea is the fact that in the presence of oxygen (VO)2P207 can be readily transformed into different V5+ phases such as ~-VOPO4, t~-VOPO4, Y-VOPO4, and 5-VOPO4 [6]. The amount of V5+ depends on the oxygen partial pressure and the temperature. Consequently, under reaction conditions, small amounts of V5+ phases may be present at the catalyst surface [12]. Recent TAP-2 transient response studies have shown that (VO)2P207 can readily absorb oxygen into its lattice, and then efficiently channel it to the active catalytic site. It is not clear however how this oxygen is stored in the lattice. It is also well known that pure V OPO4 phases do not perform as well as "reactor equilibrated" (VO)2P207 for n-butane oxidation [6]. This fact indicates that a VOPO4 surface by itself cannot be the active phase. Consequently, a number of workers have suggested that the VPO active site is situated at a (VO)2P207/VOPO4 interface [5,7,8]. The purpose of this paper is to examine how exposure of a "reactor equilibrated" catalyst to different gas phase oxygen treatments influences the selectivity and rate of n-

200 butane conversion, and to determine if the reaction selectivity is associated with the formation of VOP04 phases.

2. Experimental 2.1. Catalyst Preparation. "Reactor-equilibrated" VPO catalysts were prepared by a nonaqueous procedure detailed in previous papers [ 14], and operated at steady-state conditions (1.5% n-butane, 15 psig reactor inlet pressure and 2000 GHSV) for approximately 3000 hours. Under steady-state conditions the catalyst gave selectivities to MA of approximately 66% at 78% conversion. XRD analysis of the reactorequilibrated samples showed that they were monophasic (VO)2P2OT. Chemical analysis gave a P/V ratio of 1.01 and vanadium oxidation state of 4.02. The samples had a BET surface area of 16.5 m2/gm. 2.2. TAP-2 Reactor System. Steady-state reaction studies, transient response experiments, and oxygen activation experiments were performed with a TAP-2 multifunctional reactor system. Details of the TAP-2 system have been presented in previous publications [5,14]. The system is comprised of 1) a high-throughput, liquid nitrogen trapped, ultra-high vacuum system, 2) a microreactor-oven assembly with temperature controller, 3) a heatable gas manifold with five input ports containing four pulse valves, and one manual bleed valve, 5) a valve control module for actuating the pulse valves, 6) a gas blending station for preparing reactant mixtures from gases and liquids, 7) a pressure transducer oven, 8) a quadrupole mass spectrometer (QMS), 9) a Pentium PC computer based control and data acquisition system, and 10) a slide valve assembly with heated exhaust line. Reactant gases can be introduced into the microreactor as steady flows or transient inputs with pulse widths (FWHH) of = 250/~sec. Reaction products are analyzed in realtime using a quadrupole mass spectrometer. Pulses from separate valves can be introduced as sets of pulses of predetermined length or in a pump-probe format alternating between two valves. For operation at atmospheric pressures the vacuum system is isolated from the microreactor using a slide valve. The slide valve contains an adjustable leak valve that controls the amount of reactor effluent that enters the vacuum system. The portion of the effluent that does not escape through the leak valve exits through an external vent that contains an adjustable pressure regulator. When the reactor is operated at atmospheric pressures the mass spectral data can be collected in a standard mass intensity versus mass number format. When the reactor is operated at vacuum pressures the slide valve is retracted, and the microreactor vents directly into the vacuum system. In this case transient response data are collected one mass peak at time in a mass intensity versus time format. The high pumping speed of the vacuum system, and the near proximity of the quadrupole to the microreactor insure that pulses measured by the quadrupole reflect the true microreactor transient response. 2.3. Collection of Raman Spectra. Raman spectra were obtained using a modified SPEX double grating spectrometer, and a Coherent Innova 90 argon ion (At+) laser operating at 514.5 nanometers (nm). The spectrometer was operated with a PCbased data acquisition and control system that was developed in-house. Reactor equilibrated VPO samples produced detectable Raman scattering using 30-60 mw of light (measured at the sample) and a scan rate of .02 nm/s. Catalyst samples from reaction studies were removed from the TAP-2 microreactor, immediately placed "as is" in a sealed quartz tube under one atmosphere of air. The quartz tube was rotated at 500-1000 rpm during data acquisition to prevent excess heating of the sample, and to give a spectra indicative of the average composition.

201 In a typical experiment the Raman signal was averaged for 10-20 scans to increase the signal to noise ratio. Experiments were performed to determine if any sample degradation was caused by exposure to the laser light, and none was observed for laser intensities up to 150 mw. Physical mixtures of reactor equilibrated VPO and different VOPO4 phases were used to gauge the detection limit of the system. Physical mixtures containing 4% or more of a VOPO4 phase produced peaks characteristic of the phase. 2.4. Reaction Procedures. All reaction studies were performed with reactorequilibrated catalyst samples. A standard catalyst charge weighed between 100 and 125 milligrams with an average particle diameter of 275 microns. The catalyst bed was centered between two beds of quartz particles with the same average particle diameter as the catalyst. The quartz beds were used to thermally insulates the catalyst and reduce the axial temperature gradient (< 5~ across the catalyst bed. Steady-state flow experiments were performed at 673 K and = 1 atmosphere pressure using a gas blend of 88% Argon, 10% oxygen and 2% n-butane. Molar flow rates were set to insure turbulent flow [14]. The reactor effluent was monitored by leaking a small amount into the TAP-2 vacuum system and collecting the mass spectrum. Oxygen uptake experiments were performed at constant pressure( = 1 atmosphere) and constant temperature in the TAP-2 microreactor. In a typical flow oxidation a reactor equilibrated catalyst was exposed to a flow of pure oxygen for a period of one hour. Oxidation temperatures ranged from 683 K to 803 K. After a period of one hour, the reactor was evacuated and the temperature was lowered to reaction temperature (653 K). Transient response experiments were performed immediately after oxidation treatments or steady-state reaction studies without exposing the catalyst to ambient conditions. All pulse response experiments using n-butane were performed with a gas blend containing a mixture of 78% n-butane and 12% argon. Typical pulse intensifies were in the range of 1014 molecules per pulse. The argon pulse response was determined to be independent of the pulse intensity. Under these conditions gas transport through the microreactor can be described by Knudsen diffusion. The argon, CO2, n-butane, and maleic anhydride responses were collected at m/e values of 40, 44, 58, and 98 respectively. From continuous flow experiments using pure reagents it was determined that the QMS signals at m/e = 40, 58, and 98 are unique to argon, n-butane and maleic anhydride respectively. The area of the transient responses at these mass numbers is directly proportional to the amount of each species. The experimentally observed CO2 response contained contributions from n-butane and maleic anhydride. To obtain the true CO2 response these contributions were subtracted out. Nbutane conversion and the relative selectivities to maleic anhydride and CO2 were obtained by measuring the areas of their respective response curves. 3. Results

Figure 1 shows a set of transient responses curves for a) n-butane (m/e = 58), b) CO2 (m/e = 44), and c) maleic anhydride (m/e = 98) when a 3.5/1 C4H10/Ar mixture is pulsed over an oxygen-treated catalyst sample maintained at 653 K. The oxygen treated catalyst was prepared by oxidizing a reactor equilibrated catalyst sample at atmospheric pressure for 1 hour at 723 K. Subsequently, the catalyst was exposed to a series of 10000 equally intense pulses of the C4HI0/Ar mixture. The labeled curves (Ox) in Figures la-c were acquired after 1000 pulses had been injected into the catalyst bed. The unlabeled curves were acquired after 8000 pulses had been injected. In going from the 1000th pulse to the 8000th pulse the n-butane response increases = 2.7 times and becomes broader. These changes indicate that 2.7 times more n-butane passes through the reactor unreacted, or that the n-butane conversion decreases 2.7 times. Conversely, the MA and CO2 responses decrease and become broader. The decrease in the response is due mainly to a change in the adsorption properties of the VPO surface as oxygen is removed. The MA

202 pulse width increases because MA is adsorbed more strongly on a V 4+ surface than a on V 5+ surface [141. Thus, the changes in size and shape of the response curves reflect a change in the kinetic state of the VPO catalyst.

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203 Both samples were prepared from the same batch of reactor equilibrated VPO, but one was oxidized for 1 hour at 723 K, and the other for 1 hour at 798 K. The responses indicate the relative n-butane conversion (2a) and MA production (2b) for the same sample size, and the same n-butane pulse intensity. The n-butane response is ~ 2 times smaller for the 723 K sample indicating that conversion is 2 times higher. The MA response is ~ 10 times larger on the 723 K sample indicating that MA production is 10 times higher, and the MA selectivity is 5 times higher. The height normalized MA responses displayed in the inset in Figure 2b show that the shape of MA response curve changes very little in comparison to the change in conversion. This change should be compared with the shape change shown in Figure lc. 1

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Figure 2a. N-butane responses over two oxygen-treated VPO samples, b. MA responses over the same two samples. Both samples were prepared from the same VPO, but one was oxygen-treated at 723 K and other at 798 K. Figure 3 displays Raman spectra of the samples used to produce the response curves presented in Figure 2 taken after exposure to ~, 30 n-butane pulse. The spectra (3a) of the sample oxidized at 723 K is characteristic of (VO)2P207, and exhibits no features indicative of crystalline vS+ compounds. The spectra (3b) of the sample oxidized at 798 K displays bands characteristic of (VO)2P207, and crystalline VOPO4 phases. In general, V 5+ bands were detected when the oxidation temperature exceeded 773 K. Figures 4a-g' display plots of the areas of the MA, CO2, and n-butane transient responses versus pulse number when a 3.5/1 C4I-I10/Ar mixture is pulsed over seven different oxygen-treated catalyst samples at 653 K. The seven catalysts were prepared from the same batch of reactor equilibrated VPO, by oxidizing each sample at a different temperature, and 1 atmosphere of oxygen for 1 hour. The seven oxidation temperatures, corresponding to the seven sets of MA, CO2, and n-butane data, are in order, 683, 703, 723, 7 4 3 , 7 6 3 , 7 8 3 , and 803 K. The sets of three curves reflect the changes in MA and CO2 production, and in n-butane conversion with pulse number. The MA curves (Figures 2a,b,c,d,e) of catalyst samples oxidized at or below 763 K display a pronounced

204 maximum that occurs between 500 and 1500 pulses of n-butane. The rise in the MA production coincides with a rapid decrease in the CO2 production (Figures 2a',b',c',d',e'). Figures 5a-f plot the product of the MA and n-butane pulse areas (i.e., the relative MA selectivity) versus the n-butane pulse number. Taken together, these plots show that the MA selectivity depends on the pulse number, and the oxygen treatment temperature

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Previous TAP-2 studies [5,14] have shown that reactor-equilibrated VPO readily adsorbs oxygen at reaction temperatures, and that the oxidation rate increases with temperature. Conversely, when an oxygen-treated catalyst is heated in vacuum it evolves

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208 were exposed to extended pulse reductions with n-butane, the phases were not depleted. These results indicate that these phases are not involved in the catalytic process. The series of curves displayed in Figures 4 and 5 show that during the initial pulse reduction of an oxygen-treated catalyst, MA production increases while n-butane conversion and CO2 production decrease. The CO2 curve is characterized by a slope change that occurs when the MA curve reaches a maximum. This behavior indicates that the production of CO2, prior to the MA maximum, involves a different type of oxygen species than is involved in the formation of MA. This result supports theoretical studies that indicate that more than one type of oxygen adspecies is present on an oxygen covered VPO surface [91. The amount of oxygen consumed during the CO2 decrease and the MA increase can be estimated from the n-butane conversion. Assuming that on average, each converted n-butane molecule consumes 10 oxygen atoms, the total number of oxygen atoms consumed per pulse is = 1015. In the most selective oxygen-treated sample (Tox = 723 K) the MA response reaches a maximum after = 1000 pulses or 1018 oxygen atoms have been consumed. In a typical 100 mg sample there are = 1018 surface vanadium atoms [5] so that the number of oxygen atoms consumed is approximately equal to number of surface vanadium atoms. The total amount of oxygen consumed during the entire reduction process is = 1019 atoms taking into account the decrease in n-butane conversion. Storage of this amount of oxygen by the conversion of V4+ species into V5+ species would require the transformation of ~- 10 monolayers of V 4+ species, and would change the VPO oxidation state from 4.02 to 4.12. This increase is consistent with previous chemical analyses of oxygen-treated VPO samples [5]. These results indicate that significant quantities of oxygen can be stored in the VPO lattice and used in the selective conversion of n-butane. The stored oxygen does not however, appear to be associated with a crystalline VOPO4 phase. The financial support provided by the National Science Foundation Grant Number CTS-9322829, Huntsman Chemical Company, and Amoco Foundation is gratefully acknowledged. The authors would also like to acknowledge Professor Gregory Yablonsky for many fruitful discussions. References

1. G. Centi, F. Trifiro, J. Ebner, V. Franchetti, Chem Rev., 88 (1988) 55. 2. J. Ebner, J. Gleaves, In Oxygen Complexes and Oxygen Activation by Transition Metals, A. Martell and D. Sawyer Eds., Plenum Pub.:(1988) 273. 3. G. Centi, Catalysis Today, 16 (1993) 5, and references therein. 4. Y. Zhang, R, Sneeded, J. Volta, Catalysis Today, 16 (1993) 39. 5. Y. Schuurman, J. Gleaves, J. Ebner, M. Mummey, In New Developments in Selective Oxidation II, V. Corberan and S. Vic Bellon Eds., Elsevier Pub.: Amsterdam (1994) 203. 6 Y. Zhang, M. Forissier, R. Sneeden, J. Vedrine, J. Volta, J. Catalysis, 145 (1994) 256. 7. M. Abon, K. Bere, A. Tuel, P. Delichere, J. Catalysis, 156 (1995) 28. 8. Y. Zhang, M. Forissier, J. Vedrine, J Volta, J. Catalysis, 145 (1994) 267. 9. P. Agaskar, L. Decaul, R. Grasselli, Catalysis Letters, 23 (1994) 339. 10. S. Albonetti, F. Cavani, F. Trifiro, P. Venturoli, G. Calestani, M. Lopez Granados, J. Fierro, J. Catalysis, 160 (1996) 52. 11. G. Hutchings, A. Desmartin-Chomel, R. Olier, J. Volta, Nature, 368 (1994) 41. 12. R. Overbeek, M. Versluijs-Helder, P. Wamnga, E. Bosma, J. Geus, In New Developments in Selective Oxidation II, V. Corberan and S. Vic Bellon Eds., Elsevier PUb.: Amsterdam (1994) 183. 13. B. Abdelouahab, R. Olier, N. Guilhaume, F. Lefebvre, and J.C. Volta, J. Catalysis, 134 (1992) 151. 14. Y. Schuurman, and J. T. Gleaves, Ind. & Eng. Chem. Research, 33 (1994) 2935.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

209

Understanding the microstructural transformation mechanism which takes place during the activation of vanadium phosphorus oxide catalysts G r a h a m J. Hutchings a, Andrew Burrows b, Sujata Sajip b, Christopher J. Kiely a'b, Kossi E Bere c, J e a n - C l a u d e Volta c, Alain Tuel c and Michel Abon c aLeverhulme Centre for Innovative Catalysis, D e p a r t m e n t of Chemistry, University of Liverpool, Liverpool, L69 3BX, United Kingdom. bDepartment of Materials Science and Engineering, University of Liverpool, Liverpool, L69 3BX, United Kingdom. CInstitut de Recherches sur la Catalyse, CNRS, 2 Avenue Albert Einstein, 69626, Villeurbanne Cedex, France.

S t r u c t u r a l characterisation studies on an undoped set of VPO samples have allowed us to follow the structural evolution of the catalyst during the activation procedure. The initial VOHPO4.0.5H20 precursor has a platelike morphology with an [001] surface normal. As the transformation proceeds a direct topotactic t r a n s f o r m a t i o n of the [001] VOHPO4.0.5H20 to [100] (VO)2P207 occurs at the periphery of the platelet. In the interior of the platelet, a more complex indirect transformation sequence occurs. D a r k field imaging experiments show t h a t regions exist where the VOHPO4.0.5H20 precursor t r a n s f o r m s epitaxially into [100] 5-VOPO4. As the activation time increases, the domains of 5-VOPO4, which are embedded in a disordered matrix, s h r i n k and further t r a n s f o r m to the final (VO)2P207 phase. An a t t e m p t has been made to correlate m e a s u r e d catalytic performance data with the catalyst microstructure at various stages of the transformation process. It is found t h a t there is a distinct parallel between improving catalytic performance and a decrease in the a m o u n t of V 5§ phases present. F u r t h e r m o r e , we show t h a t when Co is added to the catalyst as a promoter, it is initially homogeneously distributed throughout the hemihydrate platelet. As activation proceeds, however, the Co is seen to have limited solubility in the (VO)2P207 phase and preferentially segregates into, as well as structurally stabilising, the disordered matrix material. 1. I N T R O D U C T I O N

V a n a d i u m phosphate (VPO) compounds are important industrial catalysts used in the conversion of n-butane to maleic anhydride[l]. Although it is now generally accepted t h a t the reaction requires the presence of both V S§ and V 4§ cations in close proximity, the precise n a t u r e of the active site in this catalyst is still a m a t t e r for

210 speculation [2,3]. The catalysts are produced by activating a hemihydrate compound, VOHPOa.0.5H20, in an n-butane/air mixture at about 400~ for an extended time period. The resultant catalyst often consists of a complex mixture of vanadium phosphorus oxide phases {i.e. (VO)2P2OT, aI-VOPO4, ~II-VOPO4, T-VOPO4, 5-VOPO4 and VO(POa)2}. Some researchers [4] favour a single compound, (VO)2P2OT, to be the sole active phase and have indicated that the presence of other phases may be due to incomplete activation. However, in-situ Raman experiments by Hutchings et al [5] suggest t h a t specific combinations of some V 5§ phases (aii and 5-VOPO4) and a V 4§ phase ((VO)2P207) phase are a necessary requirement if the catalyst is to simultaneously exhibit good activity and selectivity. The picture is complicated further by evidence of 'disordered phases' or 'poorly crystallized pyrophosphate' material in VPO catalysts prepared in organic media, which is the preferred method for the preparation of industrial catalysts [6]. Such disordered phases are also likely candidates for regions that exhibit V 5§ and V 4§ ions in close proximity. The relative proportions of the phases present depends on a number of conditions such as activation temperature, activation time and precise composition of the activation atmosphere [7]. The phase distribution and morphology of the activated material also seem to depend critically on the precise preparation route used to form the hemihydrate precursor [8] and whether or not the material has been promoted by a dopant such as Fe or Co. The details of the solid state transformation processes taking place during catalyst activation need to be understood before an optimum set of transformation parameters can be defined. In this paper we present the results of a combined transmission electron microscopy (TEM) and powder X-ray diffraction (XRD) study of the activation of an undoped and Co doped VOHPO4.0.5H20 precursor prepared in an organic medium.

2. EXPERIMENTAL 2.1 Catalyst Preparation The precursor was prepared by adding V205 (11.8g) to isobutanol (250ml). HaPO4 (16.49g, 85%) was then introduced and the whole mixture was refluxed for 16h. The light blue suspension was then separated from the organic solution by filtration and washed with isobutanol (200ml) and ethanol (150ml, 100%). The resulting solid was refluxed in water (9ml/g solid), filtered hot and dried in air (ll0~ 16h). The XRD pattern of the precursor showed that a well crystallised VOHPO4.0.5H20 phase was obtained. Starting with the same powdered precursor in a standard laboratory microreactor, four separate experiments were carried out using the same reaction mixture (n-C4Hlo/O~Ie : 1.6/18/80.4) and flow conditions (2.4 1.h1 with VSHV=1500hI). Note t h a t the composition of the gas mixture was chosen to be similar to t h a t used to activate industrial VPO catalysts. The temperature was ramped up from room t e m p e r a t u r e to 400~ at a constant rate of 0.5~ -1. Although the initial t r e a t m e n t was identical for all four experiments, the time on stream at 400~ was systematically varied ; namely 0.1h., 8h., 84h. and 132h. The four 'activated' catalysts are denoted VPO-0.1, VPO-8, VPO-84 and VPO-132 respectively. The catalysts were quenched after these different activation times by cooling the reactor rapidly with reactants being present. Reactor products during each of the four activation runs were analysed using on-line chromatography. Carbon mass balances were typically 98-102% for all data cited.

211 A set of three Co-doped precursors (with 1, 2 and 5 wt% Co) were prepared by dissolving the required amount of cobalt acetylacetonate in isobutanol prior to the operation of refluxing with isobutanol and 85% H3PO 4. The subsequent filtration, washing and doping procedures were identical to that employed for the undoped precursor. These doped catalysts were then activated for 25h at 400~ under the same reaction mixture and flow conditions as described previously. 2.2. C a t a l y s t C h a r a c t e r i s a t i o n A combination of physical techniques were employed to characterise the catalyst microstructure. XRD analysis was performed using a SIEMENS D500 diffractometer operating with a CuK~ source. BET surface area measurements using nitrogen adsorption at liquid nitrogen temperatures were also carried out. Samples suitable for transmission electron microscopy analysis were prepared by dispersing the catalyst powder onto a lacey carbon film supported on a copper mesh grid. TEM observations were made in a JEOL 2000EX high resolution electron microscope operating at 200kV. 3. R E S U L T S AND D I S C U S S I O N 3.1. C h a r a c t e r i s a t i o n of the h e m i h y d r a t e p h a s e The XRD pattern of the undoped hemihydrate precursor was characteristic of well crystallised VOHPO4.0.5H20 as shown in figure l(a). When observed in the TEM each individual hemihydrate crystallite exhibited a characteristic rhomboid platelike morphology as shown in Figure l(b). The lateral dimensions along the major and minor axes of the rhomboid were typically about 2~m by ll~m respectively. The platelet thickness was usually in the 0.03 to 0.1~m range. Selected area diffraction patterns taken normal to one of these platelet confirmed that the major and minor axes of the rhomboid correspond to the [100] and [010] directions of the VOHPO4.0.5H20 crystal structure respectively.

Figure 1. (a) XRD pattern and (b) bright field image of the hemihydrate material. All the Co-doped catalyst precursors had a platelike morphology very similar to the undoped material. In addition, their XRD spectra were essentially identical to that

212 shown in Figure l(a) for the undoped VOHPO4.0.5H20. The chemical composition of the doped materials was monitored using energy dispersive X,ray (EDX) analysis, which confirmed t h a t the Co was homogeneously distributed throughout the h e m i h y d r a t e platelet.

3.2. Catalytic performance measurements The catalytic performance data for the four undoped samples activated for different time periods are summarized in Table 1. The n-butane conversion increases from 22 to 65% with increasing activation time. We also observe a decrease in the selectivity for CO and CO2. It is particularly noticeable t h a t the selectivity to CO drops off very rapidly in the first 10 hours. When the variation in surface areas is t a k e n into account we find t h a t the intrinsic activity, VMA, for maleic anhydride production increases steadily with time on s t r e a m up to 84 hours and then tends to level off as the catalyst becomes stabilized. Table 1 Catalytic performance data for the undoped VPO materials Selectivity (%) Time

Intrinsic

BET Area

on Stream (h)

Conv(%)

MA

CO 2

CO

Activity (10 -s mol cm -2)

(m2gi)

VPO-0.1

0.1

10.5

22

34

47

17

0.43

VPO-8

8

7.6

26

48

36

15

0.99

VPO-84

84

14.8

55

66

20

12

1.48

VPO-132

132

19.4

65

69

19

11

1.39

Catalyst

The catalytic performance of the Co-doped catalysts was compared to t h a t of an undoped VPO catalyst activated for the same time period. The results of our analysis are presented in Table 2, where it is clear t h a t the addition of Co in all cases has a beneficial effect on both the selectivity to maleic anhydride production and the specific activity of the VPO catalyst. The most significant improvement, however, was noted for the catalyst with the lowest cobalt loading. Table 2 Catalytic performance data for the Co-doped VPO materials (25h on stream) Catalyst

BET Area (mZg_l)

Conv (%)

Selectivity to MA (%)

Specific Activity (mol MA/mZh 1)

VPO

13.0

16

50.3

0.43

VPO: 1at%Co

16.3

36

71.4

0.99

VPO: 2at%Co

10.0

10

63.0

1.48

VPO: 5at%Co

10.0

10

62.4

1.39

213 3.3. S t r u c t u r a l e v o l u t i o n of t h e u n d o p e d c a t a l y s t a c t i v a t e d in b u t a n e / a i r XRD spectra for the VPO-O.1, VPO-8, VPO-84 and VPO-132 activated catalysts are presented in Figure 2. The materials can all be classed as 'poorly crystalline', although the peaks in the spectra for VPO-84 and VPO-132 do gradually become a little sharper. A comparison of the spectra in figure 2 with those obtained from crystalline standards of (VO)2PzO7 and the various V O P O 4 phases [9], does clearly show the presence of (VO)2P207 as characterised by the 200, 024 and 032 reflections at 23 ~ 28.45 ~ and 29.94 ~ respectively. Any peaks from VOPO4 phases, if present, are hidden in the background noise of the spectra. The XRD spectra therefore indicate t h a t the amount and degree of crystallinity of (VO)2P20 v present are steadily increasing with activation time. 150

150

(a)

100

100

50

50

0

,

5

=~15o

,

.

,

,

,

,

,

,

(c)

,

5

10 15 20 25 30 35 40 45 50

1"0 1'5 20 2} 3'0 3'5 4"0 4"5 5'0

(b)

.,.

150 100 100 50

50 ,

5

,

,

,

,

,

,

,

,

0

9

,

,

,

,

9

,

,

,

10 15 20 25 30 35 40 45 50 5 10 15 20 25 30 35 40 45 50 20/degrees

Figure 2. XRD patterns (a) VPO-0.1, (b) VPO-8, (c) VPO-84 and (d) VPO-132. Low magnification transmission electron micrographs of typical crystallites from the VPO-0.1, VPO-8, VPO-84 and VPO-132 samples are shown in figures 3(a), (b), (c) and (d) respectively. It is clear from this sequence of micrographs t h a t the characteristic rhomboid platelet morphology of the hemihydrate is retained to varying extents in all the specimens. The VPO-0.1 sample (fig.3(a)) shows a number of subtle changes from the crystalline hemihydrate platelet. Firstly a number of circular features, which are probably internal voids, have appeared. Secondly, isolated patches of the platelet show very characteristic fissures which seem to be crystallographic in origin because they tend to align along the minor axis of the rhomboid platelet. As the activation time increases to 8 hours (VPO-8; Fig.3(b)) the density of the circular void features has increased markedly and a distinct dark fringe about 25nm thick is beginning to form along the periphery of the platelet. The fringe appears darker t h a n the centre of the platelet due to diffraction contrast, implying that well crystallized material is nucleating at the platelet rim. The material in the interior of the platelet is much more disordered in character and is very sensitive to electron beam damage. As activation progresses further (VPO-84; Fig.3(c)) the crystalline fringe at the periphery is now continuous and has coarsened to a thickness of 50-70 nm.

214 a

b

......

!!:i:~i!~i:~i!~i:i'ii:il~ i:!,~:i~~,:

Figure 3. Bright field transmission electron micrographs of platelet morphologies in (a) VPO-0.1, (b) VPO-8, (c) VPO-84 and (d) VPO-132 activated catalysts. Detailed higher magnification studies confirmed that this periphery comprised small crystallites of (VO)2P2OT. In addition, large holes in the interior of the platelet of up to 300nm in diameter are now apparent. The residual interior material is still rather disordered and beam sensitive. Finally, the VPO-132 sample in figure 3(d) shows the 'end-state' where in addition to the well crystallized rim, the material in the interior of what remains of the platelet appears more crystalline in character. Selected area diffraction patterns obtained from the VPO-0.1 sample [7] confirmed the coexistence of VOHPO4.0.5H20, (VO)2P207 and 5-VOPO4. Furthermore these studies showed that the orientations of all three phases are epitaxially related. For instance, the epitaxial orientation relationship between VOHPO4.0.5H20 and (VO)2P207 is; [001]hemi // [100] pyr~ and [010] hemi// [010] pyr~ This orientation relationship has been reported previously [9]

215 in terms of the topotactic transformation that can occur between VOHPO4.0.5H20 and (VO)2P207. Our work however [7] also demonstrates a previously unreported epitaxial relationship between VOHPO4.0.5H20 and ~-VOPO 4 in which [001] hemi// [100] delta and [010]hemi // [001] delta

Figure 4. (a) Bright field image showing a typical platelet from the VPO 0.1 sample. Corresponding dark field micrographs taken in (b) the gpyr~ reflection of (VO)2P207 and (c) the gdelta=022 reflection of 5-VOPO 4. 5-VOPO 4 has been shown [7] to be suffer severe beam damage, and to overcome this we have carried some low dose dark field imaging experiments on the VPO-0.1 sample in an attempt to locate this phase spatially. Figure 4(a) shows a bright field image from a typical platelet from the VPO-0.1 sample. The corresponding dark field image shown in Figure 4(b) was taken in the 024 (VO)2P207 reflection, in which a thin peripheral fringe of (VO)zP207 crystallites is clearly seen. If the experiment is

216 repeated using the 022 5-VOPO 4 reflection then occasional crystalline patches about 100-200nm in diameter can be seen in the interior of the platelet (fig 4(c)). Furthermore, these domains of crystalline 5-VOPQ phase nearly always seem to be associated with the regions of the platelet showing crystallographic fissure features. Dark field imaging using a reflection associated with the hemihydrate phase just gives rise to a very diffuse image of the entire interior of the platelet. In summary our dark field imaging experiments suggest that after 0.1 hours on stream, crystalline [001] oriented pyrophosphate is just beginning to form at the rim of the platelet, whereas the interior seems to consist of very disordered hemihydrate phase in which discrete domains of 5-VOPO 4 have nucleated. These domains of 5-VOPO 4 and the residual hemihydrate material subsequently appear to progressively convert to the pyrophosphate phase with increasing time on stream. Simultaneously the pyrophosphate phase which forms epitaxially at the edge of the platelet gradually coarsens and thickens.

Figure 5. (a) Typical morphology observed in the 1, 2 and 5at% Co-doped VPO catalysts. EDX spectra obtained from the disordered interior material (b) and a blocky (VO)2PzOv crystallite at the platelet edge (c).

217 The major morphology observed in the 1, 2 and 5at% Co containing catalysts is shown in figure 5(a). The interior of the platelet after 25h of activation is largely disordered and exhibits large circular perforations. However, the blocky (VO)2P207 crystallites not only appear at the rim, but are also seen to decorate the flat platelet surfaces. EDX spectra obtained from the disordered interior material and a blocky (VO)2P207 crystallite at the platelet edge are shown in Figure 5(b) and (c) respectively. It is clear that the (VO)2P207 crystallites do not contain Co (at least at the 0.1at% detectability limit of the technique) which suggests that the dopant is preferentially segregating into the disordered phase. Observations on Co-doped samples which have been activated for extended time periods indicate that a rather large proportion of disordered phase is still retained even after 150h of activation. Hence, it seems that the dissolved Co may have a tendency to structurally stabilise the disordered phase. EDX characterisation of the 2 and 5at% Co-doped VPO samples showed similar concentrations of Co in the disordered regions of the platelet to that observed in the lat%Co-doped VPO sample. In the former of the two materials, the excess Co was found to be present as a secondary phase of vanadium doped cobalt phosphate. This suggests that there is a limited solubility of Co in the disordered VPO material and explains why increasing the Co loading does not necessarily lead to improved catalytic performance.

4. C o n c l u d i n g c o m m e n t s Under the particular butane/air transformation conditions we have used, there are two routes by which the undoped hemihydrate phase can convert to the pyrophosphate phase. At the edge of the platelet a direct topotactic transformation between VOHPO4.0.5H20 and (VO)2P207 occurs. In the interior of the platelet the transformation between the two phases can occur indirectly via an intermediate 5-VOPO 4 phase. As these two types of transformation can apparently occur side by side, it is highly likely that small changes in the transformation conditions or the precursor morphology may critically affect the proportion of direct to indirect transformation occurring. This may go some way to explaining why catalyst samples that have been prepared via different routes and which have not been fully equilibrated can show considerably different relative proportions of crystalline VOPO4, crystalline (VO)2P207 and disordered mixed V4+-V5+ phases. A further key feature which emerges from our study is that as the time on stream increases, the domains of crystalline 8-VOPO 4 gradually reduce to the pyrophosphate phase. We have not as yet been able to image this second stage in the transformation to determine whether or not a topotactic change between 8 - V O P O 4 and (VO)2P207 occurs. It is interesting to note that 31p NMR MAS and Raman evidence exists [9,11] to suggest that the 8 - V O P O 4 phase may in some circumstances convert to (VO)2P207 via a n ( z i i - V O P O 4 intermediate. Our current results do however demonstrate a distinct parallel between improving catalytic performances (n-butane conversion, activity and selectivity for MA production) and a decrease in the amount of V ~§ present. It seems likely that in fully equilibrated catalysts the V~+-V4§ redox couples that are needed for maleic anhydride production are either associated with V 5§ sites existing on the surface of crystalline (VO)2P207 or with the V5+-V4+ centres present in the persistent 'disordered pyrophosphate' phase. Our observations on Co-containing materials show that although the dopant is homogeneously distributed in the hemihydrate precursor phase, it has only a very

218 limited solubility in crystalline (VO)2P2OT. This in turn leads to a Co enrichment in, and structural stabilisation of, the "disordered" VPO component of the activated catalyst. The optimum 1at% Co loading level appears to correspond roughly to the composition at which all the Co can be incorporated in the disordered phase without leading to secondary cobalt phosphate formation. 5. A C K N O W L E D G E M E N T S

This work has been funded under an EEC BRITE-EURAM research programme (contract number BRPR-CT95-0046). REFERENCES

1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11.

G.Centi, (Ed), "Forum on vanadyl pyrophosphate catalysts", Catalysis Today, 16, (1994), Elsevier, Amsterdam. P.LGai and K.Kourtakis, Science, 267, (1995), 661. G.Centi, Catal.Today, 16, (1994), 1. J.R.Ebner and M.R.Thompson, Catal.Today, 18, (1994), 51. G.J.Hutchings, A.Desmartin-Chomel, R.Olier and J.C.Volta, Nature, 368, (1994), 41. M.T.Sananes, A.Tuel, G.J.Hutchings and J.C.Volta, J.Catal., 148, (1994), 395. C.J.Kiely, A.Burrows, G.J. Hutchings, K.E. Bere, J.C. Volta, A. Tuel and M. Abon, Discuss. Faraday Soc. No. 105 (1997) in press. C.J.Kiely, A.Burrows, S.Sajip, G.J.Hutchings, M.T.Sananes, A.Tuel and J.C.Volta, J.Catal, 162 (1996) 31. F.Benabdelouahab, R.Olier, N.Guilhaume, F.Lefevre and J.C.Volta, J.Catal., 134, (1992), 151. E.Bordes, Catal.Today, 1, (1977), 499. M.Abon, K.E.Bere, A.Tuel and P.Delichere, J.Catal, 38, (1988), 83.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

Structural and Catalytic A s p e c t s o f Phosphates

Some NASICON-

219

Based Mixed Metal

P.A. Agaskar a, R.K. Grasselli a'b, D.J. Buttrey b and B. White b Central Research Laboratory, Mobil Research and Development Corporation, PO Box 1025, Princeton, NJ 08543-1025, USA *

a

b Department of Chemical Engineering, University of Delaware, Newark, DE 19716-3116, USA

Some binary and ternary metal phosphates, e.g. NbVP3012, NbTiP3012, NbTil_xVxP3012, SbVP3012, SbTiP30~2 and SbTi~_xVxP3Ol2, possessing empty NASICON structures were synthesized, structurally characterized by x-ray powder diffraction, and a phase having the nominal composition of NbTi0.vY0.25P3012 tested for the oxidation of n-butane. The main aim of the investigation was to prepare compositions which would have structurally isolated vanadium moieties, thereby facilitating the partial oxidation of paraffins to oxidation products involving less than 14 electrons. Although encouraging results were obtained by synthesizing novel compositions with structurally defined topologies, the isolation of vanadium sites was insufficient in the Nb-Ti-V-P-oxide phase, which was catalytically investigated, to obtain the desired, anticipated results. The composition produced maleic anhydride (max. yield 30.6% at 57.8 conversion) as the sole non-COx partial oxidation product. At this juncture it is believed that extraneous, V- rich amorphous overlayers observed by TEM are partially responsible for the catalytic results obtained. Future studies are underway to further refine the syntheses methods and isostructural, key catalytic element substitutions, to achieve the desired goals of rational catalyst design and thereby influencing the reaction channels of oxidation reactions.

1. INTRODUCTION The rational design of catalysts has been a desired aim of catalyst researchers for a long time. Our current attempt at this goal centers on the partial oxidation of paraffins, and entails the incorporation of key catalytic elements into a structural framework which by its very nature would favor structural isolation of such catalytic functionalities. It is well known by now, that vanadium is one of the key elements for the oxidative activation of paraffins [1-5]. It is also well known, that structural isolation of catalytic moieties is desirable to achieve selectivity to useful oxidized products, thereby preventing overoxidation to waste products, CO and CO2 [6-8].

* Former address

220 With this background in mind, we chose vanadium as our paraffin activating element and the NASICON structure as the framework into which the vanadium and other catalytic moieties would be incorporated. The reaction to be studied was chosen to be the oxidation of n-butane. The NASICON structure was chosen because it can be readily synthesized, is thermally very stable, and can accommodate a large fraction of vacancies and cation substitutions [9-12]. In addition, this structure possesses two features which should be important for the catalyst design as envisioned above. First, it is a phosphate and hence expected, owing to its acidic nature, to stabilize the lower oxidation states of transition metals, e.g., V4+; second, owing to its structure, layered octahedral metal centers with variable valence are separated from each other by redox inactive tetrahedral phosphate groups, i.e., the structure provides for isolation of descrete layers. With these design parameters in mind, we synthesized a number of vanadium containing NASICONS, characterized them and checked the catalytic properties of some of them for the oxidation of n-butane.

2. EXPERIMENTAL

2.1 Synthesis and Characterization of Compositions a) Preparation of NASICON Phases. The phases were prepared by coprecipitation as described by Agaskar and Grasselli [13]. The following compositions were prepared in this manner: NbVP3012, NbTiP3Ot2, NbTi0.sV0.sP3Ol2, NbTi0.75V0.25P3O12, SbVP3012, SbTiP3012, SbTiP3012 and SbTi0.sV0.sP3012. The detailed description of the preparation of one of these phases is given below. b) Preparation of Nb Ti0.75V0.25P3012. This composition was prepared by coprecipitation as follows [13]: Ti2(C204)3 x 10H20 (0.015 mole) was mixed with V205 (0.005 mole) and Nb(C204)5 (0.04 mole) to which 2 5 0 ml distilled water was added and the mixture stirred and heated, until a blue-green slurry was obtained. To this slurry, 85 wt% H3PO 4 (0.12 mole) was added with additional distilled water (150 ml), and the mixture evaporated to dryness. The dry residue was ground, placed into a quartz crucible, heated to 450 ~ and held at that temperature for 10 hours. The calcined mixture was reground and calcined at 900 ~ for 2 hours. c) Surface Area Measurements. The BET surface area of the Nb-Ti-V-P-oxide catalyst described above was measured by a Micromeritics ASAP 2400 N2 physisorption apparatus. The surface area of the sample as synthesized measured 16.9 mZ/g. d) X-Ray Diffraction Measurements. XRD measurements of the Nb-Ti-V-P-oxide catalyst were performed at room temperature on a Rigaku instrument, using Cu Ka radiation (~=1.54178 A~ scanning conditions of 0.04 ~ step size at 2 s/step, automatic baseline correction, 11 point binomial smoothing and data accumulation from 2 to 52 ~ 20 . The indexing was performed using LaB6 (NIST SRM 660) as an internal standard. e) Magnetic Susceptibility Measurements. The magnetic susceptibility measurements were performed with a Johnson Matthey Magnetic Susceptibility Balance. The balance utilizes the Evans method for measuring susceptibilities [14]. The measurements reported here do not include diamagnetic corrections.

221

2.2 Catalyst Evaluation The Nb-Ti-V-P-O-catalyst was evaluated in a fixed bed microreactor unit described elsewhere [ 15], modified in such a manner, that all reactor effluent lines were heated to -200 C to prevent condensation of products. The product stream was analyzed by an on-line Varian Model 3600 GC for oxygenates and C5+ hydrocarbons and a Carle refinary gas analyzer for the lower hydrocarbons and fixed gases. The composition of the feed was adjusted to contain about 1.5% by volume n-butane in air, and hence below the lower explosion limit of this mixture. The butane and oxygen conversions are reported as the mole percentages of each in the feed, converted to products. The yields of carbon containing products are defined as the mole percentages of n-butane converted to the respective products. The selectivity of each product is defined as the ratio of the yield of each product divided by butane conversion. The carbon, hydrogen and oxygen balances are 100 +/- 5 %, and all results are normalized on a no loss carbon basis. Blank runs over acid-washed quartz chips showed no reaction under conditions similar to those made with the catalyst.

3. RESULTS and DISCUSSION 3.1 Structural Properties of Compositions The synthesized compositions were characterized by X-ray diffraction, and the results are summarized in Tables 1 and 2. Table 1 X-ray diffraction powder patterns of Nb-based mixed metal phosphates having NASICON structures

NbVP3012 d 6.185 4.431 4.311 3.714 3.087 2.793 2.508 1.968 1.935 1.856

NbTiP3012

NbTio.sVo.sP3012

I/Io

d

I/Io

d

I/Io

50 100 55 84 79 45 20 13 13 12

6.14 4.41 4.28 3.693 3.071 2.781 2.494 1.961 1.925 1.847

56 86 59 100 67 59 20 16 16 15

6.109 4.599 4.291 3.685 3.063 2.784 2.481 2.058 1.929 1.831

31 18 10 63 53 29 32 5 9 5

NbTi0.75V0.25P3012 d

I/Io

6.129 24 4.587 3 4.300 100 3.685 50 3.063 49 2.784 20 2.485 29 1.931 14 1.908 8 1.850 3

222 Table 2 X-ray diffraction powder patterns of Sb-based mixed metal phosphates having NASICON structures

SbVP3012 d 6.006 4.364 4.143 3.611 3.008 2.751 2.436 2.392 2.186 2.108 2.074 2.006 1.933 1.875 1.807

SbTiP3012

SbTio.sVo.sP3012

I/Io

d

d

50 1O0 59 53 54 44 25 21 4 7 4 5 6 17 14

6.01 4.36 4.16 3.615 3.013 2.748 2.442 2.405 2.186 2.108 2.101 2.010 1.934 1.876 1.810

I/Io 58 1O0 58 79 72 59 25 22 4 7 7 8 16 17 17

6.026 4.364 4.162 3.619 3.013 2.751 2.442 2.401 2.186 2.108 2.083 2.010 1.935 1.879 1.811

I/Io 59 1O0 53 62 59 51 25 19 3 6 2 5 9 19 12

It is apparent from the X-ray diffraction patterns presented (Tables 1 and 2) that the structures of the mixed metal phosphates synthesized here possess the empty NASICON structure. A single crystal of nominal composition SbVP3OI2with sufficient size for four-circle X-ray diffraction study was isolated. The unit cell is" a = 8.287, b = 8.287, c = 22.086, c~ = 90.00, 13 = 90.00, 3, = 120.00. The fractional coordinates are" Sb/V (1): xyz = 0.3333, 0.6667, 0.3100; Sb/V (2)" xyz = 0.6667, 0.3333, 0.1867; P" x y z = 0.2856, 0.2862, 0.2501; O (1)" xyz = 0.4511, 0.2523, 2355; O (2): xyz = 0.3579, 0.4973, 0.2554; O (3): xyz = 0.1349; 0.4786, 0.3579; O (4): xyz = 0.5209, 0.1172, 0.1373. Structural refinement reveals two interesting observations: First, the c-glide operation present in the parent NASICON structure ( space group R 3c ) is absent here ( space group R 3 ), which implies that neighboring octahedral layers are crystallographically unrelated. Second, the occupancies of the octahedral sites within the two distinct layers are indistinguishable from a random 50%V / 50%Sb mixture. The ramifications of this observation in the context of our catalytic anticipations need to be further contemplated and influenced, if need be, to achieve the desired catalytic goals. Most of the NASICON phases which we synthesized are relatively pure. However, some do contain small amounts (i.e., < 5%) of extraneous impurities, which have to be further determined and eliminated in future investigations. It is to be anticipated, that such impurities m

223 might deletariously affect the catalytic properties under investigation, and in particular influence the realization of the catalytic concepts which we proposed for this study and which we tried to verify here. One of the less contaminated phases was NbTio.vsVo.25P3012, although not entirely free of impurity phases, which was chosen for the catalytic investigation of nbutane oxidation. The results are presented below. Magnetic susceptibility measurements of two of the synthesized compositions strongly suggest, that the majority oxidation state of vanadium in these NASICON phases is V 4+. The respective measured values are 1.57 BM for NbVP3012 and 2.13 BM for SbVP3012. Additional measurements, including XPS are necessary to more exactly define the oxidation state of the vanadium.

3.2 Structural and Catalytic Properties of NbTi0.7sV0.2sP3012 The composition of NbTi075V025P3012 exhibits an XRD pattern characteristic of a NASICON structure (Table 1 and Figure 1) and has a surface area of 16.9 m 2/g. Its structural features lie intermediate between those of NbTiP3OI2 and NbVP3012 NASICONS, as might be expected. The catalytic properties of this phase were investigated for the oxidation of n-butane and are summarized in Table 3 and Figure 2. The results reveal that this composition possesses catalytic properties, activates n-butane under oxidizing conditions and converts it to maleic anhydride and waste products, CO and CO2. The sole non-COx product is maleic anhydride.

'

I

~

10.0

'

'"

'

I

20.0

'

'

'

'

I

'

30.0

Degrees 20

'

'

'

I'

'

40.0

'

'

'

I'

'

50.0

Figure 1. XRD pattern ofNbTio.75Vo 25P30~2 exhibiting a NASICON structure

224 Table 1: Oxidation of n-butane over NbTi0.75V0.2sP3012 (V/P - 1/12) Experimental Conditions Temperature (~ Feed Flow Rate (cc/min) n-C4~ %) O2/n-C4~ ratio) WHSV (per hour) Contact Time (see) Conversion Oxygen n-C4 ~ Selectivity Maleic Anhydride CO CO2 Yield Maleic Anhydride CO CO2 Balances C H O

450 142.7 2 11.9 0.09 2.4

500 142.7 2 11.9 0.09 2.4

450 71.2 1 15.1 0.04 4.9

500 71.2 1 15.1 0.04 4.9

2.9 8.6

9.1 26.2

6.2 24.0

22.3 57.8

71.1 20.2 8.8

57.5 33.9 8.6

66.6 23.3 9.3

52.9 37.3 12.4

6.1 1.7 0.8

15.0 8.9 2.3

16.0 5.6 2.2

30.6 21.6 7.2

99.7 99.7 98.4

100.9 100.9 100.0

100.1 100.1 99.9

99.3 99.3 103.3

As is customary for redox reactions, the yield of the useful product (maleic anhydride) rises with (n-butane) conversion, while selectivity to maleic anhydride declines with conversion. Under the conditions studied, the highest maleic anhydride yield of 30.6%, at a conversion of 57.8%, was obtained at 500 ~ using a feed composition of 1 n-C4/15 O2 / 84 N2, a contact time of 4.9 sec, and a WHSV of 0.04/hr. Under comparable conversion conditions we obtained with the most celebrated, n-butane to maleic anhydride catalyst (VO)zP207 [16], a maleic anhydride yield of 38.5%. We conclude therefore, that the Nb-Ti-V-P-oxide catalyst investigated is relatively active for the oxidation of paraffins, since it has a V/P ratio of only 1/12 as compared to a 1/1 ratio for (VO)zP207. Of course, we had hoped that the V in the NASICON structure would be sufficiently site isolated to yield products less oxidized than maleic anhydride from n-butane. However, unfortunately that does not appear to be the case. One explanation for this might be that there are still too many adjacent V atoms, i.e., (V-O-V)n moieties, where n > 0. Nonetheless, the NASICON structure provides for some desired V site isolation, however, apparently not complete and hence not sufficient to achieve our desired catalytic goal. Another observed fact is, that the Nb-Ti-V-P-oxide under investigation shows an amorphous overlayer via TEM which is enriched in vanadium. The (V/P)surface > (g/P)particle" One can reason that at the temperature of 900 ~ required to obtain the NASICON structure, the more

225 40

/ / / / / / /

O

E

30

LIJ r'n

.

ICl >...

/

"l"-9 20 Z <

/// /

(.) m

U.I .._1

9

<

1~

l.l..

10

/ / /

O

c3 _J LLI

>-

0

"1"

'

0

I

10

' 2 ; ' 3 ; ' 4 1 0 n-C,

o

' 5 ; ' 6 0

CONVERSION

Figure 2. Maleic anhydride yield vs n-butane conversion with NbTi0.75V0.25P30~2 as catalyst volatile vanadium preferentially migrates to the surface. In order to prevent such surface enrichment, it will be necessary in the future to control the synthesis environment. It might be possible to slow down the observed surface enrichment of vanadium by controlling the partial pressure of oxygen over the solid sample. It is reasoned, that under mildly reducing conditions (e.g., nitrogen, containing only residual oxygen, typically in the range -4 < log pO2 (atm) < -3), the in situ formation of a lower valent vanadium will slow down, or possibly prevent the undesirable vanadium to-the-surface migration, since the lower valent vanadium species are much less likely to migrate than the highly oxidized vanadium species. Such studies are currently underway in our laboratory. Additionally, it will be necessary to test our NASICON phases under milder reaction conditions; particularly at lower temperatures and greater hydrocabon dilution. Such conditions would be more conducive to yield less oxidized useful intermediates than those employed in this study. 4. CONCLUSIONS Vanadium containing NASICON compositions were synthesized, structurally characterized, and a composition of the empirical formula NbTio.75Vo.25P30~2 tested for the catalytic oxidation of n-butane. The study was undertaken with the premise to rationally engineer compositions which by choice of key catalytic elements and their placement in a chosen structure, might influence the reaction channel of given oxidation reactions. It was reasoned, that placing vanadium, a known paraffin activating element into a NASICON structure might

226 result in sufficient site isolation of vanadium, so as to lead to solids which might catalyze the oxidation of paraffins in a controlled way, giving partial oxidation products involving only a few electrons. While partial oxidation of n-butane occurred over NbTio.75Vo.25 P3012 leading to maleic anhydride as the sole partial oxidation product, the 14 electron oxidation was not exactly planned. Two explanations for the latter occurrence are advanced in the paper and possible remedies to channel the oxidation reaction into a less aggressive oxidation path are given. The latter include a focused approach to the synthesis of ternary and quaternary NASICON systems under controlled conditions, leading to expected site isolated small vanadium clusters, with completely isolated vanadium centers as an upper limit of site isolation in the supporting framework, and thus lower intermediate oxidation products such as furan from C4 hydrocarbons. Studies are currently under way in our laboratories to explore the hypotheses advanced here.

REFERENCES

1.a.A.T. Guttmann, R.K. Grasselli, J.F. Brazdil and D.D. Suresh, US Patent No. 4 746 641 (1988). b. R. Catani, G. Centi, F. Trifiro and R.K. Grasselli, Ind. Eng. Chem. Res. 31 (1992) 107. c. A. Andersson, S.L.T. Andersson, G. Centi, R.K. Grasselli, M. Sanati and F. Trifiro, Appl. Catal. A, 113 (1994) 43. 2.a.M.C. Kung and H.H. Kung, J. Catal., 134 (1992) 668. b. A. Corma, J.M. Nieto Lopez and N. Paredes, J. Catal. 144 (1993) 425. 3.a. Y-C. Kim, W. Ueda and Y. Moro-oka, Catal. Today, 13 (1992) 673). b. J.P. Bartek, A.M. Ebner and J.F. Brazdil, US Patent No. 5 198 580 (1993). 4.a.M. Ai, J. Catal., 101 (1986) 389. b. M. Ai, Catal. Today, 12 (1992) 679. 5.a.J.N. Michaels, D.L. Stern and R.K. Grasselli, Catal. Lett. 42 (1996) 135; 139. b. D.L. Stern, J.N. Michaels L. DeCaul and R.K. Grasselli, Appl. Catal. (1997) in press. 6. J.L. Callahan and R.K. Grasselli, AIChE J, 9 (1963) 755. 7. R.K. Grasselli and D.D. Suresh, J. Catal. 25 (1972) 273. 8. J. Nilsson, A.R. Lana-Canovas, S. Hansen and A. Andersson, J. Catal. 160 (1996) 224. 9. P. Hagenmuller, "Solid Electrolytes" Acad. Press, New York, W. van Gool (ed.), (1978). 10. A. E1 Jazouli, etal., C.R. Acad. Sc., Paris, t. 300, Serie II, 11, (1985) 493. 11. G.V.S. Rao, U.V. Varadaraju, K.A. Thomas and B. Sivashankar, J. Solid State Chem. 70 (1987) 101. 12. A. Sereghini, etal., J. Chem. Soc., Farad. Trans., 87 (1991) 2487. 13. P.A. Agaskar and R.K. Grasselli, US Patent No. 5 354 722 (1994). 14. D.F. Evans, Physics E. Sci. Instr., 7 (1974) 247. 15. D.L. Stern and R.K. Grasselli, J. Catal. 167 (1997) in press. 16.a.G. Centi, F. Trifiro and V.M. Franchetti, Chem. Rev. 88 (1988) 55. b. G. Centi, Catal. Today 16 (1993) 1. c. P.A. Agaskar and R.K. Grasselli, Catal. Lett. 23 (1994) 339.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

227

S e l e c t i v e R e a c t i v i t y of O x y g e n A d a t o m s on M o ( 1 1 2 ) for M e t h a n o l O x i d a t i o n Ken-ichi Fukui, Katsuya Motoda, and Yasuhiro Iwasawa Department of Chemistry, Graduate School of Science, The University of Tokyo, Hongo, Bunkyo-ku, Tokyo 113, Japan

Abstract

The selective oxidation of methanol on a Mo(112) surface was investigated by temperatureprogrammed reaction (TPR) and catalytic reaction in a constant flow condition of CH3OH and O2 (10-6-10-5 Pa). Low energy electron diffraction (LEED) and Auger electron spectroscopy (AES) were used to characterize the surface structure and the amount of elements on the surface. It has been found that formaldehyde (HzCO) was a major product during TPR of methanol on a Mo(112)-p(1 • surface (0o=1.0), while CH4, H2, C(a), and O(a) were the products at lower coverages of preadsorbed oxygen. Besides, this reaction proceeded without formation of H 2 0 and was considered to be a simple dehydrogenation (CH30(a)--> HzCO (g)+ 1/2H2(g)). Excess oxygen adatoms on Mo(112)-p(1• which were not incorporated into the p(1• structure, enhanced the selectivity to formaldehyde from 50 % to 90 % and lowered the activation energy of the methanol oxidation. Such oxygen adatoms were more reactive than the oxygen atoms of the p( 1 • structure and reacted with the methoxy species to form H20 by the oxidative dehydrogenation mechanism (CH30(a) + 1/20(a)~ H2CO(g) + 1/2 H20(g)). In a constant flow of methanol, the reaction proceeded several cycles but was deactivated by C(a) accumulated on the surface. The selective oxidation of methanol in flow conditions of CH3OH and 02 successfully proceeded on the Mo(112)-p(1• surface without deactivation.

1. I n t r o d u c t i o n

Control of reaction paths on catalyst surfaces by optimizing the structure and electronic properties is a key issue to be solved in surface science. Iron/molybdenum oxides are used as industrial catalysts for methanol oxidation to form formaldehyde selectively. The iron /molybdenum oxide catalyst consists of Fez(MoO4)3 and MOO3, and shows kinetics and selectivity similar to those of MoO3 for methanol oxidation [1]. It suggests that Mo-O sites play an important role in the reaction. MoO3 has a layered structure along a (010) plane, but the (010) surface is not reactive because it has no unsaturated Mo site [1]. On Mo metal surfaces such as (100) [2,3] and (112) [4], major products in methanol reactions were H2 and CO. Therefore, we considered that partial oxidation of Mo sites is needed for the selective oxidation of methanol. We have reported that methanol reaction pathways on Mo(112) could

228

van der Waals snhere \

oxygen adatom \ \ [1 TO]

oxygen atom

ii, ,!ii

!,, s,,

I

......

, il I

....

iiii,.

....

ti!i!,-

It

p(1 x2)-O

I

0 adatom + p(1 x2)-O

Figure 1. Models for oxygen-modified Mo(112) surfaces. be controlled by modification of the surface by oxygen atoms [4-6]. Formaldehyde (H2CO) was a major product during temperature-programmed reaction (TPR) of methanol on a Mo(l12)-p(l• surface (0o=1.0), while CH4, HE, C(a), and O(a) were the products on surfaces with lower coverages of preadsorbed oxygen. Besides, the reaction on Mo(112)p(1• proceeded without formation of H20 (CH30(a)--~ HECO(g)+ 1/2HE(g)). We have suggested that formaldehyde was formed due to suppression of C-O bonds of methoxy intermediates by selective blocking of the second-layer Mo atoms with the oxygen atoms. Oxygen modification of metal surfaces have been examined on methanol reactions. On some metal surfaces such as Cu(110) [7], Cu(111) [8], Cu(100) [9,10], Ag(110) [11], Ru(001) [12], Rh(111) [13], and Fe(100) [14], oxygen atoms enhanced the formation of methoxy intermediate by extracting the hydroxyl hydrogen of methanol to form OH(a). Another effect of oxygen modification was to stabilize the methoxy species as observed on Ni(110) [15], Mo(100) [3], and W(112) [16]. On Fe(100) surface, the stabilization of methoxy by oxygen atoms resulted in a change of selectivity [ 14,17]. The Mo(112) surface has a ridge-and-trough structure, where the top layer Mo atoms form close-packed atomic rows along the [ 111] direction separated by 0.445 nm from each other and oxygen atoms are expected to occupy quasi-3-fold sites composed of one second-layer and two first-layer Mo atoms [18]. Left-half of Figure 1 shows a model of the p(1• surface (0o= 1.0) which was proposed on the basis of LEED patterns and CO titration experiments [ 19]. Every second Mo row is coordinated by oxygen atoms on both sides(Mo2c), while the other Mo rows are not directly coordinated (MONc). This structure preserves adsorption sites on MoNc for admolecules such as CO, methanol, and ammonia. Selective blocking of the second-layer Mo atoms by oxygen atoms suppressed bondbreaking of C-O or N-H, resulting in m

CH30 [4-6] or NHx (x~2) [20] species persisting on the surface up to 500 K. In this study, we show that excess oxygen adatoms on Mo(112)-p(1• which are not incorporated into the p(1 • structure, enhance the selectivity to formaldehyde and lower the activation energy of the methanol oxidation. The selective oxidation of methanol in a flow of CH3OH and 0 2 successfully proceeds on Mo(112)-p(1• without deactivation.

229

2. Experimental The experiments were performed in an ultrahigh vacuum chamber which was equipped with a low energy electron diffraction (LEED) optics, which was also used for Auger electron spectroscopy (AES), and a quadrupole mass spectrometer (QMS) for TPR. A Mo(112) sample was cleaned by cycles of At+ sputtering and annealing to 1300 K. The sample could be cooled to 150 K by liq.N2 and resistively heated at linear sweep rates of 0.5-15 K/s. The cleanliness and surface order were checked by AES and LEED. The clean surface exhibited a sharp and well-contrasted p(1 • 1) LEED pattern, indicating that the surface preserves the bulk truncated structure. The p(1 • structure (0o-=1.0) was prepared by exposing the clean surface to 2 L (1 L=1.33• Pa.s) of O2 at 300 K and subsequent annealing to 600 K [19]. The oxygen coverage was monitored by AES and LEED patterns.

3. Results and Discussion 3.1. Temperature-programmed reactions (TPR) TPR spectra from the Mo(112)-p(1 • surface exposed to methanol at 200 K (Figure 2A) showed simultaneous desorption of HzCO, CI--I4, CO, and H2 at 560 K. We showed that the composition of the species remaining on the surface above 400 K was C:O:H = 1:1:3 and that the hydroxyl hydrogen recombinatively desorbed as H2 below 400 K [4]. We considered, therefore, that the intermediate was methoxy species as supposed on other metal surfaces [2,3, 7,8,11-13,15,17,21-23] or on MoO3 [1,24]. Methoxy species were also observed on oxygen-modified Mo(110) by X-ray photoelectron spectroscopy (XPS) [25] and on oxygenmodified Mo(100) by high-resolution electron energy loss spectroscopy (HREELS) [3]. The A, r~

I

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,

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I

' .,

,

'

~--.,.,~_

_~.._,..~

I

16 amu (CH4)

-

~

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I

'

i

I

'

I

16 amu (CH4) 28 amu

(CO)

2 ainu (H2) r

._~ O9u~ ~

32 amu (CH3OH)

~;

30 amu (H2CO)

200

400

600

T/K

800

1000

18 amu (H20) x2 30 ainu (H2CO) 200

,

I

400

,

I

,

600

T/K

I

800

,

I

1000

Figure 2. TPR spectra after exposing (A) a Mo(ll2)-p(lx2)-O surface and (B) a Mo(112)p(l•

surface with 0.15 ML of oxygen adatoms to 4 L of CH3OH at 200 K. The heating

rate was 5 K/s.

230 Table 1 The product distribution in TPR for the methanol reaction around 560 K on the Mo(112) surfaces modified with oxygen. Product s peci es

Y ield / ML p( 1x 2)-O

Oxygen adatoms (00' = 0.15)

surface (00 = 1.0)

+ p(1 x2)-O surface (00 = 1.0)

H2(g )

0.10

0.01

H2CO(g )

0.09

0.06

H20(g ) CO(g)

0 0.02

0.04 << 0.01

CH4(g )

0.04

<< 0.01

C(a)

0.03

0

O(a)

0.07

0 a)

a) 00, after TPR was 0.11.

adsorption sites of methoxy species are considered to be the MoNc rows in Figure 1 because of considerable steric blocking at the MoEc sites. For the methoxy decomposition on Mo( 112)p(1• selectivity to HECO was 50 % among C-containing products (Table 1). As noted above, the reaction was a simple dehydrogenation of methanol because no desorption of H 20 was observed at any temperature. Some of the methoxy species decomposed to C(a) (17 %) and O(a), and recombinative desorption of CO was observed at 800 K. Except for this recombinative desorption, the reaction products did not include the oxygen atoms in the p(1• structure. Therefore, the oxygen atoms worked as modifiers on the surface. It must be noted again that formation of H2CO was not observed on oxygen-modified Mo(112) surfaces with lower oxygen coverages [4]. The blocking of the second-layer Mo atoms, which are considered to be effective to dissociate C-O bonds [26,27], by oxygen atoms was needed to change the selectivity. Stabilization of methoxy species and enhanced selectivity to HECO by preadsorbed oxygen atoms were also reported on oxygen-modified Fe(100) surfaces [ 14,17]. Further adsorption of oxygen on the p(1 • surface at 300 K preserved relatively sharp subspots of p(1• while a little increase of background intensity was observed. The maximum coverage of the oxygen adatoms by exposure to 02 at 300 K was 0.5 ML. Therefore we assume that the oxygen adatoms adsorb on the MoNc rows of p(1 • structure as shown in the right half of Figure 1. Then the adsorption sites of the oxygen adatoms are competing with those of methoxy species. Figure 2B shows TPR spectra of methanol with oxygen adatoms(0o'=0.15). The oxygen adatoms were adsorbed on Mo(112)-p(1• at 300 K and then the surface was exposed to methanol at 200 K. The spectra show that oxygen adatoms (0o'=0.15) on p( 1• changed the selectivity; formation of CH4 and CO was suppressed and H2CO was formed with nearly 90 % selectivity at 565 K. Recombinative desorption peak of CO from C(a) and O(a) at 800 K was not observed. These results indicate that C-O bonds of methoxy intermediates are further stabilized by coadsorbed oxygen adatoms. It is to be noted

231 that desorption of H20 was observed around 565 K by reaction of oxygen adatoms. The reaction, therefore, changed to the oxidative dehydrogenation (CH3OH + 1/2 O2---" HzCO + H20 ) which is usually observed on oxide surfaces such as MoO3 or Fez(MoO4)3 [1,24]. We confirmed by separate experiments using NH3 as a reactant that oxygen adatoms reacted to form H20 at 520 K [28], while oxygen atoms incorporated in the p( 1 x2)-O structure reacted at 650 K [20]. Heating the p( 1• surface with oxygen adatoms over 803 K caused destruction of the p(l• structure and gave a distorted p(1 • LEED pattern. Once the p(1• structure was broken, oxygen adatoms became less reactive and H20 was not formed below 600 K any more [28]. TPR spectra of methanol on the p( 1 • surface showed that formation of CH4, CO, and C(a) increased compared to no preannealing, and selectivity to H 2CO became similar to that on the p( 1• surface without oxygen adatoms. When deuterated methanol was used as a reactant, the simultaneous desorption of TPR products at 560-565 K in Figures 2A and 2B shifted to higher temperature. It suggests that the rate-determining step of the reaction is cleavage of the C-H bond of methoxy species in both cases. We measured the peak temperature Tm of the reaction as a function of heating rate (2-15 K/s) and calculated the activation energy Ea for the reaction, assuming that the rate of the reaction-limited desorption is of the first order with respect to the coverage of methoxy on the surfaces. The activation energy for the reaction of CH30(a) on Mo(112)-p(1 • without oxygen adatoms was 177+_5 kJ mol-1 The value for CD30(a) increased by ca.10 kJ mol-1, which corresponds to the difference in the zero point energy of C-H and C-D bond. Coadsorption of oxygen adatoms lowered the activation energy by 30-40 kJ mol-1 both for CH30(a) and CD30(a). It suggests that the hydrogen atom is extracted by the oxygen adatoms themselves or by the MONc atoms in the first layer Mo, which are influenced by adsorption of the oxygen adatoms.

3.2. Catalytic reactions in flow conditions We also examined catalytic reactions on Mo(112)-p(1 • in constant flow conditions of CH3OH and O2 (10-6-10-5 Pa). We measured the amount of reaction products by a temperature-jump method as shown in Figure 3, because some mass fragments of CH3OH overlapped with reaction products and background level of each species tended to rise during the gas flow. The sample temperature was kept at 450 K and then CH3OH and O2 were admitted to the chamber by two variable leak valves while monitoring the partial pressure of each gas. After the partial pressures reached desired values, the sample temperature was jumped to the reaction temperature, kept for several minutes, and decreased to 450 K. No reaction occurred at 450 K, therefore, the area of a mass signal over the base line, which is bound between the signals at 450 K, corresponds to the amount of a product from the surface. The temperature jump was repeated and data were accumulated. Yields of products for typical conditions were summarized in Table 2. The amounts of C(a) and O(a) were measured by AES after stopping the reaction gas dose and evacuating the system. When only CH3OH was supplied to the Mo(112)-p(1 • surface above 560 K, the reaction products were HzCO, CH4, CO, H2, H 2 0 and a negligible amount of C2I-I6. The reaction rate gradually decreased at any temperature (560-800 K) and pressure (10-6-10- 5 Pa) examined. No reply to a temperature jump was observed after 870 s in the case of TR=560 K and PCmOH=2.1• Pa. But in any case, more than 1 ML of methanol reacted on the

232

'

I

'

I

'

I

'

I

'

N,' 700 460 16 amu (OH4) 28 amu (CO)

e-

x:i L_ t~ t~

e18 ainu (H20)

(/3 F

"

30 amu (H2CO)

z;

0

,

I

40

,

!

80

i

I

120

I

I

160

,

200

Time / s Figure 3. QMS response in a temperature-jump measurement during the catalytic reaction of methanol on Mo( 112)-p(1 x2)-O at Pcmoi4=2.1 x 105 Pa and Po2=6.5x 10-6 Pa.

Table 2 Yields of products during the catalytic reactions of methanol on Mo( 112)-p(1x2)-O. PCH3OH

PO2

/ Pa

/ Pa

TR / K Time / s

Yield / ML H2CO(g) CH4(g ) CO(g) H2(g ) H20(g ) C(a) O(a)

2.1x10-5

560

870

2.0

1.0

0.2

2.5

2.0

1.05

< 0.05

-6

560

1580

8.5

1.2

0.5

2.8

8.2

0.55

0.45

8.1• -6 1.6x10 -5 2.1x10-5

560 700

1900 a) 550

5.4 2.7

0 0.9

0.3 0.3

0.2 3.0

5.6 2.1

0.10 1.10

0.30 < 0.05

2.1x10 -5 6.5x10 -6

700

1790 a)

16.7

0

1.2

0

19.3

0.05

0.15

2.1x10 "5 6.5•

a) The reaction has not been deactivated and proceeded further in the condition.

233 surface, which means that the reaction proceeded catalytically. As show in Table 2, the selectivity to H2CO was 47 % at 560 K, which is close to the value obtained by TPR of methanol on the p(1• surface without the oxygen adatoms (Table 1). After the reaction, p(1• subspots of LEED were not observed due to a considerable increase in the background intensity. The results of AES showed that the surface was covered with C(a) (0r Therefore, the reaction ceased due to accumulation of C(a) by nonselective decomposition of CH3OH. A remarkable difference between TPR and the constant-flow reaction was formation of H20. As noted above, the oxygen adatoms on the MoNr sites reacted with H(a) to form H20 at 520 K if H(a) was supplied sufficiently. Once O(a) is formed by decomposition of CH30(a), it can react to form H20 with H(a) above 520 K. But during TPR, the reaction is initiated by the C-H bond scission of the methoxy species and cleavage of the C-O bond may occur later, which may result in insufficient supply of H(a) to O(a). As expected by the results of TPR in Figure 2, deactivation of the methanol dehydrogenation in the flow conditions was suppressed by providing oxygen adatoms on the Mo(l12)p(1• surface. Providing 6.5• 10-6 Pa of O2 at 560 K, the amount of methanol reacted on the surface was over twice as many as that without oxygen supply and the selectivity to H2CO increased from 47 % to 79 %. But nonselective decomposition of methanol to C(a) also occurred in this condition and the reaction eventually stopped. At a higher ratio of 02 to CH3OH ( PCH3OI-I=8.lx 10-6 Pa, Po2 = 1.6• 10-5 Pa), the selectivity to H2CO increased to 93 % and the reaction seemed to proceed catalytically without significant deactivation. Thus, the oxygen adatoms retain the C-O bond of methoxy on the p(1 • surface and the selective formation of H2CO can be achieved. The condition listed at the bottom of Table 2 showed the best result in our experiments. Much more than 10 ML of H2CO was formed with a selectivity of 93 %. Note that these yields were the minimum values because no deactivation was observed even after 3200 s. Relatively sharp subspots of p(l• were observed and accumulation of C(a) was not observed by AES. In this condition, the reaction rate had the first order relation with PCmOH and had nearly 0th order relation with Po2. The activation energy for the reaction in a flow condition (PcmoH=8.4• Pa) was estimated by QMS response at the initial temperature-jump as a function of the reaction temperature. Addition of oxygen (Po2=2.6• 10-5 Pa) lowered the activation energy by 15 kJ tool-1. Thus, the role of oxygen adatoms to enhance the selectivity to H2CO and to decrease the activation energy of the reaction was confirmed similarly to the TPR experiments.

4.

Conclusions

The results presented here indicate that the methanol reaction path is dramatically sensitive to the coverage and arrangement of the oxygen adatoms and the surface metal structure. On Mo(112)-p(1• methoxy species were stabilized by the oxygen modifiers in the p(1 • phase, which selectively blocked the Mo atoms with high coordination. This type of modification contributed to the formation of H2CO on the surface, but with 50 % selectivity. The coadsorbed oxygen adatoms enhanced the selectivity to H2CO over 90 % and the reaction proceeded catalytically in the flow conditions on the Mo(112)-p(1• surface without significant deactivation. Besides, the oxygen adatoms decreased the activation energy for the selective oxidation. Thus, the present results suggest that we can control the reaction path by

234 designing the reaction field with the two types of coadsorbed oxygens, modifier atoms and adatoms.

Acknowledgement This work has been supported by CREST (Core Research for Evolutional Science and Technology) of Japan Science and Technology Corporation (JST).

References 1. U. Chowdhry, A. Ferretti, L. E. Firment, C. J. Machiels, F. Ohuchi, A. W. Sleight and R.H. Staley, Appl. Surf. Sci. 19 (1984) 360. 2. E. I. Ko and R. J. Madix, Surf. Sci. 112 (1981) 373. 3. S. L. Miles, S. L. Bernasek and J. L. Gland, J. Phys. Chem. 87 (1983) 1626. 4. K. Fukui, T. Aruga and Y. I wasawa, Surf. Sci. 295 (1993) 160. 5. T. Aruga, K. Fukui and Y. I wasawa, J. Am. Chem. Soc. 114 (1992) 4911. 6. T. Aruga, K. Fukui and Y. Iwasawa, in: Catalytic Selective Oxidation, Vol.523, eds. S. T. Oyama and J.W.Hightower (American Chemical Society, 1993) p. 110. 7. I. E. Wachs and R. J. Madix, J. Catal. 53 (1978) 208. 8. J. N. Russell, Jr., S. M. Gates and J. T. Yates, Jr., Surf. Sci. 163 (1985) 516. 9. B. A. Sexton, Surf. Sci. 88 (1979) 299. 10. R. Ryberg, Phys. Rev. B 31 (1985) 2545. 11. I. E. Wachs and R. J. Madix, Surf. Sci. 76 (1978) 531. 12. J. Hrbek, R. De paola and F. M. Hoffmann, Surf. Sci. 166 (1986) 361. 13. F. Solymosi, T. I. Tarn6czi and Berk6,A, J. Phys. Chem. 88 (1984) 6170. 14. J.-P. Lu, M. Albert, S. L. Bernasek and D. J. Dwyer, Surf. Sci. 239 (1990) 49. 15. S. R. Bare, J. A. Stroscio and W. Ho, Surf. Sci. 155 (1985) L281. 16. J. B. Benziger and R. E. Preston, J. Phys. Chem. 89 (1985) 5002. 17. J.-P. Lu, M. Albert and S. L. Bernasek, Surf. Sci. 218 (1989) 1. 18. H. Bu, O. Grizzi, M. Shi and J. W. Rabalais, Phys. Rev. B 40 (1989) 10147. 19. K. Fukui, T. Aruga and Y. Iwasawa, Surf. Sci. 281 (1993) 241. 20. T. Aruga, K. Tateno, K. Fukui and Y. Iwasawa, Surf. Sci. 324 (1995) 17. 21. F. M. Leibsle, S. M. Francis, R. Davis, N. Xiang, S. Haq and M. Bowker, Phys. Rev. Lett. 72 (1994) 2569. 22. F. M. Leibsle, S. M. Francis, S. Haq and M. Bowker, Surf. Sci. 318 (1994) 46. 23. S. M. Francis, F. M. Leibsle, S. Haq, N. Xiang and M. Bowker, Surf. Sci. 315 (1994) 284. 24. W. E. Farneth, F. Ohuchi, R. H. Staley, U. Chowdhry and A. W. Sleight, J. Phys. Chem. 89 (1985) 2493. 25. J. G. Serafin and C. M. Friend, J. Am. Chem. Soc. 111 (1989) 8967. 26. L. M. Falicov and G. A. Somorjai, Proc. Nat. Acad. Sci. USA 82 (1985) 2207. 27. J. Wang and R. I. Masel, J. Catal. 126 (1990) 519. 28. K. Fukui, K. Motoda and Y. Iwasawa, to be published.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

235

M e c h a n i s t i c studies of alkane partial oxidation reactions on nickel oxide by m o d e m surface science t e c h n i q u e s # Nancy R. Gleason~ and Francisco Zaera* Department of Chemistry, University of California, Riverside, CA 92521, USA.

The reactions of alkyl iodides, precursors to alkyl surface moieties, with adsorbed oxygen have been studied on Ni(100) by using temperature-programmed desorption, x-ray photoelectron spectroscopy, and ion scattering spectroscopy. The oxygen coverage was varied from submonolayer values to exposures which lead to surface oxidation. The reaction of 2-propyl iodide on surfaces completely covered with oxygen was found to lead to the complete oxidation of the alkyl iodide, but for low oxygen precoverages some acetone production was observed as well. The mechanism for the formation of acetone from 2-propyl iodide was investigated, and the reactions of other alkyl groups on both oxygen- and hydroxide-covered surfaces were also studied. 1. INTRODUCTION The partial oxidation of alkanes is an important industrial reaction for the manufacturing of oxygenated hydrocarbons such as alcohols, aldehydes, and ketones, which in turn are the precursors for the synthesis of higher molecular-weight hydrocarbons. A variety of metal- and metal oxide-based catalysts can be used to oxidize saturated hydrocarbons, but the challenge to successfully producing oxygenated products is in stopping the reaction before total oxidation occurs [ 1]: while the partial oxidation of hydrocarbons is thermodynamically feasible, complete oxidation to carbon dioxide and water is far more energetically favorable, so high yields for partial oxidation products can only be achieved by controlling the kinetics of both pathways. Much research in both the catalytic and surface science communities has been aimed at determining the conditions which favor partial over total oxidation, i.e., at finding ways to selectively control the reaction products [2]. It is generally accepted that the initial rate-limiting step in the partial oxidation of alkanes is the initial scission of a C-H bond to generate alkyl species on the catalyst surface [3, 4]. Alkanes have very low sticking coefficients, and therefore require the high temperature and pressure conditions used in industrial catalytic processes for their activation. Given that these conditions are difficult (if not impossible) to emulate under the ultra-high vacuum environment (UHV) normally used in surface science studies, the C-H activation step needs to be bypassed there in order to reach reasonable coverages of alkyl species on metal surfaces. One approach to achieve this is via the adsorption and decomposition of the corresponding alkyl iodides, since the C-I bonds can be readily activated to yield significant concentrations of the desired alkyl surface species [5-8]. A variety of alkyl moieties can be produced on # Funded by a grant from the Department of Energy, Basic Energy Sciences. Present address: Department of Chemistry, Canisius College, Buffalo, NY 14208, USA. * Corresponding author.

236 surfaces this way, the selection being only limited by the availability of the corresponding alkyl iodide precursors. In the present study the reaction of 2-propyl iodide with oxygen was investigated on Ni(100) surfaces. A variety of products were observed to desorb from the O/Ni(100) system, andthe selectivity among them was found to depend strongly on the coverage of oxygen. For low oxygen coverages acetone is produced in addition to hydrogen, propane, and propene, but at oxygen surface concentrations close to monolayer saturation (0.50 ML) neither acetone nor hydrogen are detected, and the amounts of propane and propene produced from the iodide are substantially reduced. Furthermore, high oxygen exposures lead to oxidation of the nickel surface [9], at which point no hydrocarbons desorb at all; only the products associated with total oxidation, namely, CO, CO2, and H20, are observed. It was also determined that the C-I bond-scission occurs in a temperature range similar to that on the clean metal, between 120180 K, suggesting that 2-propyl fragments are created on the surface at low temperatures, and that the presence of adsorbed oxygen does not alter the kinetics of that dissociation step. These results are quite different to those from the O/Rh(111) system [ 10, 11 ], indicating that a different reaction mechanism operates on nickel. In particular, we propose that the conversion to 2-propyl iodide involves the formation of a 2-propoxide species, which then yields acetone via a rate-limiting 13-hydride elimination step. 2. E X P E R I M E N T A L All experiments were done in a stainless-steel ultra-high-vacuum chamber with a base pressure of lx 10-10 Torr equipped to do temperature-programmed desorption (TPD), X-ray photoelectron (XPS), and ion-scattering (ISS) spectroscopies [12, 13]. TPD spectra were obtained by simultaneously monitoring the mass spectrometer signal of up to 15 masses with an interfaced computer while heating the sample at a rate of 10 K/s. XPS spectra were taken by using an A1 anode and a hemispherical electron-energy analyzer with an overall energy resolution of about 1.2 eV full width at half maximum. The binding energy scale was calibrated against the Pt 4f7/2 and Cu 2p3/2 core levels. ISS spectra were obtained using the same hemispherical energy analyzer as for XPS but with the voltage biases reversed to detect ions rather than electrons, and with a scattering geometry such that the angle between the ion source and the analyzer was 115 ~ [ 14]. A 1-2 ~A 500 eV He+ ion beam was focused to a spot size of about 2 mm diameter on the crystal, and the kinetic energy of the scattered ions was monitored with an interfaced computer. Control experiments on oxygen-covered Ni(100) surfaces, prepared by dosing 3.0 L of 02 at 300 K, showed that sputtering by the 500 eV He+ beam was insignificant [ 15]. Finally, because of the potential problems with the quantitative use of ISS, several methods were used to calibrate the ISS signal, and the conclusion was reached that work function changes due to the various adsorbates (which ultimately affect the neutralization probabilities of the outgoing ions) are negligible in this system and do not influence the quantitative analysis of the data reported here in any significant fashion [ 16]. The nickel (100) single crystal was cut, oriented, and polished using standard procedures, and mounted on a manipulator by spot-welding it to tantalum support wires in contact with a liquid nitrogen reservoir; the sample could be resistively heated to 1200 K and cooled rapidly back to 90 K with this arrangement. The surface temperature was monitored by a chromel-alumel thermocouple spot-welded to the edge of the crystal. Surface cleaning was done by cycles of Ar+ ion bombardment and annealing to 1200 K as well as oxygen treatments (to remove atomic carbon) until no impurities were detected by XPS or ISS. The 2-propyl iodide was obtained from Alfa Products (98% purity), protected from light, and subjected to several freeze-pump-thaw cycles before using; its purity was routinely checked by mass spectrometry. Compressed oxygen (99.999%), argon (99.999%), and hydrogen (99.999%) gases were obtained from Matheson and used as supplied. All gas exposures were

237

done by backfilling of the vacuum chamber via leak valves, and are reported in Langmuirs (1 L = lx 10-6 Torr-s), not corrected for efficiency differences in the ionization gauge. For the studies on the reaction of the alkyl iodides with O/Ni(100), the surface was prepared as follows: (1) Different amounts of molecular oxygen were dosed on the surface at 300 K, a temperature at which oxygen dissociates. (2) The sample was cooled below 100 K. (3) In the case of OH covered surfaces, the hydroxide groups were prepared by then dosing water. (4) The 2-propyl iodide was adsorbed. A constant 4.0 L exposure of the iodide was used for the TPD studies as a function of oxygen coverage. (5) For the XPS and ISS annealing studies, the sample was finally heated rapidly to the indicated temperature, at a rate of nearly 10 K/s in order to reproduce the surface conditions that exist during the TPD experiments, and immediately allowed to cool to below 100 K to freeze the relevant intermediates. 3. R E S U L T S AND DISCUSSION

3.0L

02+ xL 2-C3H71/Ni(100)

1 --

~"T?",

I:1. o

: 0Ni

ISS

E. = 5 0 0 e V H e +

0

2 4 2-C3H71 Exp. I L

Fig. 1. Ion scattering spectra (ISS) as a function of 2-propyl iodide exposure on O/Ni(100) surfaces prepared by adsorption of 3.0 L of 02 at 300 K. The inset shows the ISS peak areas normalized to the signal of the Ni and O peaks for the surface with 0.0 L 2-C3H7I. These data highlight the fact that the alkyl iodides prefer to adsorb on the metal sites.

llJ I•J

A ILl Z

0.1

~^ ~ 100

2.0 4.0

~ I

i

200

300

'

i

400

Kinetic Energy / eV

'

500

The nature of the adsorption sites for 2-propyl iodide on the O/Ni(100) surfaces were probed by ISS first. A set of experiments were done as a function of 2-propyl iodide exposure for given fixed oxygen coverages in order to probe the location of the binding sites for the iodide -- either Ni or O sites. Figure 1 shows the 2-propyl iodide coverage-dependent ISS data obtained for a fixed 3.0 L pre-exposure to oxygen. The top trace, labelled 0.0 L, corresponds to the oxygen-dosed surface before any 2-propyl iodide adsorption, and is the spectrum to which all other traces are compared and normalized. The peaks at 245 and 410 eV kinetic energies correspond to O and Ni, respectively, as determined by using standard elastic collision theory and by taking into account the geometry of our system. Upon exposures of the oxygen-dosed surface to small amounts of 2-C3H7I, the Ni signal decreases, while the O signal remains almost constant. Specifically, over 80% of the Ni peak disappears after a 1.0 L alkyl iodide exposure, but the oxygen peak still retains over 60% of its initial intensity at that point. Nevertheless, most of the oxygen signal is lost after a 2.0 L 2-C3H7I dose, and no Ni or O peaks are seen after a 4.0 L exposure. The coverage dependence of the ISS signals is better illustrated in the inset, in which plots of the normalized Ni and O ISS peak intensities are displayed versus the alkyl iodide exposure. This figure highlights the selective titration of the nickel sites for low 2-propyl iodide exposures, which means that the alkyl iodide binds

238 preferentially to nickel sites. It is important to note that the exposure at which the oxygen signal starts to be significantly attenuated, around 2.0 L of 2-propyl iodide, is the same required to detect any acetone by TPD (see below). The adsorption of 2-propyl iodide is molecular at 100 K regardless of the initial oxygen coverage. This is clearly demonstrated by the I 3d XPS data obtained for 2-propyl iodide adsorbed on both clean and oxygen-precovered Ni(100) surfaces [17]. However, heating the 02 + 2-C3H7I-dosed surfaces between 120 and 180 K induces the dissociation of the C-I bond, as also determined by XPS. We propose that this bond-activation step occurs on the Ni sites, because: (1) Ni is the preferred adsorption site; (2) the temperature range for the C-I bond-cleavage on the oxygen-covered surfaces is the same as that on the clean surface; and (3) as the number of nickel sites decreases with increasing 00, the amount of 2propyl iodide that dissociates decreases. We also assume that the majority species that form at low temperatures upon the breaking of the C-I bond are 2-propyl fragments attached to nickel atoms. 2-Propyl Iodide on O/Ni(100) T P D Effect of O x y g e n on Product Selectivity u)

Clean Ni(100)

3.0 L O2/Ni(100) Submonolayer

40.0 L 02/Ni(100 ) 3 ML NiO

Metal Function

Partial Oxidation

Total Oxidation

,1-1 =m

t~

(U

!_

!_

Water

I

'

100

'

,) '

I

'

300

'

~Hydrogen '

I

'

500

'

9

I

'

700

'

Hydrogen '

100

'

'

I

'

'

'

I

'

9

'

I

'

300 500 700 Temperature / K

'

I

'

100

'

'

I

'

300

'

'

I

'

500

'

'

I

'

'

700

Fig. 2. Temperature-programmeddesorption (TPD) spectra from 4.0 L of 2-C3H71 adsorbed on Ni(100) surfaces predosed with various amounts of oxygen. Three regimes are observed for this system: (1) that for the clean nickel, where only the hydrogenation-dehydrogenationsteps typical of transition metals are seen (left); (2) that for nickel oxide, where there is little reactivity, and where only complete oxidation is observed (right); and (3) that for an intermediateoxygen surface coverage, where some partial oxidation is manifestedby the appearance of a TPD peak for acetone around 350 K (center). The reaction of the resulting 2-propyl groups with oxygen on Ni(100) was characterized as a function of oxygen coverage by temperature-programmed desorption (TPD). A variety of species were found to desorb from this surface, namely, hydrogen, water, carbon monoxide, propene, propane, carbon dioxide, acetone, and the original hydrocarbon molecule. The distribution of reaction products was found to be strongly dependent on 00, as shown in Figure 2, which summarizes the TPD profiles for the main desorbing species after three different oxygen predoses. It can be seen there that the thermal chemistry of the

239 resulting 2-propyl species on the oxygen-treated Ni(100) follows three distinct pathways depending on the nature of the oxygenated surface. On the one end, the clean nickel behaves in the same way as many other metals, that is, it promotes both 13-hydride and reductive elimination steps to yield propene and propane, respectively [7, 13, 18]. At the other extreme, NiO films as thin as 1-2 ML thick passivate the metal and induce total oxidation to CO, CO2, and H20. It is at the intermediate coverages obtained after doses of less than 10.0 L of 02 (which yield atomic oxygen coverages below approximately 80% of monolayer saturation [9]) where the most interesting chemistry is seen, because a small amount of partial oxidation to acetone is detected. Figure 3 compares the molecular (left) and acetone (right) TPD data for the reaction of 3.0 L of oxygen (approximately 0.30 ML of O atoms, or 60% of monolayer saturation) with varying amounts of 2-propyl iodide on Ni(100). A 0.5 L exposure of 2-propyl iodide leads to the desorption of hydrogen, propene and propane, but not acetone, and results in TPD traces quite similar to those obtained from the same 2-C3H7I dose on the clean surface. The onset of acetone formation is seen as a small peak around 350 K only after a 2.0 L alkyl halide dose, and the molecular desorption data shows that monolayer saturation of 2-propyl iodide on this surface occurs between 2.0 and 4.0 L. Notice in particular that the 2.0 L mark corresponds to the point at which all the nickel sites become occupied (see Figure 1). This suggests that, in order for acetone to be produced, a particular surface ensemble is required with the 2-propyl groups adsorbed next to oxygen atoms [ 19-21 ].

II

3.0 L 0 2 -I- X L 2-C3H71/Ni(100 ) T P D

. m

Molecular Desorption

:3

zi

Acetone (xl0 scale)

!._

t~ 61

L._

~t

u~ 61 In

13. t~

2"C3H71 Exposure(x) / L

\

J 2.0

t._

13.

0.5 I

'

100

'

'

I

'

300

'

'

I

'

500

'

'

I

'

700

'

I

'

100

'

'

I

'

300

Temperature / K

'

'

I

'

500

'

'

I

'

700

'

Fig. 3. Molecular (left) and acetone (right) TPD spectra from 0.5, 2.0, and 4.0 L of 2-propyl iodide adsorbed on Ni(100) pretreated with a fixed 3.0 L oxygen dose at 300 K. The results from this figure indicate that acetone production starts at the point where the nickel adsorption sites become saturated, suggesting that the proximity of alkyl and oxygen groups on the surface is a requisite for

partial oxidation reactions. Several pieces of evidence point to the fact that the next step in the reaction that leads to the formation of acetone is the insertion of an oxygen atom into the metal-carbon bond. Figure 4 shows three of them. The left frame displays ISS data obtained after annealing the 2propyl iodide + oxygen system to different temperatures. The main result from this work is highlighted in the inset, which indicates that the O ISS peak never recovers the initial intensity seen before the alkyl halide dose. This suggests that at temperatures slightly above 200 K some of the alkyl groups that remain on the surface migrate to sites on top of the

240

chemisorbed oxygen atoms, blocking the latter from the incoming probing ions. The two other frames of Figure 4 provide XPS and TPD evidence for the similar behavior of 2-propyl iodide and 2-propanol on oxygen-covered Ni(100) surfaces; since propanol is known to produce 2-propoxide groups at low temperatures [22], the same intermediate is inferred to be involved in the case of the alkyl halide. In particular, the C ls XPS traces present in both cases the small shoulder around 285.7 eV binding energy most likely associated with the carbon atom adjacent to the oxygen, and both acetone TPD traces display similar peaks around 320 K. Notice also that molecular acetone desorbs at much lower temperatures from these surfaces, which means that acetone detection in the TPD experiments with the iodo alkane is reaction (not desorption) limited. E v i d e n c e for P r o p o x i d e F o r m a t i o n from 2-Propyl Iodide on O/Ni(100) S u r f a c e P r e p a r e d by D o s i n g 3.0 L 0 2 at 3 0 0 K

ISS 1

"1

C l s XPS

half-covered

Still

i\

|

(middlecarbon in propoxide)

'=/"ee-e'------e

n

o =-./ . . . . 100 TI K

i

~1

I

300

,

, ~ . 4.o L 2-%H~1

2001 15Ol

.......

200

/~

7L~I 4.o L2-C~H,, 4L.~_~I heatedto 170 K

~.~~--~'x ,

320 K

5.0 L 2-C3H7OH ::' heated to 200 K ,:

, 700

I ~ , I.~-~--~~J'

i

Acetone TPD

,oo !

i

400

.

Kinetic Energy / eV

i

500

!"

'1

280

,

,

'

9 ' ,

i

285

.

.

.

.

i

,

290 100

Binding Energy I eV

,'

,

i

,

300

,

,

i

500

9

,

,

i

,

700

Temperature I K

,

9

900

Fig. 4. Left: ISS annealing data for 4.0 L of 2-propyl iodide adsorbed on Ni(100) predosed with 3.0 L of 02 at 300 K. The inset indicates how the oxygen ISS signal never reaches its original value from before alkyl halide adsorption, suggesting that some of the oxygen atoms become covered by alkyl groups. Center: C ls XPS data for 2-propanol (top) and 2-propyl iodide (bottom), adsorbed on oxygen-precovered Ni(100) surfaces and then annealed to induce 2-propoxide formation. The alkoxide intermediate is identified by the signal at 285.7 eV binding energy that corresponds to the middle carbon atom. Right: acetone TPD spectra from O-precovered surfaces dosed with acetone, propene, 2-propanol, and 2-propyl iodide. The similar behavior seen in the latter two cases suggest that 2-propanol and 2-propyl moieties react via a common intermediate. Also, the absence of any acetone desorption from adsorbed propene discards such a species as a possible intermediate in the partial oxidation of propyl groups in this case.

Lastly, 2-propoxide intermediates undergo a 13-hydride elimination step above 300 K to yield the final acetone product. The selectivity of this step is best illustrated by the acetone TPD traces shown in Figure 5, which were obtained with the partially labelled CD3CHICD3 isotopomer of 2-propyl iodide. The only acetone detected in these experiments is the perdeutero species, which means that the hydrogen-removing step is regiospecific and involves only the secondary middle hydrogen. These experiments also confirm that propene is not involved in the production of acetone in this system. The data presented so far suggest that a few requirements need to be met in order for partial oxidation reactions to take place on oxides, namely: (1) There is a need for metal

241 atoms to be exposed on the surface. This in fact was a requisite here only because the Ni sites are the ones that facilitate the dissociation of the C-I bond in the alkyl halides (a reaction that is not relevant for the oxidation of alkanes), but they may also be necessary to induce the initial C-H bond-activation in alkanes. (2) The migration of alkyl groups attached to metal atoms seems to be somewhat limited, which means that they need to form next to oxygen atoms for the partial oxidation process to proceed. (3) The partial oxidation reaction involves alkoxide intermediates, and therefore may be favored by anything that stabilizes such species on the surface. (4) The rate-limiting step appears to be the last 13-hydride elimination, so it is important to facilitate the fast desorption of the ketone (or aldehyde) products, because otherwise they may decompose on the surface immediately after their formation.

4.0 L CD3CHICD 3 on O/Ni(100), TPD Surface prepared by dosing 3.0 L 02 at 300 K CD3Ncll/CD3 ~ ........

-1/2H2

Ni

~

O

>300 K

Ni

Acetone

Propoxide

i

100

'

'

'

i

Fig. 5. Acetone TPD traces for the main isotopomers expected from the reaction of CD3CHICD3 with oxygen on Ni(100) surfaces. The exclusive formation of perdeutero acetone in this case indicates the high selectivity towards a 13-hydride elimination step from the 2-propoxide intermediate, and rules out a mechanism where an initial 13-hydrideelimination from 2propyl groups on Ni sites is followed by oxygen incorporation.

CD3Nf/C~

'

'

'

i

'

'

'

!

,

300 500 700 Temperature / K

It should be noted that the experiments reported here were carried out on oxygendosed nickel surfaces, not on nickel oxide. Nevertheless, it appears that the conclusions enumerated above may be applicable to substoichiometric oxide surfaces as well [19-21]. The possibility of extrapolating our surface science studies to more realistic systems is supported by the data in Figure 6, which shows TPD data illustrating the ability of different surfaces towards acetone production from propyl moieties. As mentioned before, the clean nickel surface only yields propane and propene; oxygen atoms are needed to produce oxygenated products (bottom trace). The second trace from the bottom reproduces the acetone desorption profile shown in Figure 2 for surfaces covered with oxygen submonolayers, which are active systems for partial oxidation processes. Next up there is another flat trace, that obtained from a stoichiometric nickel oxide film: no acetone is produced there because of the inert character of that surface. It is the last (top) trace the one that shows the promise of our studies. The NiO film was annealed in this case to high (>600 K) temperatures prior to alkyl halide dosing in order to induce the diffusion of some oxygen atoms into the bulk and the formation of a substoichiometric oxide top layer. The TPD data proves that such a surface is still somewhat active towards acetone formation. This appears to imply that even subsurface oxygen may be capable of migrating to the surface in order to insert itself into the Ni-C bond during the 2-propoxide formation step. One final observation is worth reporting here, that is, that concerning the role of OH surface groups in the partial oxidation process. Figure 7 provides TPD traces for the respective desorption of acetaldehyde and acetone from ethyl and 2-propyl iodides adsorbed

242

on oxygen- (left) and hydroxide- (right) covered nickel surfaces. Two important observations can be highlighted from these data: (1) The formation of acetone from 2-propyl groups is enhanced by the presence of OH groups on the surface, to the point of yielding TPD traces similar to those seen for the case of 2-propanol (compare with Figure 4); and (2) Some acetaldehyde is produced during the oxidation of ethyl groups if hydroxide species are present on the surface. In contrast, no aldehydes were detected in TPD experiments with methyl, ethyl, 1-propyl, or 1-butyl iodides on purely oxygen-precovered nickel surfaces, perhaps because the resulting alkoxides decompose on the surface before yielding the desired products. The mechanism by which the OH groups enhance the ability of the surface to induce partial oxidation reactions is still under investigation.

Effect of Oxygen Environment on Partial Oxidation Selectivity Acetone TPD from 2-C3H71 on O/Ni(100) 355 K

Fig. 6.Acetone TPD for 2-propyl iodide adsorbed on Ni(100) surfaces after different oxygen treatments. The bottom and third traces indicate that neither the clean nickel metal nor stoichiometric NiO films are capable of inducing the partial oxidation of 2propyl groups to acetone. In contrast, the second and four traces show that substoichiometric nickel oxides are able to promote such a reaction even if the oxygen atoms are in the subsurface region.

Subsurface

x8

Oxide 55 K

Clean metal I

100

'

'

'

I

'

'

300

'

I

500

'

'

'

I

700 Temperature / K

'

'

'

I

900

4. CONCLUSIONS It was shown here that, under the right conditions, substoichiometric nickel oxides may be capable of catalyzing partial oxidation reactions. A number of criteria were identified for establishing the activity of this pathway, namely: (1) Nickel atoms need to be exposed in order to promote the initial alkane activation and to stabilize the resulting alkyl surface groups. (2) Oxygen atoms need to be present in the proximity of the alkyl groups in order for the insertion step that leads to alkoxide formation to take place. This oxygen could be located in the immediate subsurface region. (3) Alkenes appear to not be direct intermediates in the conversion of alkyl groups to aldehydes or ketones. This is an interesting observation, because alkenes can indeed be converted to such products catalytically [23]. It could be suggested that perhaps alkenes convert to alkyl groups on the surface before undergoing oxygen incorporation. (4) Surface hydroxide groups appear to enhance the partial oxidation pathway, either because they facilitate the formation of alkoxides, or because they help in the limiting H-abstraction step. (5) The formation of ketones seems to be easier than the production of aldehydes. This is somewhat encouraging, because secondary C-H bonds are weaker than primary ones and therefore easier to break; alkane activation is likely to yield branched alkyl surface moieties.

243 | !

Effect of OH Surface Groups on Selectivity for Partial Oxidation TPD on O/Ni(100) 3 L 0 2 dose at 300 K

TPD on O+OH/Ni(100) 0.5 L 0 2 4- 2.0 L H20

~

360 K

355 K ,tv~'~

acetone from 3.0 L 2-C H_,l

/ -/ J

acetaldehyde from 2.0 L C2H51A 228 K 100

300

500

700

900100

300

500

700

900

Temperature / K Fig. 7. TPD traces from alkyl iodides adsorbed on oxygen- (left) and hydroxide- (right) precovered Ni(100) surfaces. The bottom traces correspond to the formation of acetaldehyde from ethyl iodide, while those on top display the desorption of acetone from 2-propyl iodide conversion. The enhancing power of OH surface groups towards partial oxidation pathways is indicated by two observations from these data: (1) the yield for acetone increases to the point of resembling that seen with 2-propanol; and (2) some acetaldehyde is detected as well. It is at the present time unclear if the OH groups favor the formation of alkoxide intermediates or the subsequent [3-hydride elimination step.

A s c h e m a t i c representation of the m e c h a n i s m p r o p o s e d here for the partial oxidation of 2-propyl iodide on o x y g e n - c o v e r e d Ni(100) surfaces is p r o v i d e d as S c h e m e 1 below. W e believe that s o m e of these ideas m a y be applicable not only to other alkyl groups, but also to other oxide surfaces and to more realistic catalytic conditions. CH3CH2CH3 (g) CH3CH2CH3(g) +

CH3~HCH3 I O

I

CH3CH=CH2(g) 225 K xO CnH

CH3CH=CH2 ~)

/CH3CHCI-~

/ /-

.o

200 K

CH3"rH / CH3

Migration

-150-200K

O

CH3 13'--H

~

>400

H2(g) + C O ~NI~

CH3 C ~)

> 3 0 0 K ' - ~

CH3x /CH3 C (g) 3SSK--

Propoxide

l Hydroxyl proton abstraction CH3CHCH3 O

I

OH

Scheme 1. Proposed mechanism for the partial oxidation of 2-propyl surface species on O/Ni(100) at low oxygen coverages. A similar mechanism is likely to operate during the catalytic conversion of alkanes on substoichiometric nickel oxides.

244 REFERENCES

[1] [21 [3] [41 [51 [6] [71 [81 [9] [10] [11] [12] [13] [ 14] [15] [16] [ 17] [18] [19] [20] [21] [22] [23]

R. J. Madix and J. T. Roberts, in Surface Reactions, Eds. R. J. Madix (Springer-Verlag, Berlin, 1994) pp. 2. J. H. Lunsford, Angew. Chem. Int. Ed. Engl. 34 (1995) 970. G. A. Somorjai, Catal. Rev.-Sci. Eng. 23 (1981) 189. R. Pitchai and K. Klier, Catal. Rev.-Sci. Eng. 28 (1986) 13. F. Zaera, Acc. Chem. Res. 25 (1992) 260. X.-L. Zhou, X.-L. Zhu and J. M. White, Acc. Chem. Res. 23 (1990) 327. F. Zaera, Chem. Rev. 95 (1995) 2651. B. E. Bent, Chem. Rev. 96 (1996) 1361. C. R. Brundle, in The Chemical Physics of Solid Surfaces and Heterogeneous Catalysis, Vol. 3A (Chemisorption Systems), Eds. D. A. King and D. P. Woodruff (Elsevier, Amsterdam, 1990) pp. 132. C.W.J. Bol and C. M. Friend, J. Phys. Chem. 99 (1995) 11930. C.W.J. Bol and C. M. Friend, J. Am. Chem. Soc. 117 (1995) 11572. F. Zaera, Surf. Sci. 219 (1989) 453. S. Tjandra and F. Zaera, J. Am. Chem. Soc. 117 (1995) 9749. F. Zaera, Langmuir 7 (1991) 1188. N.R. Gleason and F. Zaera, J. Catal. (1997) in press. N.R. Gleason and F. Zaera, Surf. Sci. (1997) in press. N. R. Gleason and F. Zaera, 211 th American Chemical Society National Meeting, New Orleans, 1996, Paper No. COLL. S. Tjandra and F. Zaera, Langmuir 10 (1994) 2640. J. L. Callahan and R. K. Grasselli, AIChe J. 9 (1963) 755. R. K. Grasselli and J. D. Burrington, Adv. Catal. 30 (1981) 133. J. Nilsson, A. R. Lana-Canovas, S. Hansen and A. Anderson, J. Catal. 160 (1996) 244. X. Xu and C. M. Friend, Surf. Sci. 260 (1992) 14. H.H. Kung, Adv. Catal. 40 (1994) 1.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

245

S t r u c t u r e and Catalysis of LixNi2_xO2 Oxide Systems for Oxidative Coupling of M e t h a n e T. l~liyazaki a, T. Doi b, T. M i y a m a e c, and I. M a t s u u r a a a D e p a r t m e n t of C h e m i s t r y , F a c u l t y of Science, T o y a m a University, Gofuku, T o y a m a 930, J a p a n byokkaichi R e s e a r c h L a b o r a t o r y , Tosoh Corp., Yokkaichi 510, J a p a n CInstitute for M o l e c u l a r Science, Myodaiji, Okazaki 444, J a p a n

P h o t o e l e c t r o n s p e c t r o s c o p y ( X P S and U P S ) studies have b e e n applied to t h e investigation of active oxygen species of solid solution series LixNi2_ • w i t h wide range of 0<x_< 1 for u n d e r s t a n d i n g t h e origin of high a c t i v i t y and C2selectivity of t h e oxidative coupling of m e t h a n e . F r o m the Comparison of these s p e c t r a l features, the surface oxygen of Li0.3Nil.702 w i t h a cubic s t r u c t u r e was assigned to 0 2 - species, while t h a t of LiNiO2 w i t h a h e x a g o n a l s t r u c t u r e was a t t r i b u t e d to O- species a s s o c i a t e d w i t h t h e c a t a l y t i c a c t i v i t y for the oxidative coupling of m e t h a n e . F r o m the r e l a t i o n s h i p b e t w e e n these s t r u c t u r e s a n d the c a t a l y t i c activity, it was inferred t h a t t h e a n i s o t r o p i c layered s u p e r s t r u c t u r e of LiNiO2 largely affected the f o r m a t i o n of the active oxygen species for the oxidative coupling reaction.

1. I N T R O D U C T I O N C a t a l y s t s used for t h e oxidative coupling of m e t h a n e ( O C M ) have b e e n widely i n v e s t i g a t e d for t h e chemical utilization of n a t u r a l gas. T h e role of t h e c a t a l y s t is p r i m a r i l y to provide a r e a c t i o n p a t h w a y which facilitates the a b s t r a c t i o n of one h y d r o g e n from m e t h a n e w i t h surface oxygen species while s u p p r e s s i n g deep oxidative reactions. Since Keller and B h a s i n i n v e s t i g a t e d t h e possibility of s y n t h e s i z i n g C 2 - h y d r o c m ' b o n s by t h e oxidation of m e t h a n e using an oxide c a t a l y s t , m a n y different c a t a l y t i c s y s t e m s have b e e n e x a m ined [1]. Interestingly, all r e p o r t e d selective c a t a l y s t s t o w a r d C 2 - h y d r o c a r b o n s in m e t h a n e conversion contain an alkali c o m p o n e n t . One of the m o s t active and selective c a t a l y s t is m a g n e s i u m oxide d o p e d w i t h alkali m e t a l ions such as Li 4. A l i t h i u m - d o p e d m a g n e s i u m oxide, L i / M g O , has b e e n extensively s t u d i e d as a c a t a l y s t for the O C M reaction. Lunsford et al. p r o p o s e d t h a t the active centers of L i / M g O c a t a l y s t is [Li ~O-] species on the c a t a l y s t surface [2]. O t s u k a et al. also r e p o r t e d t h e s t u d y of alkali-doped t r a n s i t i o n m e t a l

246 oxides [3]. This r e p o r t p o i n t e d out t h a t LiNiO2 contained reducible comp o n e n t s h a d good efficiency w i t h this reaction b a s e d on a redox m e c h a n i s m involving lattice oxygen anions. Gaffney et al. s u g g e s t e d t h a t N a p r o m o t e d P r 6 O l l c a t a l y s t forms N a 2 0 2 ( 0 2 2 - , peroxide anion) as an active species by a redox reaction w i t h P r 6 O l l [4]. T h e selectivity a n d a c t i v i t y of the O C M r e a c t i o n is m a i n l y d e t e r m i n e d by the n a t u r e of surface oxygen species, which is affected by t h e electrical p r o p e r t i e s of the m e t a l oxides used as c a t a l y s t . T h u s , a d e t a i l e d s t u d y of t h e oxygen species on the c a t a l y s t surface is essential for u n d e r s t a n d i n g t h e origin of t h e catalysis. Accordingly, t h e p u r p o s e of this investigation is to m a k e out t h e r e l a t i o n s h i p b e t w e e n t h e a c t i v i t y of the O C I ~ a n d surface oxygen species of LiNiO2 c a t a l y s t s y s t e m . An u n d e r s t a n d ing of t h e s e c a t a l y s t p r o p e r t i e s and t h e i r c a t a l y t i c p e r f o r m a n c e would m o s t p r o b a b l y provide a f u n d a m e n t a l knowledge for designing c a t a l y s t s to increase C2-selectivity. T h e s t r u c t u r e of LiNiO2 are classified into the hexagonal t y p e , l i t h i u m a n d nickel ions in LiNiO2 a l t e r n a t e l y c o n f i g u r a t e d on the (111) planes of rock-salt type. F r o m a different point of view, electrochemically, LiNiO2 has a t t r a c t e d m u c h a t t e n t i o n due to t h e c a p a b i l i t y of i n t e r c a l a t i o n w i t h alkali m e t a l s a n d a p p l i c a t i o n to l i t h i u m - b a s e d b a t t e r i e s [5]. In this p a p e r , we r e p o r t a s t u d y of t h e LixNi2_xO2 s y s t e m c a t a l y s t s w i t h rock-salt s t r u c t u r e . C a t a l y s t s were c h a r a c t e r i z e d by X R D a n d X P S and t e s t e d in t h e oxidative coupling of m e t h a n e in order to e l u c i d a t e the different s t r u c t u r e and surface species.

2. E X P E R I M E N T A L T h e l i t h i u m nickelate(III) c a t a l y s t used in the e x p e r i m e n t s was p r e p a r e d by a solid-state reaction of an i n t i m a t e m i x t u r e of LiNO3 and Ni(OH)2 . T h e L i - i m p r e g n a t e d Ni(OH)2 was calcined in air at 873K for 2h a n d s u b s e q u e n t l y at 1073K for 6h. T h e s y n t h e s i z e d p o w d e r were identified by m e a n s of powder X - r a y diffraction ( X R D , S h i m a d z u XD-3A) [6]. T h e O C M r e a c t i o n was carried out in a fixed-bed r e a c t o r using 0.3g of the c a t a l y s t a n d was monit o r e d by gas c h r o m a t o g r a p h y . T h e flow r a t e s and t h e m i x e d ratio of gas, for all e x p e r i m e n t s , were u n d e r the a t m o s p h e r i c p r e s s u r e as follows 95 0 m l / m i n ; CH4:O2:He----2:l:3, respectively. T h e electronic d e n s i t y of LiMO2 ( M - - N i , Co, Fe, a n d In) a n d n o n s t o i c h i o m e t r i c LixNi2_xO2 (0_<x_
247 the Fermi edge of gold films e v a p o r a t e d in situ.

3. R E S U L T S

AND

DISCUSSION

3.1 C r y s t a l s t r u c t u r e of LixNie_xO2 X-ray diffraction p a t t e r n s of the solid solution series LixNi2_xO2 are shown in F i g u r e 1. FI'om the p o w d e r X-ray diffraction p a t t e r n s , the s t r u c t u r e of L i x N i 2 _ x O 2 ( 0 < x < l ) could be s t r u c t u r a l l y classified into some groups, t h o u g h the basic s t r u c t u r e of" these c o m p o u n d s consists of the rock-salt type. T h e solid solution series LixNi2_xO2 ( 0 < x < 0 . 6 ) , Lio.3Nil ~O2 and Lio.6Nil.402 are shown in F i g . l - ( a ) and F i g . l - ( b ) , have a typical cubic t y p e s t r u c t u r e analogous to the s t r u c t u r e of NiO. On the o t h e r hand, the X R D p a t t e r n s of LiNiO2 (Fig. l-e) can be seen by the a p p e m ' a n c e of the s u p e r l a t t i c e peaks at t w o - t h e t a values of 18.33 ~ (003), 36.2 ~ (101), 48.5 ~ (105), and 58.2 ~ (107) and (009) in a h e x a g o n a l setting. C o m p a r e d with these X R D p a t t e r n s , LixNi2_• for the range of 0 . 6 < x < 1 , Li0.rNil.302, Li0.9Nil.lO2 are shown in Fig.l-(c) and F i g . l (d), m a y be an incomplete hexagonal type. F r o m these results, the lattice s t r u c t u r e t r a n s f o r m e d from a cubic t y p e to a hexagonal t y p e as lithium cont e n t was increasingly doped. T h e r e is considerable evidence to show t h a t the t r a n s f o r m a t i o n should occur at x--0.6 to be shown in F i g . l - ( b ) . T h e ionic configuration of the solid solution series LixNi2_xO2 are w r i t t e n by the formula Lix(Ni 3. )x(Ni 2. )2_x02 . NiO c o r r e s p o n d s to LixNi2_xO2 ( x ~ 0 ) is classified into

{003)

(107), (009) /

(006), (102)

I {lO4)

(e) ('~

I

I (1o8)

l('os)

I.t

,,.A_._...~L--~ . . . . . . .

(d)

o..

]

(,,o)

9

A

,

b 1.

(b)

g~

w

l { .....

(a) .

dl

10

A

I(220)

( I L _ ~ (200) ,

i

20

,

I.

30

9

.I

.,

40

I

50

,

l

60

9

J

70

2theta / deg. F i g u r e 1. X-ray diffraction p a t t e r n s of the solid solution LixNi2_xO2. (a) Lio.3Nil.rO2; (b) Lio.6Nil.402; (c)Lio.TNil.302; (d) Lio.9Nil.102" (e) LiNiO2.

248 a cubic s t r u c t u r e in w h i c h nickel cations a n d o x y g e n anions are c o n f i g u r a t e d on t h e (111) planes of t h e rock-salt t y p e . C r y s t a l s t r u c t u r e of L i N i 0 2 a n d LixNi2_xO2 solid solution are s h o w n in F i g u r e 2. NiO has a rock-salt t y p e s t r u c t u r e in which Ni cations are c o o r d i n a t e d in t h e s i x - c o o r d i n a t e l o c a t i o n f o r m e d in t h e cubic closest p a c k i n g (ccp) of oxygen ions. As for LixNi2 xO2 solid s o l u t i o n in w h i c h Li ions are dissolved into NiO, t h e solid solution w i t h t h e r a n g e of x<.0.6 has a cubic s t r u c t u r e , t h a t is a d i s o r d e r e d rock-salt s t r u c . . . .

(c)

0

0

o Ni

9 Li

9 Ni or Li

F i g u r e 2. T h e s t r u c t u r e of t h e rock-salt s t r u c t u r e . (a) C u b i c s t r u c t u r e , (b) I n c o m p l e t e h e x a g o n a l s t r u c t u r e a n d (c) H e x a g o n a l s t r u c t u r e . 573K 773K

2931

1093K

Lio.7NiuO 2

. 893K [293Kl "~ I- I 11033K

II

" ~

--~

i

9

l0

il II

[[ f[I

II03K

.

l

20

I

Lil.oNil.oO2

9

I

30 40 2 theta / deg.

i

i

50

60

=-

ar

.

9

10

.

I

20

9

~

!

9

a

30 40 2 theta / deg.

50

60

F i g u r e 3. T h e t e m p e r a t u r e d e p e n d e n c e on X R D p a t t e r n s of Lio.rNil.302 a n d LiNiO2 c a t a l y s t s . t u r e , this c r y s t a l p h a s e is r e p l a c e d by one in which t h e Li a n d Ni c a t i o n s are s e g r e g a t e d on a l t e r n a t e (111) planes a n d are r a n d o m l y d i s t r i b u t e d over t h e available o c t a h e d r a l m e t a l sites of t h e rock-salt s t r u c t u r e ( s e e Fig. 2-a). On t h e o t h e r h a n d , solid s o l u t i o n s w i t h t h e r a n g e of 0 . 6 < x < l have a h e x a g o n a l

249 layered s t r u c t u r e w i t h t h e oxygen layers ((111) direction of t h e rock-salt t y p e s t r u c t u r e ) m i d w a y b e t w e e n cation layers and these cations are a r r a n g e d in a l t e r n a t i n g layers of p u r e Ni a n d m i x e d layers of r a n d o m l y d i s t r i b u t e d Li and Ni cations (see Fig. 2-b). T h e s t r u c t u r e of LiNiO2 has a h e x a g o n a l s y s t e m (~-LiFeO2 t y p e s t r u c t u r e in which M 3~ (Ni) cation layers and Li cation layers are a l t e r n a t e l y a n d r e g u l a r l y a r r a n g e d to oxygen ion layers (see Fig. 2-c). T h e t e m p e r a t u r e d e p e n d e n c e of X R D p a t t e r n s of Li0.rNil.302 a n d LiNiO2 are shown in F i g u r e 3. T h e i n c o m p l e t e h e x a g o n a l s t r u c t u r e was t r a n s f o r m e d to t h e cubic lattice at 973K~ while t h e s t r u c t u r e of LiNiO2 was held on t h e h e x a g o n a l t y p e above the t e m p e r a t u r e at 1073K. 3.2 C2-selectivity of the OCM reaction T h e results of m e t h a n e conversion and the selectivity of C 2 - h y d r o c a r b o n s are c o m p a r e d in Table 1. LixNi2_xO2 c a t a l y s t s w i t h different l i t h i u m cont e n t for t h e range of 0 < x < l . 0 were p r e s e n t e d . It is k n o w n t h a t the selective O C M reaction g e n e r a l l y p r o c e e d s only at t e m p e r a t u r e above ca. 900K. T h e intrinsic r e q u i r e m e n t of such high t e m p e r a t u r e for t h e OCIVI r e a c t i o n is still not fully u n d e r s t o o d . T h e c a t a l y t i c r u n s were carried out u n d e r conditions in which t h e C2-selectivity was achieved at the m a x i m u m . F i g u r e 4 shows a plot of the selectivity of C 2 - h y d r o c a r b o n s as functions of l i t h i u m c o n t e n t s x in Li• T h e f o r m a t i o n of C 2 - h y d r o c a r b o n s is initially observed at t h e l i t h i u m c o n t e n t of x - 0.2. F r o m these results~ it is found t h a t the solid solutions of LixNi2_xO2 w i t h x > 0 . 6 having t h e h e x a g o n a l s t r u c t u r e b e c o m e a c a t a l y s t which has high selectivity to t h e O C M reaction. However~ even in the Table 1 C a t a l y t i c a c t i v i t y of t h e oxidative coupling of m e t h a n e over LixNi2_xO2. Catalyst NiO Li0.3 Ni 1.rO2 Li0.6 Ni 1.4 02 Li0.rNil.3 02 Li0.9Nil.lO~ LiNiO2

React. Temp. (K)

CH4- Conv. (%)

C2-Select. (%)

Structure

993 1053 993 1053 993 1053 993 1053 993 1053

21.4 33.8 31.7 33.6 33.3 34.1 37.2 36.1 32.2 42.5

32.0 27.3 40.0 25.4 54.1 20.7 51.4 25.1 53.7 63.6

Cubic C u b ic Cubic Cubic Cubic Hexagonal Cubic Hexagonal Cubic Hexagonal Hexagonal

case of these selective catalysts~ its selectivity d r o p s a b r u p t l y at a high t e m p e r a t u r e at 1053K as shown in Table 1. T h e e x p e r i m e n t a l result over Li0 7Nil.302 is shown Figure 5. T h e s t r u c t u r e of LixNi2_xO2 ( 0 . 7 < x < 0 . 9 ) t r a n s f o r m e d from

250

---Hexagonal type --~

Cubic type

v

70

I!!:: :::i(

60 k,

iii l i ! )...e...~--o--'~ 'i ,ii,i[l

50

! !

9-. 40 ..~ -~ 3O rfl ',.,, ~ 20

,

.

,

,I

.......

:~-~*~_~:

l

a. ~**

~ ~i~,!

D

,

i~i:: ::!

,

l:L::::i

10 00

0.0

0.2

0.4 0.6 X (LixNi2_xO2)

Figure 4. The relationship between LixNi2_xO2 ( 0 < x < l ) catalyst 9 100

A

0.8

C2-selectivity

1.0

and the structure of

O2-conv. O..._o

80

t=,

A

60 .~

C2-select.

.,~

~9

4020

yo / O / o / CH4-conv.

O

.._...O 900

950

1000

1050

Temperature / K

Figure 5. The temperature d e p e n d e n c e catalyst.

o f methane oxidation o v e r L i 0 . ? N i 1 . 3 0 2

251 a hexagonal t y p e to a cubic t y p e at higher t e m p e r a t u r e t h a n 973K. It should be noted t h a t the C2-selectivity may be correlated to the crystal s t r u c t u r e of these catalysts. The catalytic activity of the O C M reaction should be predictable by the n a t u r e of the oxygen species on the catalyst surface. In addition, the electronic states of the surface oxygen could be largely related to the bulk s t r u c t u r e in the crystal.

3.3 A c t i v e sites for t h e O C M r e a c t i o n A typical p h o t o e l e c t r o n spectroscopy e x p e r i m e n t was conducted as follows for u n d e r s t a n d i n g of the catalyst surface. The characterization of the surface oxygen species on these oxides were investigated using XPS and UPS. Figure 6 displays representative typical X P S spectra in the O ls region for Lil.0Nil.002 (Fig. 6-a), Li0.3Ni1.702 (Fig. 6-b) and Ni304 (Fig. 6-c). The O ls spectra in Fig. 6 reveal the presence of at least two different types of the near-surface oxygen species on these catalysts. Two structures, denoted by the characters A and B, were observed at 529.4 and 531.3 eV in the XPS O ls s p e c t r a of Ni304, and at 529.2 and 531.4 eV in those of Li0.3Nil.702. The peak A located at 529.3+0.1 eV could be a t t r i b u t e d to 0 2 - species which the catalytic n a t u r e must be nonactive for the O C M reaction [8]. It seems reasonable to u n d e r s t a n d t h a t the main existence of the 0 2 - species on Ni304 oxides are nonselective for the O C M reaction. A n o t h e r O ls peaks (peak B) at higher binding energy in these oxides may be associated with the catalytic

A

Ols

~

(c)

(b)

m

(a) I

538

9

I

536

,

I

,

I

,

I

.

I

,

i - - ,

530 528 526 Binding Energy / eV

534

532

I

524

,

I

522

Figure 6. X P S O ls spectra of LiNiO2 system. (a) LiNiO2; (b) Lio3Nil.rO2; (c) Ni3Oa.

252 activity for the O C M reaction. T h e peak intensity ratio A / B of the Li0.3Nil.rO2 was given to be 36/64 in the last column on Table 2, and the oxygen species of the peak B is reasonably existence on the Li0.3Ni1.rO2 catalyst surface. B u t this catalyst was low activity for the O C M reaction as described previous section. In t h e case of LiNiO2 with the active and high C2-selectivity catalyst for the O C M reaction, a peak B' in the XPS s p e c t r u m of LiNiO2 was observed at Table 2 The O ls level of the binding energy of LixNi2_xO2 catalysts. Crystal

Structure

Ni304 Li0.3Nil.rO2 Li0.vNil.302 LiNiO2

Cubic Hexagonal Hexagonal

Ols (eV) Peak A 529.4 529.2 529.1 529.2

Ols (eV) Peak B 531.3 531.4 531.6 531.9

Intensity ratio A/B 63/37 36/64 21//79 11//89

531.9 eV to higher binding energy level of the peak B and largely related to the selective O C M reaction. The peak B' was also observed at 531.8 eV in the X P S s p e c t r u m of LiCoO2 which exhibit high .activity for the O C M reaction [9]. The slight difference in the binding energy between the peak B and the peak B' m a y be often a b e t t e r indication of the n a t u r e of oxygen species r a t h e r t h a n the absolute value of the binding energy. Thus, the a p p e a r a n c e of the oxygen species is a t t r i b u t a b l e to the high C2-selectivity for the OC]~I reaction. C

r~

m

(b)

(a) "

20

I

15

"

"

1'0

I

5

"

'

EF=()

Binding energy / eV Figure 7. U P S s p e c t r a of the solid solution LixNi2_xO2. (a) Lio.3Nil.rO2; (b) LiNiO2.

253 T h e results of X P S m e a s u r e m e n t for t h e solid solution series LixNi2_xO2 are s u m m a r i z e d in Table 2. T h e peaks B of Li07Ni1302 w i t h a i n c o m p l e t e hexagonal s t r u c t u r e was also observed at 531.6 eV to slight higher b i n d i n g energy. T h e i m p o r t a n t point to note t h a t the crystal s t r u c t u r e of these c a t a l y s t also largely affected to t h e electronic states of the surface oxygen. Figure 7 shows U P S s p e c t r a of Li0.3Ni1702 and LiNiO2. These s p e c t r a were m e a s u r e d at t h e incident hp of 40 eV w i t h reference to EF as t h e zero on the e n e r g y scale. Several structures~ d e n o t e d by the c h a r a c t e r s A-C~ were observed at 2.0, 6.0, and 10.5 eV in those of LiNiO2 (Fig. 6-a)~ a n d at Eb ---- 2.2, 6.0 and 11.0 eV in t h e U P S s p e c t r a of Li0.3Nil.702 (Fig. 6-b), T h e s t r u c t u r e located at Eb -- 2.0~2.2 eV can be assigned the 3d b a n d of nickel oxide~ a n d the b r o a d peak at 6.0 eV in the valence b a n d can be assigned to the O 2- species associated w i t h NiO [10]. It was suggested from these results t h a t the bulk s t r u c t u r e largely affected t h e surface oxygen species which play an i m p o r t a n t role for the selective of t h e O C M reaction. Figure 8 shows the location of oxygen ions, Ni and Li cations viewed t w o - d i m e n s i o n a l l y along the vertical direction of the cubic closest packing of oxygen ions of LixNi2_xO2 series. As l i t h i u m c o n t e n t was increasingly d o p e d , t h e lattice s t r u c t u r e t r a n s f o r m e d from a cubic t y p e (Fig. 8-a) to a hexagonal t y p e (Fig. 8-c) t h r o u g h a interm e d i a t e t y p e (Fig. 8-b) b e t w e e n these s t r u c t u r e . T h e oxygen species in the cubic lattice is placed in a h o m o g e n e o u s charge d e n s i t y w i t h a r a n d o m mixed Li ~~ Ni 2 ~ and Ni 3~ . On the o t h e r hand, the oxygen species in a hexagonal lattice is placed in a anisotropic charge density b e t w e e n m o n o c a t i o n Li ~ layer and t r i c a t i o n Ni 3~- layer. It is inferred from the relationship b e t w e e n t h e i r s t r u c t u r e of LixNi2_xO2 a n d t h e selectivity of the O C M t h a t Li + and Ni 3~ cations t w o - d i m e n s i o n a l l y configurated as shown in Fig. 8-c is the i m p o r t a n t

0

0

o

Ni

9

Li

9

Ni

or

Li

Figure 8. Active oxygen species of the solid solution series LixNi2_xO2. (a) Cubic type; ( b ) I n c o m p l e t e hexagonal type; (c) H e x a g o n a l type. role for the oxidative reaction. T h e f o r m a t i o n of active O- species occurs in t h e lattice oxygen and on t h e catalyst surface. It is p r o p o s e d t h a t an active O - species of the c a t a l y s t surface e x t r a c t s a h y d r o g e n from m e t h a n e , and accordingly t h e coupling of m e t h y l radicals take place.

254

CONCLUSIONS In this work, we present the results for LixNi2_xO2 ( 0 < x < l ) catalysts of the O C M reaction. In order to u n d e r s t a n d the surface catalytic active species, the characterization of the oxygen species of LixNi2_xO2 oxides is examined by the combined approach of XPS and UPS. From the X P S O ls and U P S spectral analysis, the main oxygen species of LixNi2_xO2 (x<0.6) with a cubic s t r u c t u r e can be assigned to 0 2 - , while those of Lil.0Nil.002 with a hexagonal s t r u c t u r e can be a t t r i b u t e d to catalytic active O- species for the O C M reaction. These results indicate t h a t the electronic structure and the catalytic activity for the O C M reaction strongly depends on crystal s t r u c t u r e of LixNi2_xO2 catalyst systems.

ACKNOWLEDGMENT The a u t h o r s acknowledge the staffs of the U V S O R Facility and IMS for their helpful advice and technical support.

References [1] G. E. Keller and M. M . Bhasin, J. Catal., r3 (~9s2) 9. [2] T. Ito J. X. Wang, C. H. Lin, and J. H. Lunsford, J. Am. Chem. Soc., 107 (1985) 5062. [3] M. H a t a n o and K. Otsuka, Shokubai, 29 (1987) 46. [4] A. M. Gaffney, C. A. Jones, J. J. Leonard, and J. A. Sofranko, J. Catal., 114 (1988) 422. [5] K. Mizushima, P. C. Jones, P. J. Wiseman, and J. B. Goodenough, Mater. Res. Bull., 15 (1980) 783. [6] J. R. Dahn, U. von Sacken, ics, 44 (1990) 87.

and C. A. Michal,

Solid States Ion-

[7] K. Seki in: Optical techniques to characterize polymer ed. H. Baessler (Elsevier, A m s t e r d a m , 1989) pp.l15-180.

system,

[8] C. L. Padro, W. E. Grasso, G. T. Baronetti, A. A. Cstro and O. A. Scelza, " N e w Developments in Selective Oxidation", 82 (1994) 411. [9] T. Miyazaki, T. Doi, M. Kato, T. Miyake, and I. M a t u u r a , manuscript submitted. [10] M. S. Hegde and M. Ayyoob, Surf. Sci. Lett., 173 (1986) L635.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

255

R e a c t i o n i n d u c e d spreading o f metal oxides: in situ R a m a n spectroscopic studies during o x i d a t i o n reactions Y. Cai +*, C.-B. Wang + and I. E. Wachs + Zettlemoyer Center for Surface Studies and Department of Chemical Engineering, Lehigh University, Bethlehem, PA 18015, USA 1. INTRODUCTION Metal oxides, particularly oxides of vanadium and molybdenum, are widely used as catalysts in numerous industrial applications (oxidation, oxidative dehydrogenation, dehydrogenation, olefin metathesis, olefin polymerization, selective catalytic reduction (SCR) and hydrodesulfurization (HDS)/hydrodenitrogenation (HDN)). In many of the applications, these metal oxides are supported on a high surface area metal oxide substrate, such a s A1203 or TiO2, to form an active surface metal oxide species (two-dimensional metal oxide overlayer). These supported metal oxide catalysts are typically prepared by impregnation of the corresponding soluble metal oxide salts, followed by drying and calcination at elevated temperatures (400-600~ There have been recent reports in the literature that an alternative route to the preparation of such supported metal oxide catalysts involves solid-state reactions, thermal spreading or spontaneous dispersion, of physical mixtures of the pure metal oxides at temperatures of 400-500~ [1-6]. However, no studies have examined the direct influence of the reaction environment on the spreading of metal oxides. In the present investigation, a new phenomenon of reaction induced spreading of crystalline MoO3 and V205 on oxide supports is observed at temperatures much lower than that required for thermal spreading via solid-state reactions, 200-250~ vs. 400-500~ 2. E X P E R I M E N T A L

The M o O 3 and V205 crystalline powders were obtained from Aldrich. The T i O 2 support (P-25; 55 m2/g) was purchased from Degussa and the SnO2 support (3.7 m2/g) from Aldrich. The physical mixtures of the binary oxides were prepared by combining an appropriate amount of MoO3 (or V205) with 5.0 g of TiO 2 (or SnO2) and 150 ml of pentane (Aldrich) in a beaker, and vibrating the mixtures for 15 minutes in an ultrasonic bath. After evaporation of pentane, the samples were dried for 16 hours at 100 ~ in air. No further treatments were performed.

*Current address: United Catalysts Inc., 1227 South 12th Street, Louisville, KY 40210. +Financial support of the Division of Basic Energy Sciences of the Department of Energy, grant DEFG02-93ER 14350, is gratefully acknowledged.

256 Thermal spreading was studied by treating the catalysts in an oven at a constant temperature for 1 h. Both loose powder catalysts and samples pressed into self-supporting wafers were investigated in the current studies. After the thermal treatments, Raman spectra of the catalysts were recorded in an ambient environment or in a controlled oxygen atmosphere. In the latter case, the samples were heated in oxygen (Linde Gas, ultra-high purity) in a dehydration cell for 45 minutes at 300~ and then cooled to about 50~ to obtain the Raman spectra. Additional details about the Raman spectrometer can be found elsewhere [5]. Reaction induced spreading during alcohol oxidation (methanol, ethanol and 2-butanol) was also investigated with both loose powder catalysts and samples pressed into self-supporting wafers. The self-supporting wafers were investigated with an in situ Raman spectrometer system that allowed monitoring of the changes during the reaction. A 100-200 mg sample disc was placed in the sample holder, which is mounted into a ceramic shaft rotating at 1500 rpm (see reference [7] for additional experimental details). The catalysts were initially heated in an O2/He =16/84 stream at 230~ for 30 minutes before recording the reference spectrum. A reaction gas mixture of CHaOH/OJHe - 4/16/80 was subsequently introduced into the in situ cell at a flow rate of 100 ml(STP)/min and the in situ Raman spectra were collected as a function of time. At the end of the reaction experiment, the methanol was removed from the gas stream and the catalyst was reoxidized in the O2/He stream. Alcohol oxidation over the loose powder catalysts were conducted in a fixed bed reactor at 230~ and atmospheric pressure. The details of the reactor system was previously described elsewhere [8]. The catalysts were pretreated in a flow of O2/He for 15 min prior to oxidation reaction. A reactant stream of C H a O H / O E / H e = 6/13/81 with a total flow rate of 100 ml/min was used for methanol oxidation. For ethanol and 2-butanol oxidation, a gaseous mixture of OR/He (13/81; ml/min) containing saturated ethanol or 2-butanol vapor at ambient temperature was introduced into the reactor. Analyses of reactants and products were carried out by an on-line Hewlett Packard 5890B GC. The spent catalysts were also characterized by Raman spectrometer. 3. RESULTS

3.1. Thermal spreading The ambient Raman spectra of a catalyst pellet containing a 4% MoO3/TiO 2 physical mixture after being exposed to different thermal treatments in a furnace are shown in Figure 1. The starting physical mixture only exhibited the Raman bands of crystalline MoO 3 (strong bands at about 814 and 990 cm -1) and the TiO 2 support (strong bands below 700 cm-1). After the one-hour 400 ~ thermal treatments, in both dry air and wet air, the Raman bands of crystalline M o O 3 predominate and only a trace of hydrated surface molybdenum oxide species is observed (Raman band at 956 cm-1). The Raman band of the surface molybdenum oxide species slightly increased as the temperature of the thermal treatments increased to 500~ Comparison of the initial Raman spectrum with the spectrum of the 500 ~ thermally treated catalyst pellet showed that most of the molybdenum oxide was still present as M o O 3 crystals and that only a small amount of surface molybdenum oxide species was present. In contrast to the thermal treatments with the catalyst pellet, almost complete dispersion of crystalline M o O 3 w a s observed after a 500~ treatment with the 4% MoOa/TiO 2 catalyst in the loose powder form. Thermal treatment of the loose powder at 400~ also showed a significant disperion of crystalline M o O 3 o n the titania support. This demonstrates that strong mass transfer limitations exist when the catalyst is in the form of a pellet, but not in the form of the loose powder. However, the design of the current in

257 situ Raman cell requires that the catalyst be in the form of a pellet and, consequently, the thermally treated pellet is the appropriate reference for the in situ Raman studies. Analogous thermal treatments, in dry as well as wet air, with a catalyst pellet and catalyst powder consisting of a 4% V2OJTiO 2 physical mixture revealed that no thermal dispersion of V205 occurred for these samples. Thus, thermal treatments of physical mixtures of MoO3/TiO2 and VzOs/TiO 2 at 400-500~ for 1 hour (1) result in the formation of surface molybdena species, (2) do not result in the formation of surface vanadia species and (3) the concentrations of surface species are significantly enhanced for catalysts in loose powder form compared to catalysts in pellet form.

Thermal Spreading of 4% MoO3/TiO z

814

Physical Mixture in Pellet Form (1 h)

A

I/) e=

990

s <

956

(d)

>i

w ~= C

=.,=.

(c)

o.,=.

(b) (a) 1100

1000

900

800

Raman Shift (cm "1)

Figure 1. Ambient Raman spectra of physical mixture of 4% MoO3/TiO2 catalyst pellet after different thermal treatments: (a) dry air (400~ 1 h), (b) wet air (400~ lh), (c) dry air (450~ 1 h) and (d) dry air (500~ 1 h).

258

4% MoOs+Ti02, Physical Mixture (230~

988 A

C: ::)

-969

s

,< w C q)

->

m

, i

814

q)

m

n,

x0.25

11 O0

1000

900

800

Raman Shift (crn "1)

Figure 2. In situ Raman spectra of physical mixture of 4% MoO3/TiO 2 catalyst pellet during methanol oxidation at 230~ 9(a) before methanol oxidation, (b) 20 min, (c) 1 h, (d) 3 h, (e) 5 h, (f) after reaction, oxidation of catalyst for 30 min and (g) after reaction, oxidation of catalyst for lh. 3.2. M e t h a n o l Oxidation

The in situ Raman spectra of a catalyst pellet consisting of a 4% MoOa/TiO2 physical mixture are shown during methanol oxidation at 230~ in Figure 2. The starting sample, Fig. 2a, only possesses the strong Raman bands of crystalline M o O 3 at 814 and 988 cm ~. Upon exposure to methanol oxidation conditions, Fig. 2b-e, the sharp Raman bands due to crystalline MoO 3 slowly diminish with reaction time and a new broad Raman band at 969 cm ~ is formed. The in situ Raman band at 969 cm ~ has previously been assigned to a surface molybdenum oxide coordinated methoxy, CHBO, species [9]. Upon switching to a flowing oxygen stream in the absence of methanol, Fig. 2f and g, the Raman band at 969 cm -~ shifted to about 990 cm -~ reflecting the decomposition of the surface methoxy-molybdate complex to a dehydrated surface molybdenum oxide species [9]. Simultaneously, there was also an increase in the Raman bands of crystalline MoO3 due to the oxidation of the partially reduced M o O 3 particles during the methanol oxidation reaction. Raising the reaction temperature to 300~ for about an hour, figure

259 not shown, resulted in the complete disappearance of the crystalline MoO3 Raman bands and only the appearance of the Raman bands associated with the surface molybdenum oxide species. Reoxidation of the sample at 300~ also resulted in the appearance of weak crystalline MoO3 Raman bands revealing that some residual reduced crystallites still remained and that higher temperature treatments are required to completely disperse the MoO3 crystals on the titania support. Essentially the same MoO3 dispersion behavior was observed during methanol oxidation with catalyst pellets consisting of 0.5-1% MoO3/SnO2 physical mixtures. Thus, the above in situ Raman studies clearly demonstrate that reaction induced spreading of crystalline MoO3 readily occurs over oxide supports during methanol oxidation at very mild temperatures, 230 ~C.

4% V2Os/TiO2,Physical Mixture

-791

(230~

-1022

A

C

E',

(e)

-E

<

>,1 rW

(c)

C (b) al

990

11 O0

1000

900

800

Raman Shift (cm "1)

Figure 3. In situ Raman spectra of physical mixture of 4% V2Os/TiO 2 catalyst pellet during methanol oxidation at 230~ 9(a) before reaction, (b) 30 min, (c) 1 h, (d) 3 h, (e) after reaction, oxidation of catalyst for 30 min, (f) after reaction, oxidation of catalyst for 1 h.

260

100

80

60 eCm~

r~

~,. 40

2O

I

0

50

'

I

100

'

Time (min)

I

150

'

200

Figure 4. Oxidation of alcohols over catalysts in loose powder form at 230 ~ as a function of reaction time: (a) methanol oxidation over 4% MoO3/YiO 2physical mixture, (b) methanol oxidation over 4% VzOs/TiO 2 physical mixture, (c) ethanol oxidation over 4% MoO3/TiO2 physical mixture and (d) 2-butanol oxidation over 4% MoO3/TiO 2 physical mixture. The in situ Raman spectra of a catalyst pellet consisting of a 4% V2Os/TiO2 physical mixture are shown during methanol oxidation at 230~ in Figure 3. The starting sample, Figure 3a, only exhibits the Raman bands of crystalline V205 at about 990 cm -1 and the titania support at about 790 cm 1. Exposure of the vanadia-titania catalyst to methanol oxidation at 230~ Figures 3b-d, completely removes the Raman bands of the V205 crystals and no new bands due to surface vanadia species are observed. The complete absence of any vanadia Raman bands suggests that the vanadia component of the catalyst was reduced (reduced vanadia gives rise to very weak Raman bands). Reoxidation of the 4% VzOs/TiO2 physical mixture catalyst pellet resulted in the appearance of a new Raman band at 1022 cm 1 associated with surface vanadia species [ 10,11 ] and the complete absence of crystalline V205 particles (no sharp Raman band at about 990 cm-1). Thus, the above in situ Raman studies clearly demonstrate that reaction induced spreading of crystalline V205 readily occurs over oxide supports during methanol oxidation at very mild temperatures, 230 ~ The catalytic behavior of the above physical mixtures, in loose powder form, were also investigated during methanol oxidation in a fixed-bed reactor as shown in Figure 4a and b. The methanol conversion over 4% MoO3/TiO2 continuously increased from about 8 to 16% during the first three hours of reaction. The corresponding methanol oxidation studies over the 4% VzOs/TiO 2 catalyst were more dramatic: the methanol conversion continuously increased from about 18 to 37% with reaction time during the first 110 minutes and then exhibited a sharp jump

261 to 100% methanol conversion at approximately 140 minutes. The jump in methanol conversion was accompanied by an increase in the temperature of the catalyst bed, approximately 244 ~C, due to the exothermic heat of reaction. Ambient Raman analysis of the spent catalysts revealed that both crystalline MoO 3 and V205 became almost completely dispersed during the methanol oxidation studies. Additional studies in an oxygen-free methanol environment further demonstrated that the dispersion of the crystalline oxides was not related to the presence of gas phase oxygen. Thus, the increase in methanol conversion as a function of time over MoO3/TiO 2 and V2OJTiO2 physical mixtures is directly related to the transformation of crystalline MoO3 and V205 into surface molybdena and vanadia species, respectively. 3.3 Ethanol and 2-butanol oxidation

The influence of oxidation reaction environments involving higher alcohols, ethanol and 2butanol, upon the catalytic behavior and dispersion of M o O 3 o n a TiO2 support was also examined. The catalytic behavior during ethanol and 2-butanol oxidation are shown in Figure 4c and d as a function of reaction time in a fixed-bed reactor containing the catalyst in loose powder form. The higher alcohols were more active than methanol and their conversions increased continuously with reaction time. In the case of ethanol oxidation, some blue Mo deposits were observed on the walls of the reactor exit due to the formation of volatile Mo species. Ambient Raman analysis of the spent catalysts revealed the presence of significant amounts of crystalline M o O 3 as well as the presence of some surface molybdena species. However, the concentrations of the surface molybdena species were much lower than that found after methanol oxidation and the surface molybdena species concentration was greater after ethanol oxidation than 2-butanol oxidation. Thus, these experiments reveal that the dispersion of crystalline MoO 3 particles on TiO2 supports during oxidation of alcohols follows the trend: methanol > ethanol > 2-butanol. 4. DISCUSSION The thermal spreading of metal oxides over oxide supports has been intensively investigated over the past decade and much is currently known about this process [ 1,5]. The driving force for the thermal spreading of metal oxides is related to the lower surface free energy of crystalline oxides such as V205 and MoO3 compared to crystalline oxide supports such as TiO2, S n O 2, A1203, etc. This process is analogous to the wetting of one solid by another induced by the forces of surface tension in order to lower the surface free energy of the system [2]. The low Tamman temperatures o f V 2 0 5 and M o O 3 (345 and 397.5~ respectively) are responsible for the efficient spreading of these metal oxides at temperatures of 400-500~ Furthermore, the spreading kinetics of the metal oxides are (1) dependent on the structure and morphology of the oxide support, (2) enhanced over well-developed crystal planes and (3) dependent on the specific gaseous environment (oxidizing vs. reducing or wet vs. dry) [5]. Under oxidizing conditions and elevated temperatures, the spreading of crystalline V205 and M o O 3 is initiated spontaneously at the metal oxide-support interface and subsequent migration occurs by surface diffusion of the metal oxides via vacancies or unoccupied sites in the two-dimensional metal oxide overlayer. Amorphous metal oxide phases are suggested as a transient form between the crystalline metal oxides and the two-dimensional metal oxide overlayers. Moisture enhances the surface diffusion of the metal oxides [1 ]. Under mildly reducing conditions, the spreading of crystalline metal oxides is significantly retarded due to the much higher Tamman temperatures of the

262 corresponding reduced crystalline metal oxides [1,5]. The present thermal treatment experiments in air revealed that extensive dispersion of MoO3 occurred at 400~ and essentially complete dispersion took place at 500~ for the loose powder physical mixture of 4% MoO3/TiO2. In contrast, very little dispersion was observed for comparable thermal treatments for the loose powder physical mixture of 4% VzOs/TiO 2. The observation that the kinetics of MoO 3 disperion are faster than the kinetics of V205 dispersion were also previously observed [12]. The lack of V205 dispersion by the thermal treatments is somewhat surprising and may be related to the structure and morphology of the titania support employed in the present investigation. The form of the physically mixed metal oxide was also found to significantly affect the dispersion kinetics due to the presence of significant mass transfer limitations in the catalyst pellet relative to the loose powder. The presence of mass transfer limitations in catalyst pellets or wafers typically employed for Raman and IR studies is welldocumented in the literature [ 13]. The present studies demonstrated that significant dispersion of M o O 3 o r V2 05 o n a titania support could not be achieved at temperatures of 500 ~ with the physically mixed oxides in the form of a catalyst pellet. Thus, dispersion of M o O 3 and V205 on oxide supports at much lower temperatures for physically mixed catalysts in the form of pellets can not be due to thermal spreading and must occur by another mechanism. The in situ Raman studies clearly demonstrate that spreading of MoO3 and V205 over different oxide supports in the form of catalyst pellets readily occurred during methanol oxidation at temperatures as low as 230~ Such low temperatures, which are below the Tamman temperatures of these oxides and the temperatures required for thermal spreading in the catalyst pellet (above 500~ implies that thermal spreading is not involved in the spreading mechanism taking place during methanol oxidation. This suggests that a strong interaction between the gas phase components and the crystalline metal oxide phases may be occurring. Formaldehyde is the major selective oxidation reaction product and is known to interact very weakly with metal oxides such as molybdates and vanadates, and adsorbed formaldehyde is readily displaced by the presence of moisture and methanol [14,15]. Moisture interacts strongly with molybdates [ 14,15] and vanadates [8], but the thermal spreading experiments did not result in significant dispersion of the crystalline metal oxides in the physically mixed catalyst pellet. The interaction of carbon dioxide with molybdates and vanadates is extremely weak and adsorption is usually not even observed at room temperature [16,17]. In contrast to these gaseous components, the interaction of methanol with molybdates and vanadates is very strong and is much stronger than moisture since adsorption of methanol can displace adsorbed moisture [14,15]. Furthermore, methanol oxidation over crystalline M o O 3 and V205 results in the deposition of molybdena and vanadia at the exit of the reactor, which possesses lower temperatures [ 18]. This observation suggests that methanol is able to strongly complex with Mo and V present in crystalline MoO3 and V205 to form volatile Mo(OCH3) n and V(OCH3)n complexes. The alkoxy complexes of vanadia and molybdena are well known and are liquids at room temperature possessing high vapor pressures. Thus, the low temperature dispersion of metal oxides over oxide supports during methanol oxidation is due to the formation of volatile metal-methoxy complexes that result in vapor phase transport of the oxides. The dispersion mechanism may also occur by surface diffusion of the metal-methoxy complex, but no such information is currently available. The absence of Mo and V deposits at the reactor exit during methanol oxidation suggests that either surface diffusion or readsorption of the volatile M-alkoxides is also taking place. In summary, a new phenomenon of reaction induced spreading of crystalline metal oxides on oxide supports is observed in the present investigation at temperatures much lower than that required for thermal spreading via

263 solid-state reactions, 200-250~ vs. 400-500~ Thermal spreading depends on the Tamman temperature of the crystalline metal oxide phases and reduced metal oxide phases possess very high Tamman temperatures which significantly retard migration [1,5]. However, essentially complete dispersion ofV205 on TiO2 was observed during methanol oxidation even though the in situ Raman spectra revealed that the vanadia was reduced under the reaction conditions (see Fig. 3). Essentially complete dispersion of M o O 3 o n TiO2 was also observed after treatment of the catalyst in an oxygen-free methanol environment. Thus, the oxidation state of the metal oxide does not appear to influence the kinetics of reaction induced spreading of crystalline metal oxides. Reaction induced spreading of MoO3 on oxide supports during oxidation of higher alcohols is significantly reduced relative to methanol oxidation (methanol > ethanol > 2-butanol). The reduced migration kinetics is most probably related to the stability and reactivity of the various alcohols. The rate determining step during the oxidation of alcohols to their corresponding aldehydes or ketones involves breaking the alpha C-H bond of the alkoxides (the carbon bonded to the alkoxy oxygen), and the stability of this bond is related to the number of additional carbon atoms coordinated to the alpha carbon: stability decreases with increasing number of carbon atoms coordinated to the alpha carbon [ 15,19]. Thus, the methoxy complex is more stable than the ethoxy complex, and the 2-propoxy complex is the least stable among these alkoxy complexes. The greater stability of the Mo-methoxy complex most probably is responsible for the greater volatility and spreading observed during methanol oxidation. The current findings that reaction induced spreading of metal oxides on oxide supports can occur during oxidation reactions at very low temperatures have important implications for commercial applications as well as fundamental studies. The oxidation of methanol to formaldehyde is industrially conducted with F e 2 ( M o O 4 ) 3 . M o O 3 catalysts that contain excess MOO3. The strong interaction between methanol and the MoO 3 component results in the stripping of the molybdena from the catalyst and its deposition as MoO3 crystalline needles at the bottom of the reactor where the temperatures are somewhat cooler. This volatilization phenomenon is responsible for catalyst deactivation (loss of activity and selectivity) and pressure build-up in such commercial reactors [20]. The opposite behavior is observed during methanol oxidation over MoO3/SiO 2 catalysts at 230~ The strong interaction of methanol with Mo and the weak interaction between surface molybdena species and the silica support results in agglomeration and crystallization of the surface molybdena species to beta-MoO3 particles during methanol oxidation [21 ]. A very important consequence of reaction induced spreading of metal oxides during alcohol oxidation is that the catalyst preparation method of many supported metal oxide systems is not critical since the same surface metal oxide species will form during reaction (especially methanol oxidation) [ 12,21 ]. Furthermore, the possibility that reaction induced spreading occurs during oxidation reactions over catalysts composed of physical mixtures needs to be very carefully investigated in such systems before other mechanisms are proposed to account for observed reactivity patterns [22]. 5. CONCLUSIONS A new phenomenon of reaction induced spreading of crystalline M o O 3 and V205 on oxide supports is observed during methanol oxidation at temperatures much lower than that required for thermal spreading via solid-state reactions, 230~ vs. 400-500~ The migration of the metal oxides appears to proceed by the formation of volatile M-(OCH3) . complexes and is not

264 influenced by the oxidation state of the metal oxide (both oxidized and reduced metal oxides are readily dispersed). The kinetics of reaction induced spreading of metal oxides during alcohol oxidation is much slower for higher alcohols because of the low stability of the corresponding M-alkoxides compared with the more stable M-methoxides. REFERENCES

1. 2. 3. 4. 5. 6.

H. Knoezinger and E. Taglauer, Catalysis, 10 (1993) 1. J. Haber, T. Machej and T. Czeppe, Surf. Sci., 151 (1985) 301. D. Honicke and J. Xu, J. Phys. Chem., 92 (1988) 4699. Y. Xie and T. Tang, adv. Catal., 37 (1990) 1. J. Haber, T. Machej, E. M. Serwicka and I. E. Wachs, Catal. Lett., 32 (1995) 101. F. Cavani, G. Centi, E. Foresti, F. Trifiro and G. Busca, J. Chem. Soc., Faraday Trans., 1, 84 (1988) 237. 7. J.-M. Jehng, H. Hu, X. Gao and I. E. Wachs, Catal. Today, 28 (1996) 335. 8. G. Deo and I. E. Wachs, J. Catal., 146 (1994) 323. 9. H. Hu and I. E. Wachs, J. Phys. Chem., 99 (1995) 10911. 10. M. A. Vuurman, I. E. Wachs and A. M. Hirt, J. Phys. Chem., 95 (1991) 9928. 11. G. Went, S. T. Oyama and A. T. Bell, J. Phys. Chem., 94 (1990) 4240. 12. T. Machej, J. Haber, A. M. Turek and I. E. Wachs, Appl. Catal., 70 (1991) 115. 13. Y. Cai and I. E. Wachs, to be published. 14. W.-H. Cheng, J. Catal., 158 (1996) 477. 15. W. Holstein and C. J. Machiels, J. Catal., 162 (1996) 118. 16. K. Segawa and W. K. Hall, J. Catal., 77 (1982) 221. 17. A. M. Turek, I. E. Wachs and E. DeCanio, J. Phys. Chem., 96 (1992) 5000. 18. G. Deo, H. Hu and I. E. Wachs, to be published. 19. W. E. Farneth, R. H. Staley and A. W. Sleight, J. Am. Chem. Soc., 108 (1986) 2327. 20. R. Pearce and W. R. Patterson, Catalysis and Chemical Processes (Wiley, New York, 1981) p. 263. 21. M. Banares, H. Hu and I. E. Wachs, J. Catal., 150 (1994) 407. 22. P. Ruiz and B. Delmon, Catal. Today, 3 (1988) 199.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

Temperature

Programmed

Desorption of Ethylene

265

and A c e t a l d e h y d e

on

U r a n i u m Oxides. E v i d e n c e o f F u r a n F o r m a t i o n from Ethylene. H. Madhavaram and H. Idriss Materials Chemistry, Department of Chemistry, The University of Auckland, Private Bag 92019, Auckland, New Zealand The reactions of acetaldehyde and ethylene have been investigated on the surfaces of UO 2 and UO3 by temperature programmed desorption (TPD). On UO 2 two molecules of acetaldehyde undergo reductive coupling to C4 olefins. This is due to the fluorite structure of UO 2, which can accommodate large numbers of excess oxygen, up to UO2.25. The vacant surface oxygen of UO 2 were titrated by N20 adsorption and were equal to 1.86 x 1019 molecules/g, representing an apparent surface area of vacant oxygen sites of 1.9 m2/g. On the other hand, ethylene-TPD on [3 UO3 indicated the desorption of acetaldehyde (490 K). In addition, an unexpected product was also observed. This product was identified as furan (C4H40, m/e 68, 39) which desorbed at ca. 550 K with a carbon selectivity of ca. 40 %. Furan formation from ethylene on UO 3 requiring the formation of one carbon-carbon bond and of one carbon-oxygen bond, is most likely accompanied by oxygen depletion from the UO 3 surfaces and subsequent reduction of U cations into lower oxidation states. The observation of furan from ethylene shows that one may obtain oxygenated products with a high carbon number from ethylene (a relatively abundant feed stock) via one single step. 1. INTRODUCTION Oxidation-Reduction reactions very often track the cation oxidation states of oxide materials [1 ]. Changing the oxidation state of a given cation is accompanied by structural change (such as from a rutile or anatase TiO 2 to corundum Ti203 or from orthorhombic V205 to rutile VO 2, i.e., changing of the coordination numbers of metal cations [2]). Another way of changing the oxidation states of cations is by creation of surface defects, where the surface looses its ordered structure [1, 2]. In the case of oxide materials several factors affect oxygen depletion (or in other words reduction of cations) the most important are the mass difference between the cation and the anion, the bond energy, and the formation of ordered or semi ordered cluster defects. While the mass difference is essentially important in the case of reduction via particle bombardment (see Sigmund theory [3]), bond energy and surface structure are most likely the dominant factors during

266 chemical reduction. The investigation of the effects of changing the structures and oxidation states of oxide materials is crucial to the understanding of their catalytic properties. The uranium oxides system is a good candidate for this investigation due to its presence in different stable and metastable structures - the main product of oxidations of uranium metal are UO2, U407, U308, U409, and UO 3 - as well as the presence of a wide range of oxidation states (from +2 to +6) [4]. The main reason for this wide range of oxidation states in U oxide (and the early actinides in general) is relativistic effects [5], which is simply a mass correction for the core electrons that lead to greater shielding of the higher orbitals, or in other words a decrease in the ionization potential and work function. Another important feature of some phases of U oxides is their possibility of accommodating large numbers of interstitial oxygen atoms in clusters, [2:2:2] clusters [6], without changing the crystal structure, such as UO 2 to UO2+x (x up to ca. 0.25); U cations in proximity of these clusters may have higher oxidation states than +4. There is also another implication to the ease of removing electrons form the outer shells, one can change the surface from one electronic state to another by reduction or by oxidation. Recently we observed that H2-reduction as + +4 well as Ar -sputtering of U308 resulted in surface cations exclusively in a U oxidation state [7]. This net change in oxidation state is unlike what one observes on early transition metal oxides such as Ti [8] and V [9] oxides. In addition, and most important, U oxides are known as good catalysts (or as components of catalysts) for serval industrial reactions such as olefins and paraffins ammoxidation [10-13], hydrocarbons dehydrogenation [14], and very recently for total oxidation of chlorinated compounds [15]. It has also been observed that U308 is active in C-C bond formation reactions such as the formation of isobutene from acetone [7]. Despite these technological importance fundamental understanding of the reactivity of U oxides surfaces towards organic reactions is lacking due to the very few amounts of work interested in this system, the most complicate oxide known [16]. This work is devoted to the understanding of the oxidation of CH2=CH 2 on UO 3 and the reduction of CH3CHO on UO 2 by temperature programmed desorption (TPD). Surface and bulk characteristics were investigated by X-ray Photoelectron Spectroscopy (XPS) and Xray Diffraction (XRD) as well as by N20 adsorption.

2. EXPERIMENTAL

TPD at atmospheric pressure was performed using a fixed-bed reactor interfaced to a high vacuum chamber equipped with a Spectra Vision quadrupole mass spectrometer (base pressure ca. 1 x 10.7 ton). The mass spectrometer is multiplexed with an IBM PC which is equipped with a programme (RGA for windows) that allows the monitoring of 12 masses simultaneously at a cycling rate of ca. 5 s. Catalysts were loaded into the reactor and heated under dry air (or hydrogen) for 2 hours at 800 K (or 10 hours, in the case of hydrogen reduction) prior to reaction. After cooling to room temperature (under H2 or air), the carrier gas was displaced by He (ultra pure) before adsorption of acetaldehyde or of ethylene. Acetaldehyde was placed in a saturator at room temperature. In order to obtain surface

267 saturation, dosing of acetaldehyde was performed while monitoring its m/e 29. A decrease in the signal (due to adsorption) followed by signal restoration is indicative of surface saturation (in ca. 2 minutes). In the case of ethylene dosing was obtained upon exposure for 5 minutes. The catalyst was then purged with He for ca. one hour at room temperature in order to remove traces of the reactant in the TPD line as well as weakly adsorbed molecules on the surface of the catalyst. The gas flow was introduced into the chamber through an -1 interface which consists of a leak valve differentially pumped to 10 torr during operating -6 conditions. A constant pressure of ca. 5 x 10 torr was maintained during all TPD runs. During the purging the m/e 29 (the highest m/e of acetaldehyde) or m/e 27 (for ethylene) were monitored and the TPD started when no change in this m/e signal was observed (after ca. one hour of purging time). The ramping rate during TPD was kept fixed at 15 K/min. The fragmentation pattems of each product were checked in order to identify unambiguously the reaction products by the method described previously [17]. This involved: (a) the separation of the desorption peaks into different domains of temperatures, (b) the analysis of the fragmentation pattern of each product separately, (c) starting from the most intense fragment for each product (m/e 29 for acetaldehyde, for example) and subtracting the corresponding amount of its fragmentation until the majority of the signals were accounted for. XPS was performed using a KRATOS XSAM-800 model with a base -9 pressure of ca. 10 torr. U (4f), O(ls), C(ls) and Ar(ls) (in the case of the Ar-ion sputtered samples) regions were scanned each run. Unreduced samples were loaded into the system without further treatment. Ex situ reduced samples (using H 2 at the same conditions as for TPD) were exposed to air (although, under oxygen free N 2 flow) for about 30 to 60 s, at room temperature, before introduction into the XPS chamber. Ar-ion sputtering was m /

performed using a direct beam KRATOS ion gun at a pressure of ca. 5 x 10 torr. Mg Ko~ radiation was used at 170 W. Collection of spectra were performed at a pass energy of 38 eV. Sample charging up to 5 eV occurred under X-ray irradiation. Binding energies were calibrated with respect to the signal of adventitious carbon (binding energy at 284.7 eV). No charging was observed with UO 3 samples. XRD data were collected using a Phillips 1130 generator, and a Phillips 1050 goniometer. XR radiation was achieved using a Cu tube (broad focus) (Kc~; X = 1.514 A) at 44 kV and 20 mA. N20 titration was performed in a pulse reactor. Pulses of N20 were introduced into the reactor at 480 K. A thermal Detector at the end of a Porapack Q column allowed the monitoring of N2 and N20 peaks. The absence of formation of N 2 accompanied by total restoration of N20 signal was indicative of total titration. This took about 4 pulses of l ml each (1 atm.) per 1 g of UO 3. UO 3 was prepared from a uranium nitrate solution by precipitation with NH 3 at pH 9. After filtration and drying at 373 K over night the powder was calcined at 673 K for 5 hours. XRD indicated a pure 13UO 3. Polysrystalline U30 8 (from BDH No. 26216) was used as received.

268 3. RESULTS 3.1. Surface and bulk characterisation of UO 2 and UO 3.

XRD spectra of UO3, U308 and H2-reduced UO 3 are presented in Figure 1. H2-reduction of UO 3 at 800 K for 10 hours resulted in transformation of the monoclinic phase of [3 UO3 into the orthorhombic fluorite structure of UO 2, although some orthorhombic ~ U308 is also present. Similar results were observed from H2-reduction of U308 [7], with complete transformation of o~U308 to UO 2, however. Table 1 Titration of oxygen vacancies by N20 adsorption on unreduced UO 3 and H2-reduced UO 3 (UO2). Reactor temperature 480 K, BET surface area of UO 3 = 33 m2/g, 1 g of catalyst, reduction temperature 773 K, 16 hours. Pulse number

N 2 (molecules)

N20 (molecules)

H2-reduced UO 3 (UO2) (1 g) 1 1.37 1019

1.32 1019

2

0.33 1019

2.4 1019

3

0.14 1019

2.6 1019

4

0.02 1019

2.75 1019

5

negligible

2.78 1019

total

1.86 10

19

Unreduced UO 3 (1 g) 1

no reaction

2.78 1019

XP spectra of UO2, U308, and UO 3 were analysed elsewhere [7]. A brief description is +

given here. Figure 2 (adapted from ref. 7) shows the XPS U4f region of UO 3, Ar -sputtered UO 3 and H 2 reduced U308. Three important points need indication. First, XPS U4f7/2 and U4f5/2 of UO 3 (spectra a) are higher in binding energy than those in spectra b and c. Second, spectrum a does not contain satellites while both spectra b and c contain satellites + at 386.5 and 397.5. Third, the XPS U4f peak positions of Ar -sputtered UO 3 as well as of H2-reduced U30 8 are those of UO 2 and UO2+ x respectively (see Table 1 in ref. 7), clearly indicating that one can shift the cation oxidation state from one position to the other (from +6 in UO 3 to +4 in UO2). This is unlike early transition metal oxides where, although they + are sensitive to H 2- or Ar - reduction, the resultant surfaces still contain considerable

I

XRD H2 - reduced U 0 3

'3O8

Ar+-sputtered UO, El

a,

u

El

L

U

d\x,j H,-reduced U,08

I

I

I

I

20

40

60

80

28 Figure 1. XRD of U03, U30s, and H2 - reduced U03 (mainly U02).

WOJ

400

395

390 385 380 375 Binding Energy (eV)

Figure 2. X P S of U03, Art -sputtered UO3 (U02) and H2 - reduced U30x (UOz,x)

270 amounts of stoichiometric phases. This unique characteristics of U oxides affect its chemical reactivity (see below), particularly with regard to oxidation-reduction reactions. The pulse method of N20 was investigated on UO 3 and H2-reduced UO 3 (UO2) (Table 1). This method is successful for titration of the surface area of Cu ~ and Ag o catalysts [ 18] and we wanted to try it to titrate oxygen vacancies instead of using oxygen in order to avoid formation of multilayers of dioxygen on the surface. N20 dissociated on UO 2 but not on UO 3 (Table 1).The dissociation reaction is activated, below ca. 425 K no dissociation occurred. A temperature of 480 K was observed as optimum were N20 dissociated non catalytically (catalytic decomposition occurred at ca. 525 K and above). From Table 1 one may estimate the total surface area of potential vacant sites which may abstract oxygen 2 19 from oxygenated compounds. Assuming that one m contains 1 x 10 atoms, N20 titration data indicated a surface of ca. 1.9 m2/g, or about 6 % of the total BET surface is composed of oxygen vacancies. 3.2. Acetaldehyde-TPD on UO 2.

Figure 3 and Table 2 present the desorption products during TPD after acetaldehyde adsorption on UO 2 (H2-reduced UO3). Table 2 Carbon yield and carbon selectivity of products formed during TPD after acetaldehyde adsorption at room temperature on UO 2 Product

Desorption Temperature (K) Acetaldehyde (m/e 29) 390 Propane (m/e 39) 610 Butadiene (m/e 54) 540 butene (m/e 56) 673 Ethanol (m/e 45) 415 CO 2 (m/e 44)

730

Carbon Yield (100%) 65.9 12.2 6.3 0.9 0.7 14.0

Carbon Selectivity (100 %) 35.8 18.5 2.6 2.0 41

Serval reactions occurred evidenced by a complex desorption products. First, acetaldehyde (m/e 29, 15, 44) desorbed at 390 K followed by traces of ethanol at 415 K (2 % of carbon selectivity, table 2). Three other products were observed. Butadiene and butene desorbed at 540 and 673 K respectively with a combined carbon selectivity of 21.1%. This reaction pathway follows a reductive coupling mechanism which has been observed previously on the surfaces of TiO 2 single crystal and powder [19-21]. The formation of C4 olefins is a clear example of the capacity of UO 2 surfaces to abstract large amounts of oxygen from surface carbonyls, via pinacolates [ 19], as follow

271

2 CH3CHO + 2 U

+4

- Vint.vac.

)

CH3CH=CHCH 3 + 2 U

+4+x - Oint.

Vint.vac.: interstitial oxygen vacancy, Oint.: interstitial oxygen. 3.3. E t h y l e n e - T P D on U O 3

Figure 4 and Table 3 show the desorption products during ethylene-TPD on ~ UO 3. Table 3 Carbon yield and carbon selectivity of products formed during TPD after ethylene adsorption at room temperature on UO 3 Product

Desorption Temperature (K) Ethylene (m/e 28, 27) 400-700 Acetaldehyde (m/e 29) 480 Furan (m/e 68, 39) 550 CO 2 (m/e 44)

above 800

H20 (m/e 18)

ca. 500

Carbon Yield (100%) 85.7 8.3 6.0

Carbon Selectivity (100%) 58 42

not calculated

In addition to ethylene desorption in a large temperature domain, acetaldehyde was clearly observed evidenced by its m/e 15, 29 and 44 (Table 3). The formation of acetaldehyde from ethylene indicates the facile removal of surface oxygen on UO 3 and shows its high reactivity towards oxidation of olefins. It is important to note that during TPD there is no regeneration of surface sites (in contrast to a steady state oxidation reaction with oxygen). This reaction requires a subsequent reduction of surface cations as follow

CH2=CH 2 + U

+6

-O

~

CH3CHO + U

+4

+ VO

(480K)

Vo: surface oxygen vacancy In addition, another important product was observed, furan (C4H40, m/e 68 and 39) at 550 K with comparable yield to acetaldehyde (42 % carbon selectivity). Thus, furan formation indicated that U surfaces are also active for C-C bond formation in their oxidised form, in addition of being an active C-O bond formation catalyst. The key route to this reaction is the formation of C4 olefin (most likely butadiene) which in its turn reacts with the surface oxygen to give furan as follow

CH2=CH 2 + CH2=CH 2 + O 1

CH2=CH-CH=CH 2

+ H20

cthylene/UO

acetaldehyde/U02

G

0 0-

-I--,

E4 k

0

v1 Q)

ct1i;inol x 40 0

pl-opane x 5

Q)

k k

0

0

5 300 400 500 600 700 800

300 400 500 600 700 800

273 CH2=CH-CH-CH 2 + 201

C4H40 (furan) + H20

(550 K)

Ol: lattice oxygen

Tow further points are worth mentioning. Firstly, XRD of the used ]3 UO3 (after TPD) indicated a mixed phase materials composed mainly of [3 UO3 and o~ U308. TPD of ethylene on this used UO3 (which have been regenerated under a dry air flow at 473 K for 90 minutes) showed a furan yield very similar to that on pure 13UO3 [22]. This result (which is under further investigation) may indicate that o~ U308 is also active towards this oxidative coupling reaction. It is important to mention that U308 contains substantial amounts of U +6 cations (together with U +4 or U +5 cations [7]). Secondly, in order to understand the reaction mechanism, TPD after butadiene adsorption at room temperature on ]3 UO3 was also investigated. Furan was clearly observed together with maleic anhydride [22]. This last point reinforces the above reaction mechanism.

4. CONCLUSIONS The oxidation of ethylene has been investigated on polycrystalline 13 UO 3 surfaces. Two oxygen containing products were observed: acetaldehyde, and furan. Furan desorption which requires a C-C bond formation, most likely is formed via dimerization of two adsorbed ethylene molecules followed by cyclization with available surface oxygen. Both the formation of acetaldehyde and furan from ethylene on UO 3 are clear examples of the ease of removing oxygen atoms from UO 3 surfaces. The reduction of acetaldehyde was also investigated on UO 2. Two molecules of acetaldehyde couple together to make a symmetric olefin: butene (which undergoes further dehydrogenation to butadiene). This is similar to what has been observed on TiO 2 and CeO 2 surfaces before [ 19-21, 23 ]. These complex chemical pathways indicate the richness of the U oxides system and open routes to further investigations.

References 1. M.A. Barteau, Chem. Rev., 96 (1996) 1413 and references therein. 2. V.E. Henrich and P.A. Cox, The Surface Science of Metal Oxides, 1994, Cambridge University Press, and references therein. 3. P. Sigmund, Sputtering by Ion Bombardment: Theoretical Concepts. Topics in Applied Physiscs, 47 (1981) 9. 4. C.A. Colmenars, Prog. Solid State Chem., 9 (1975) 139. 5. M. Pepper and B.E. Bursten, Chem. Rev., 91 (1991) 271. 6. R.J.D. Tilley, Defect Crystal Chemistry, Blakie, Glasgow and London, 1986.

274 7. H. Madhavaram, P. Buckanan and H. Idriss, J. Vac. Sci. Technol. A, 1997, in press. 8. H. Idriss and M.A. Barteau, Catal. Lett., 26 (1994) 123. 9. H. Poelman, L. Fiermans, J. Vennik and G. Dalmai, Surf. Sci., 275 (1992) 351. 10. K.M. Taylor, US Patent No. 3,670,006 (1972). 11. R.K. Grasselli and R.C. Miller, US patent No. 4010188 (1977) 12. R.K. Grasselli and D.D. Suresh, J. Catal. 25 (1972) 273. 13. D.D. Suresh, M.J. Seely, J.F. Brazdil and R.K. Grasselli, US Patent No. 4855275 (1989) 14. J.M. Hermann, J. Disdier, F.G. Freira and M.F. Portela, J. Chem. Soc. Farad. Trans., 91 (1995) 2343. 15. G.J. Hutchings, C.S. Heneghan, I.D. Hudson and S.H. Taylor, Nature, 384 (1996) 341. 16. F.A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 1972, Wiley, New York, third edition. 17. H. Idriss, K.S. Kim and M.A. Barteau, J. Catal., 139 (1993) 119. 18. M. Boudart and G. Djega-Mariadassou, Kinetics of Heterogeneous Catalytic Reactions, 1984, Princeton University Press. 19. H. Idriss, K.G. Pierce and M.A. Barteau, J. Am. Chem. Soc. 116 (1994) 3063. 20. H. Idriss, K.G. Pierce and M.A. Barteau, J. Am. Chem. Soc. 113 (1991) 715. 21. J.E. Rekoske and M.A. Barteau, Ind. Eng. Chem. Res., 34 (1995) 2931. 22. H. Madhavaram and H. Idriss, work in progress. 23. H. Idriss, C. Diagne, J.P. Hindermann, A. Kiennemann and M.A. Barteau, J. Catal., 155 (1995)219.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

275

Active Sites of Vanadium-Molybdenum-Containing Catalyst for Allyl Alcohol Oxidation: ESR S t u d y i n s i t u . O.V~Krviov. N~.uen Tien Tai. B.V.Rozentuiler N.N.Seme~ov Ln.stitute o! ( . h . m l c a ! Physics, Russian A c a d e m y of Sciences, ul.Kosygina 4, Moscow, 117884, Russia Allyl alc.r)hol oxidation into acrolein on the rhombic phase of molybu,:J~uJJ~ oxide in,-,u_ified with v a n a d i u m oxide has been studied by the Kne~nou ,Kl~leLic ,1 . , aim, , t,y: ESR of V . I + ions in sit'u. It was showl,, that active sites for ti~is reaction are V 'r ions situated in the bulk of the catalyst. ~w n e a r its s~rface, but n~:.,t at the surface. Fast diffusion r)f ,?iectro_ns and a more slower diffusion of oxygen ions b e t w e e n the s,~rface and t;he bulk occur d u r i n g the reaction. AI .

.

.

.

.

.

.

.

.

.

.

!1

u.

' 'roduction. A widespread opinion exists about the m e c h a n i s m s of oxidative catalytic reactions, that active sites for these reactions are surface coordinatively u n s a t u r a t e d transition metal ions. which can be r e d u c e d and reoxidized dllring the reaction. O u r studies of sevel-a] oxidation reactions [1,2] by .tt',S.R .in situ have shown t h a t the active sites can be situated not on!v at the surface, but in the bulk of the eata!y~t. We havc~ studied CO oxidation over the paraelectric phase of BaTiO3. , which contained about 0.01% Mn z+ ions as a n a t u r a l impurity. It t u r n e d .... J. ~u,o,L, 1 . _ ~. ~~ A~2+ ,_,u~., at 400-~t~0 K the active sites for CO oxidation are ..,,., ions. .Al(,r~g with adsorption and catalytic m e a s u r e m e n t s , we have proved using in situ ESR s ! ~ c t r a studies of Mn 2+ ions t h a t t h e y are situated in the bulk of BaTif) 3. It was found out that CO adsorption proceeds in exact accordance with the stoicbiometric equation, w h e r e C)~2- is a c,]rface c,x,,~en it,-' '~, Mr, 4+ + CO + O~'~- ---~ Mn ~+ 4- CO2,

(1)

i.e. adsorption of one CQ molecule results in a p p e a r a n c e of one Mn 2+ ion. On the contrary, 02 adsorption decreased the i n t e n s i t y of the ESR spo.otrum in accordance with the s t n i e h i n m e t r y

276 a~,~ ":,-,rf-'-! ~

~: u+_"__> 2Mn~ +-+ zuo.+"-' -, _ ,,

:.~,

tzl

i o. t:tvo Mn zl- inn~ d+sapp~.ar on a d s o r p t i o n of o n e nxyg~.n mnloe~iP. (In e x p o s u r e to t h e 2CO4-O z s t m e b i o m e t r i e m i x t u r e ~.he c a t a l y t i c r e a c t i o n is pr.,..-.eeeding t h r e u g h t h e r e d o x m e c h a n i s m f o l l o w s E q u a t i o n s (1) a n d {2). The 99 ..... - ~+.- o f m , , , . ~ , , . ~ o u r c o x i d a t i o n ( E q u a t i o n 2) is a l m o s t 2 ordcr~" of iiia,~iii|,tid~, N* u , , '~ J " t h a n t,ia~ ' ' oI~ t h e r e d u c t i o n ( E q u a t i o n l ) . In tht ~, 9 "........ " ' - "

.

.

established. cal,

.

. . . . . . .

.

The

.

.

lOW

reaction

proceeds

lu

predominantly

"

t2.oliCuIJtI

over

the

" "~t 'L.IUII

i8

oxidized

a J v s++. t.

~|'ho.~o invo.~tigntions wPro. e o n t i n ~ e d in m+r stud.les of a iiyi aieohoi ]r,+,:~ .'.~er,,l,~in m , e r t b e h e x a g o n a l pb_a~e of MoO a m o d i f i e d b y ! w ! % of Vz() 5 [2 ~]. T h e reacti~;n "" "~ '~

~,,,II~O11

:' .--~ \

+' '-'

.

Lot

ii2,k)

p r o c e e d s at 400-470 K w i t h i O0% s e l e c t i v i t y . P r a c t i c a l l y all t h e V r h,t~s rneam+t'o+d w i t h ~ S R i n si~.u w ~ r e in t h e b u J k of the, e a t a j v s t . VorT.r!atjoD of o~.e V 4.b i o n d u r i n g r e d u e t i o n of t h e catalyst. ]_s a c c o m p a n i e d b y t h e disapl.x~aranee of o n e allyl alcohol m o l e c u l e ,

. ,-, . + .'--:.i.,*~ . -,** ....§

+ [ ]~ + C : + 1 1 4 0

(+~,~

+ 119.O

it w a s s h o w n [3] t h a t V 4+ ions o b s e r v e d in E S R s o e e t r a a r e u n i f o r m l y c i i s t r i b u t e d o v e r t h e wiaole b u l k of h e x a g o n a i MoO~. (Ipym,~ito. offoet~ havo. bo~n o b s e r v e d on o x i d a t i o n of t h e c a t a l y s t b y ~,y.ygen. T h i s p r o , _ ' e s : ~ e.9....n+ be d e s c r i b e d b y f h e s t o i e h i o m e t r i e e q t m t i n n : q'~r4-t._t_

I 1

.l.. , q

., ~'tI1)',~-t--

/3-~

tV,~

AL 435 K Lhe t'aLe or Lhe s t a t i o n a r y c a t a l y t i c r e a c t i o n is proport"I O I l ~ l ' i:o t h e c,.Jneentration of v a n a d i u m ions a n d to t h e o x y g e n p r e s s u r e in th~ mixture

. . . .

k.;z.i '

J

,~ .... /

*~, i.e., at ~,~e- t e m p e r a ~++...4 at t h e t e m p e r a t u r e s ,,,+s,,~' 1.{~+1...... +l , ~ n.- a~-nn ~ - 6 0 0 ~r l;ui~'~ ot" at+ai,ioij~ti- s ~ t M y t i c oxidMiui,, t h ~ h e x a g o n M p h a s e of ~+ ***vO,,~ i~ [l'Hill

o I slt~n:~e.u

inl, o

the

more

s~ame, +

~

1

rhotnbie

phase 1

[6]. Thi~-

puper

is

277

devoted to the investigation of the mechanism of allyl alcohol oxidation into acrolein on the rhombic phase of MoO~ modified with V20~.

Experimental Samples of V-Mo-oxide catalysts have been prepared by mixing (NI-I4)6MoTOz4 and NH4VO.3, drying for 4 hours at 470 K, and calcination at 550-80~2 K Catalysts with 0.2, 0.5, 12, ,~.0, and 3.0 wt% of VzO5 were prepared. Their phase coml~)sition has been d e t e r m i n e d on the in~ t r u m e n t DRON-2 with Fe-K..~-emission. ESR s t ~ c t r a have been registered with the spectrometer EPR-V constructed in the Institute of Chemical Physics and equipped with an a t t a c h m e n t for the t e m p e r a t u r e regulation of the ampol!le with the s'ample directly :in the resonator of the spectrometer. Accuracy of the . is 0.2 ~ Calibration of the spectrometer has been t e m p e r a t u r e ~e"ulation, .~. Clone with the help of solutions of stable nitroxyl radicMs in benzene. g-values and IIFS constants were d e t e r m i n e d by comparison with a Mn24-/MgO standmxl. The absolute error in the d e t e r m i n a t i o n of the spin concentration by means of double integration is '_*50,%, the relative one is +2%. The error in g-value is 0.001. and that in ItFS constants is 0.1 Gs. The V4+ ions concentration was d e t e r m i n e d via double integration using the second parallel component of the ESR signal. A flow microcatalytic set-up has been combined with the ESR-spectrometer. Catalyst samples were placed into a flow- reactor which was at the same time an ampoule for the ESR studies. Gas mixtures He+O~, He+CsHsOII. and tIe§ have been prepared with the help of special 4- and 6-way valves. Gas analysis have been performed on the Carbowax column of the gas chromatograph, the length of the column was 1 m, and its t e m p e r a t u r e was 300 K. The catalytic reaction was studied at 380-540 K directly in the heated resonator of the ESR-spectrometer, at higher t e m p e r a t u r e s it was studied outside of the resonator. '

}

.

.

9

.

.

Results and discussion. Paramagnetic centres in V-Mo-oxide catalyst. Each of the observed ESR signals consists of 24 ItFS component~ due to interaction of an unpaired electron with the v a n a d i u m nucleus (s!V I=7/2. p=99.7o~). Their dependence on t h e , t e m p e r a t u r e of calcination of 2%V2Os/MoO 3 is shown in Fig.l The si.~nal tt (gx=gy=l.95r gz=l.908) corresponds to V4+ in the hexagonal phase of MoO~, the signals A, B, and C

278

~} 3"V.,(reI.uni tz )

Figure 1. De~mndence of the intencity of ESR signa]s observed in V-MoO~ catalyst on the heating temperature.

,K 593 623 673 723 873 q23 correspond to V4+ in the rhombic phase. It is seen that the rhombic phase only exists at temperatures higher than 600 K. This result was also confirmed by the XRD study. The signals A (gx=1.976, gy=1.974, gz=l.921) and B (gx=1.974, gy=l.970, g~.=1.928) are characterized by additional tIFS, which is typical for the interaction between unpaired electron of vanadyl and the nitrogen nucleus l'14N. I=l, p=99.63%). A calculation of the orbitals with an unt~aired electron [7] shows that the A signal corresl~mds to the V imide complex, where the NH group occupies the tx.,sition of one of the oxygen ligands around vanadyl. The B-signal a p p e a ~ at a higher t e m w r a t u r e (700 K). A calculation shows, that this signal corresponds to the interaction of the unpaired electron of vanadyl with a NO molecule. Both, the A and B complexes, are formed during preparation of the samples from arr lmonium molybdates and vanadates. The signal C (gx=l.971, gy=1.969, gz=1.872) s essentially different. It is fi~rmed at high temtmratures (800-900 K). Every component of its additional ftFS consists of one intensive line in the centre and 6 lines of equal intensity, separated by equal distances. The intensities of these lin,::s are in the ratio of 100:5.1. The appearance of this signal can be explained by the formation of non-stoichiometric phases in MoOa, the so called Magnelli phases [81, where the MoOs octahedrons are connected by planes and edges, but not by apexes. Such a non-stoichiometry contracts distances between the metal cations. Interaction of an tlnpaived electron of V4~ of the first octahedron with the molybdenum nucleus ('~,97Mo, I=5/2, p=25.18 %) of the adjacent octahedron connected by an edge with the fi.,~'t one gives additional HFS. The signals A, B, and C were observed at all concentrations of VzOa. Fig.2 shows the dependence of the n u m b e r of V 4+ ions and of the ESR line width on the VaO~ content. The maximal concentration of V 4+ ions

279

al:ter the catalyst reduction by allyl alcohol has been obsex"qed for 2% o[ V205. The line width increases monotonically with increasing percent of V~.O5 The V 4t ions are distributed uniformly and separately in the catalyst bulk with the increase of VzO 5 up to 2%. At higher V205 concentrations the signal intensity increases and the lines are broadened, because of strong mutual interaction of the V 4+ ions. At higher VzO 5 concentrations clusters of V 4+ ions are formed, and the n u m b e r of V 4+ ions observed in ESR diminishes. ,......

'~ ,3

<3

%

Figure 2. Dependence of V 4+ ions n u m b e r in V-MoO 3 catalyst and ESR line width on V20 5 content.

4O

o%

z

Io

Vo05 (wt %)

Reduction of rhombic MoO 3 modified with VzO 5 by allyl alcohol. The reduction kinetics of the rhombic MoOs with 2 wt% of VzO5 ha~ been studied at 387-437 K and at 1.2-2.8% of alcohol in the gas flow. Fig.3 shows kinetic curves of the V 4+ ions increase in the catalyst bulk arid of the decrease of acroiein formation t~te. These cur'ves are described by the first order equation. Acrolein and traces of water are appear in the gas phase. No significant decrease of the oxidation degree of the main MoOs matrix was observed, ~'~ rC4H~.O (10 ~6 motQc/9.=) . . . .

[V4+](10 TM ions/g.s)

~.~ Figure 3, Kinetic curves of the change of V 4+ ions ~$ u urnber and of aerolein formation rate; [C~H5OI{] = 7,1x l 0 I7 molec./cm s,

280 A q u a n t i t a t i v e comparison of these two processes, i.e., aclx~lein formation and V 4+ ions reduction shows t h a t 1.2-1.3 CstI,~O molecules are formed [~.r one V 44 ion produced. It indicates t~ssibly some reduction of Mo sites. The effective rate constants of rhombic phase reduction are shown in Table 1. The results of the s t u d y of hexagonal Table 1. Effective rate constants [C3IIsOII= 6.7xI017 m o l e c / c m 'z. MoO~ phase

T. K

of

V-MoOs-catalyst

387

397

417

437

Ris ......... i~e;iiid2);~2~.......(ii~ Hexagonal k~.r:c.,10-as -t 0.3

(I.6 0.4

1.1 1.0

1.9 2.4

reduction:

E a, k c a l / m o l e 11 16

V-MoO~ redl~ction for" similar conditions [3] are shown in the same tublu. 'Phes,a. dut~ indicttte a r a t h e r small influence of the MoOa structure, on the rate of its reduction. The formation of one acroiein molecule per one reduced V ion shows t h a t the charge compensation of V 5+ in MoO 3 takes place at the u c_am pJex formation. Some deviations from this expence of the v~75+-~' ratio can be explained by aerolein formation on o t h e r eentres. Thus, the catalyst reduction by allyl alcohol is described by the Equation (4). This process consists of at least of two steps" the surface one (7), which proceeds with pm-tieipation of the surface oxygen ion, and bulk step ('8), w h e n the charge carriers diffuse to the bulk and reduce the V 5+ ions: V.~5+O~" + C,~HsOH -~ Vs4+[ ]so + C.~tt40 + H20 V b5+O b- -~, Vsa+Os -

(7) (8)

The process (P,) represents diffusion of vacancies from the surface to the bulk und oxygen filling up of vacancies up at. the surface. We propose, as in [2], Thst the charge compensation of V 5+ in MoOs takes place at the expense of the formation of V5+O" ejmplex. Such complex should be paramagnetic, but it was not frecorde by ESR.. Special properties of r :vel~ not important for f u r t h e r discusision. The complex V4+[ ]0 is probably located in the i m m e d i a t e proximity to s t r u c t u r e s of the crystallographic shift, or Magnelly phases.

281

rteoxidation of the reduced rhombic phase of V-MoO~ by oxygen. Fig.4 shows the V "i+ decrease kinetics during reoxidation of the reduced rhombic V-MoO3 by oxygen. Reoxidation proceeds in accordance with b~quation (5), but the vacancy [ ]0 remains at the surface. The diffusion V~,+O - § V,+[ ]bo __> V~+[ ]o + Vs+Ob-

(9)

i~, slow. No products of this reaction were registered by the chromatographic analysis. A,l analysis of these results by the method of affine transformations showed that at temperatures higher than 520 K all the kinetic curves belong to the same family and can be described by a hype.rbolic equation No/[V 4+] - keff.t + 1, where N O is the initial n u m b e r of V 4+ ions in the sample. The values of " and the ,'elevant data for the hexagonal phase of V-MoOs [4] are h'eff. given in Table 2. These data show that the hexagonal phase reoxidation proceeds much faster. For the hexagonal phase similar values of kerr. are observed at tempera.tllres approximately by 100-1200 lower.

Table 2. Rate constants of V-MoO3. reoxidation;[O2]=2.7:• MoO~phase

T, K

426

460

486

528

554

570

lt~ molec/cm 3. 593

i~i~oml~ic........k~f~ii-O:~-s~ ........................................................................ 1.0 1.7 2.6 3.8 Hexagonal keffl0 -3 s -I 1.2 3.6 7.7

A comparison of Tables 1 and 2 shows that the kinetics of V 5+ ions reductions are identical in both, rhombic and hexagonal phases of MoO~, but the reoxidation of the rhombic phase is much slower.. This result probably explains the loss of activity of the catalyst induced by overheating during its preparation. In the hexagonal phase there are numerous cavities throughout the whole crystal.. Their diameter is about 0.5 nm, i.e. they can provide oxygen diffusion and, therefore, fast reoxidation. Such cavities are absent in the rhombic phase. From the other side:, no such channels are: necessary for the reduction of V-MoO.~, because they are too small for organic molecules. The reduction apparerently proceeds at the expence of electron migration from the. surface t

282

T ' 9 "

~

"

~90l;

~

; 03

~

~.

fO

r L.

2We "--~,--,..~ 0

t

t~

20

,,

I

40

0

~.

? 97~ ~ZoY ,.

9t (m;.) "~-

Q

'

e

4

GO

Figure 4. Kinetic curves of the change of V 4+ ions n u m b e r during catalyst reoxidation; [021 = 2 . 7.~10-~ molec./cm~.

,

i

* v2o 5(wt ~;)

Figure 5. Dependence of the reduction and reoxidation rate on V~O5 content in the rhombic phase of MOO3.

surface to the bulk. Thus, the increased catalytic activity of the hexagonal V-MoO s is explained by the easiness of catalyst reoxidation. Fig,5 shows rate constants dependence for reduction and reoxidation of r hombic V-MoO~ on the VzO 5 content. Both curves have a m a x i m u m at 2.0 wt% V~Os. These regularities correlate with changes of the ESR spectra (Fig,2)' the active sites are isolated v a n a d i u m ions. But with increase of V205 content higher than 2% pail~ed ions or clusters are formed. This leads to the loss of catalytic activity.

Stationary oxidation of allyl alcohol on the rhombic phase of V-MoO 3. Fig,6 shows the kinetic curves for the acrolein formation rate on the rhombic V-MoO s and for the change of V 4+ ions content in the

'z'c . 3

o t10*%oto~-~'g-- )

4~

~0

a,0

Oil

II

9

",~

-""--'4

0

Figure 6. Kinetics of the change of V 4+ ions n u m b e r in the rhombic phase of V-MoO 3 and the rate at 525 K; [C3HsOH] = 5.5x I 017 me]co./cm s. of acrolein formation

283

C3HsOH+O2 reaction mixture at 525 K. It is seen that the stationary reaction rate increases with increasing oxygen partial pressure in the mi• The acrolein formation rate is at any m o m e n t unambiguosly related to the n u m b e r of V4+ions in the catalyst. Such a regularity can be explained by the participation of the bulk V 4+ ions in the catalytic process. The kinetics of allyl alcohol oxidation on the rhombic phase of the V-MoO3 catalyst can be described by the equation

rc3mo =

N,:k ~d.[ C~H~O H] k ox.[02]

( 11 )

k~.~.a.lC.sHsCH] + ko~.[Oz] Table 3. Rate constants of the reduction and oxidation of the rhombic phase of V-MoO~ catalyst at 525 K; [C3H5OH] = 5.5• molec./cm :3. V20 5 content (wt. %)

0.5

1.0

2.0

3.0

kt..ed. ( 10 "21 cm~s "~)

from the experiment on reduction calculated by eq.(11)

0.25 0.29

0.40 0.52

0.58 0.96

0.90 0.90

k (10-2!.cm3s -I)

from the experiment on reoxidation calculated by eq.(ll)

2.0 2.3

3.8 3.0

5.6 4.6

5.4 4.1

o

.

~

.

.

This Equation does not differ from the usual Mars-Van Krevelen redox equation. The rate constants of the separate steps of oxidation and reduction from Equation (11) are listed in Table 3. They are compared in the same Table with the rate constants determined separately from the experiments on reduction and reoxidation. The coincidence b e t w e e n the calculated and experimental rate constants confirms the proposed redox mechanism of allyl alcohol oxidation over the rhombic phase ,:ff V-MoO 8 catalyst. Conclusion. A graphic scheme of allyl alcohol oxidation is represented in Fig.7. The ESR study of this reaction in situ has shown that the active sites

284

Alcohol O z- O z- O z- O~.- O z- 0 2Mo Mo Mo Mo V 5+ O z- O 2- O2-O 2- O ~- Oz--

-,

O 2- O z- 0 2- 0 2- 0 2- 0 2Mo Mo Mo Mo V 4+ 0 2-[ ] O 2- 0 2-O- 0 2-

-O

--~

Acrolein + w a t e r 0 2- [ ] O ~-- O~- O~- O~.Mo Mo Mo Mo V 4+ O 2- O 2- O ~- O ~- O- O ~-

-->

0 2- 0 2- 0 2- 0 2- O z- 0 2Mo Mo Mo Mo V4+ 0 2- 0 2- 0 2- 0 2-[ ] O-

Figure 7. The scheme of allyl alcohol interaction with the lattice of VMoOs catalyst. of the V-MoO 8 catalyst are V 4+ ions, which are situated mainly in the bulk of the catalyst, or near the surface. These ions supply electrons for the oxygen activation. Th~ second step of the redox catalytic process, reduction, proceeds via alcohol interaction with V 5+ ions at the surface. Thus, a continuous electron exchange between the catalyst surface and the bulk takes place during the reaction. This electron exchange is accompanied by oxygen ions (or vacancies) exchange in the case of hexagonal phase of V-MoO s. This catalyst contain~ r a t h e r broad cavities which allow fast oxygen diffusion. But the rhombic phase has a more dense lattice and oxygen diffusion in this phase is slower. This slow oxygen diffusion also explains the s o m e w h a t smaller catalytic activity of the rhombic phase of V-MoOs, References.

1. B.V.Rozentuller, K.N.Spiridonov, O.V.Krylov, Kinetika i Kataliz, 22 (1981) 797 2. M.A.Makarova, B.V.Rozentuller, O.V.Kry]ov, Kinetika i Kataliz, 28 (1987) 1143 3. M.A.Makarova, B.V.Rozentuller, O.V.Krylov, Kinetika i Kataliz, 28 (1987) 1395 4. M.A.Makarova, B.V.Rozentuller, O.V.Krylov. K i n e t i k a i Kataliz, 29 (1988) 872 5. M.A.Makarova, B.V.Rozentuller, O.V.Krylov, Kinetika i Kataliz, 20 (1988) 876 6. D.P.Shashkin, M.Yu.Kutyrev, P.A.Shiryaev. Proc. 5-th Intern. Symp. on Heterogeneous Catalysis (Varna, Bulgary) V.1 (1988) 177 7. D~Kivelson, S.-K~Lee, J.Chem.Phys, 41 (1964) 1806 8. L.C.Dufour, O.Bertrand. N~Floquet, Surface Sci., 147 (1984) 396

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

285

O x i d a t i v e d e h y d r o g e n a t i o n of e t h a n e over v a n a d i u m a n d n i o b i u m oxides s u p p o r t e d catalysts P. Ciambelli a, L. Lisi b, G. Ruoppolo ~ G. Russo o and J. C. Volta d. a Dipartimento di Ingegneria Chimica e Alimentare, Universith di Salerno, 84084 Fisciano (SA), Italy. Istituto di Ricerche sulla Combustione, CNR, Napoli, Italy. oDipartimento di Ingegneria Chimica, Universith "Federico II", Napoli, Italy. d Institut de Recherches sur la Catalyse, CNRS, Villeurbanne, France.

Ethane oxidehydrogenation has been investigated over niobium and vanadium oxides supported on high surface area TiO2. The vanadia-titania catalysts are very active but with low selectivity due to their high reducibility. The selectivity to ethylene is enhanced by the presence of niobium. By changing the order of addition of vanadia and niobia to the support, catalysts with slightly different redox and acid properties are obtained. At low vanadium loading, supporting the two oxides at the same time results in the best catalytic performances, while at high loading a two steps impregnation gives the best results.

1. I N T R O D U C T I O N Oxidative dehydrogenation (ODH) of light paraffins could be an alternative route to produce high purity olefins compared to conventional methods such as dehydrogenation and cracking [1]. Paraffin dehydrogenation in the presence of oxygen is thermodynamically favoured due to water formation, however the selectivity to olefins is generally poor because of the low reactivity of the alkane compared to the formed olefin [2]. In the last decade attention has been paid to ethane ODH due to the interest in transforming an abundant component of natural gas into a more valuable product. It has been shown that catalysts based on Li and Mg mixed oxides, active in the reaction of methane coupling, are also active in ethane ODH [3-4]. A similar behavior has been also found for various rare earth oxides [5]. Catalysts based on transition metal oxides are the principal materials investigated in alkane ODH. Ethane is oxidatively dehydrogenated to ethylene with high conversion and selectivity over bulk metal oxide catalysts containing

286 Mo, V and Nb. It was proposed that the role of Nb20~ is to enhance the intrinsic activity of Mo and V and to improve the selectivity by inhibiting the oxidation of ethane to carbon oxides [6, 7]. The catalytic properties of pure V20~ are inadequate [8], whereas crystalline Nb20~ is highly selective in propane ODH [9]. Supporting niobia on a-A1203 does not result in better performance, whereas niobia supported vanadia is more active than pure niobia, maintaining the high selectivity [ 10]. Moreover, the effect of supporting vanadium oxide on either silica or alumina on the catalytic properties of V20~ in alkane ODH [2, 11-13] has been investigated. VgOdTiO9 catalysts have been widely used in selective oxidation of hydrocarbons. A new application of ethane ODH has been reported for a combination of vanadium and molybdenum as phosphates on TiO9 [14]. In this work, the activity and selectivity of catalysts based on niobium and vanadium oxides supported on high surface area anatase TiO2 in ethane ODH have been investigated. Specifically, the influence of the cooperation of vanadium and niobium oxides supported phases as components inducing respectively redox and acid properties, together with the effect of the preparation conditions on the catalytic performances have been studied. 2. E X P E R I M E N T A L

2.1. Catalyst preparation The support was pure anatase TiO2 (surface area 125 m 2 g-l) supplied by hoxide Specialties. Binary and ternary catalysts, containing respectively only supported vanadium or niobium oxide or both oxides, were prepared by wet impregnation with vanadium metavanadate (BDH Laboratory Supplies) and niobium ammonia complex (Companhia Brasileira de Metalurgia e Minera~fio) aqueous solutions. After impregnation the materials were dried at 110~ and calcined at 550~ in flowing air. The ternary catalysts were prepared either by using both precursor salts of vanadium and niobium oxides at the same time or by changing the order of addition of the two components to TiO2. Before the addition of the second component, the catalyst containing vanadium or niobium oxide was dried and calcined. These operations were repeated after the second impregnation. The supported catalysts will be denoted as xVyNb/Ti, x being the V20~ and y the Nb205 nominal content (weight percentage). Furthermore, the ternary catalysts prepared in two steps will be denoted as (I) ff vanadium oxide is deposited on the support in the first step and as (II) if niobium oxide is added in the first step. Binary and ternary catalysts with vanadium content of 1 and 6 wt% V20~ and binary and ternary catalyst with niobium content of 6 wt% Nb~O~ were prepared.

2.2. Physico-Chemical Characterization XRD analysis was performed with a PW 1710 Philips diffractometer. The BET surface areas were determined by N2 adsorption at 77K with a Carlo Erba 1900 Sorptomatic.

287 Temperature Programmed Desorption (TPD) of NH3 and Temperature Programmed Reduction (TPR) with H2 were performed using a Micromeritics TPD/TPR 2900 analyzer equipped with a TCD and coupled with a Hiden HPR 20 mass spectrometer. The samples were preheated in flowing air at 550~ for 2 h. In TPR analyses a 5% HdAr mixture (25 cm~/min) was used to reduce the sample by heating 10~ up to 670~ In NH3 TPD analyses, after saturating the sample with a 2% NH3/He mixture, ammonia was desorbed by heating 10~ up to 650~ in flowing He (25 cm~/min). In situ Raman spectra were recorded with a DILOR OMARS 89 spectrophotometer equipped with an intensified photodiode array detector. The emission line at 514.5 nm from Ar § ion laser (Spectra Physics, Model 164) was used for excitation. The power of the incident beam on the sample was 36 mW. Before the acquisition of the spectrum the sample was heated up to 400~ (2~ and kept at this temperature for 12h to obtain a complete dehydration of the surface. After cooling down to 300~ the spectra were recorded. The same procedure was used to acquire Raman signals of the pure TiO9 support. The final spectra of the catalysts were obtained by subtracting the TiO~ contribution.

2.3. Catalytic activity tests Catalytic activity tests were carried out with a fixed bed quartz microreactor at atmospheric pressure. The catalyst (particle size = 300-400~m) was placed on the porous septum of the reactor. In order to limit the occurrence of homogeneous reactions a-A120~ particles were loaded before the catalytic bed and the reactor diameter was reduced after the catalytic bed. The temperature was monitored by a type K thermocouple located in the catalyst bed. The reactor outlet gas was analyzed with a H a r t m a n n & Braun URAS 10 E electrochemical/IR continuos photometer for 02, CO and CO2 and with a Hewlett Packard 5890A gaschromatograph equipped with a flame ionization detector for C2H6 and C2Ha. Water produced by the reaction was kept by a silica gel trap in order to avoid condensation in the cold parts of the apparatus. The contact time ranged from 0.001 to 1 g s cm ~ depending on the different activity of the catalysts. The reaction temperature was 550~ The feed composition was 2% 02 and 4% C2H6 in helium. Carbon balance was always closed to within + 2%.

3. R E S U L T S AND D I S C U S S I O N By assuming a monolayer capacity of 15.3 wt% V20~ [14] and of 17.5 wt% Nb205 [15] with reference to the surface area of the support and considering that the surface coverage in the ternary catalysts is the sum of V20~ and Nb20~ single coverages, it can be concluded that the theoretical monolayer capacity was never exceeded even in the 6V6Nb/Ti catalyst. However, traces of T or TT Nb205 phases [9] were detected in this sample by XRD suggesting that some aggregation occurs when the coverage approaches the monolayer. No peaks due to TiO2 rutile phase were observed in the XRD spectra of the catalysts indicating that the calcination temperature is low enough with respect to the phase transition from anatase.

288 For all the catalysts a quite negligible loss of surface area with respect to TiO2 was observed, the values ranging between 111 and 105 m2g1.

3.1. TPR e x p e r i m e n t s The results of TPR experiments are reported in Table 1. For all the samples no significant effects of H2 reduction were observed below 330~ Similar results were obtained by Topsoe et al. [ 16] on vanadia-titania catalysts. The TiO2 support and the 6Nb/Ti catalyst undergo a very poor reduction compared to that occurring when vanadium is present in the catalyst. The TiO2 support shows a tailed peak with maximum at 554~ while the addition of niobium oxide results in the appearance of a new peak at 609~ For both V/Ti catalysts the onset and the peak temperature of reduction are lower than those of pure TiO2. The reduction of the support is still clearly detected as a distinct peak when the V20~ content is 1 wt%, whereas it is completely hidden by the more intense reduction peak of vanadium oxide at the larger V20~ content. The value of V/I-I2 ratio of about 1 suggests the reduction of vanadium from a +5 to a +3 average oxidation state probably occurring in a single step. Values of V/H2 approaching 1 were also found by Went et al. [17] for vanadia-titania catalysts and by Blasco et al. [18] for V20~ on different supports. The presence of niobium in the ternary catalysts results in shifting the peak temperatures of H2 uptake to higher values, if compared to those of the corresponding Vfl~ catalysts. The temperature shift is greater at low than at high vanadium loading, indicating a different effect of niobium on the redox properties of the ternary catalysts. Moreover, the shift is smaller for the samples impregnated in two steps with respect to the single step impregnation. The values of V/H2 are somewhat affected by the presence of niobium, but it is d~f-ficult to quantify the contribution due to the reduction of the support, especially at low vanadium loading. 3.2. TPD experiments. No ammonia oxidation products were detected during the TPD experiments, except for a very small amount of N2 observed at T>500~ with the vanadium containing catalysts, likely due to reaction with catalyst surface oxygen. This is in agreement with the easy reducibility of vanadia-titania samples found in TPR experiments. Lietti and Forzatti [19] observed ammonia oxidation at high temperature on V20~/TiO2 but on pure TiO2. In Figure 1 NH~ TPD profiles are reported. The tailed TPD peak with maximum at 190~ indicates that ammonia is desorbed from the TiO2 surface in a wide range of temperatures (100-650~ due to the presence of different adsorbed species as also found in FT-IR studies [20]. The addition of V20~ or Nb205 to TiO2 results in the modification of the ammonia desorption profile of the support, strongly depending on the metal oxide. Niobium oxide causes a slight decrease of NH~ adsorption with respect to titania only in the range 180-350~ At lower and higher temperature the two profiles are totally superimposed suggesting that the nature of acid sites on TiO2 and 6Nb/Ti catalyst is the same, probably of Lewis type as proposed by Pittman and Bell [21].

289

r,.) o 1.5-

~

o ~ l . 0

oo

-1.5

, , 6V6Nb/Ti [\/ . 6V/Ti

1V6Nb/Ti

o

,,/1V/Ti

o

9

-

1

-

1.0

~"

oo

o

o

=9 0.5 -

-

a,

0.5

o

=

o

o

r~

o'3

~ 0 . 0

-

Z

0

I

I

I

200

400

600

0

I

I

[

200

400

600

-0.0

Z

Temperature (~

Temperature (~

Figure 1. NHa TPD of TiO2, 1V/Ti, 1V6Nb/Ti, 6Nb/Ti, 6V/Ti and 6V6Nb/Ti.

Table 1 H2 u p t a k e and peak t e m p e r a t u r e in TPR experiments and NHs desorbed in TPD experiments. Tmax V/H2 ratio NHa desorbed Catalyst H2 uptake (10 .4 mol -g-') (~ (~mol .m 2) , TiO2 0.6 554 3.5 1V/Ti

1.1

496

1.0

3.3

1V6Nb/Ti

1.3

539

0.8

3.7

1V6Nb/Ti (I)

1.6

521

0.7

3.7

1V6Nb/Ti (II)

1.3

523

0.8

3.8

6V/Ti

6.1

516

1.1

2.7

6V6Nb/Ti

6.5

539

1.0

3.3

6V6Nb/Ti (I)

5.6

532

1.2

3.5

6V6Nb/Ti (II)

5.6

526

1.2

3.8

6Nb/Ti

1.1

609

3.4

Differently, the change of the NHa TPD profile of Ti02 due to the presence of v a n a d i u m suggests t h a t medium and high temperature sites of the support are preferentially covered with lower acid strength sites. This is indicated by the appearence of a sharp peak with m a x i m u m in the range 110-155~ shown by all

290 the vanadium containing samples and by the absence of any signal of NH3 at T > 500~ for the high vanadium content catalysts. It was reported [16] that at low loading vanadia interacts preferentially with the most basic hych'oxyl groups present on the titania surface. At high vanadium loading, most of Ti-OH groups are replaced by new Bronsted acid sites which give rise to a NH~ band whose intensity increases with the vanadium content. This result is in agreement with the substitution of strong Lewis acid sites of TiO2 with weaker Bronsted acid sites due to V-OH groups observed in the vanadium containing samples. The TPD profiles of ternary catalysts are also different from those of binary catalysts, showing that the presence of niobium results in a different distribution of acid sites, especially at low vanadium loading. However, both xV6Nb/Ti (I) and (II) give rise to a lower temperature signal. It is noteworthing that the presence of niobium oxide enhances the acidity of the bynary catalysts, more strongly at high vanadium loading. The amount of desorbed NH3 from the ternary catalysts is slightly affected by the preparation method, especially at high vanadium loading (Table 1).

3.3. Laser-Raman spectroscopy. In Figure 2 Laser-Raman spectra of 6Nb, 1V/Ti, 6V/Ti, 1V6Nb/Ti and 6V6Nb/Ti catalysts are reported. The spectrum of 6Nb/Ti shows a narrow peak at ca. 990 cm 1 which can be attributed to the double bond Nb=O of both tetrahedral and octahedral NbOx species [21]. A broad band, with lower intensity and centred at ca. 920 cm -1, probably due to Nb-O-Nb bridges in octahedrally coordinated species, is also present [21]. 6V/Ti catalyst shows a narrow peak at 1030 cm 1 due to the double bond V=O and a broad band at 915-920 cm 1 attributed to V-O-V bridges in polycondensed species [17]. The band at 915-920 cm -1 is absent in the spectrum of 1V/Ti sample, where tetrahedral isolated species prevail. The spectrum of 6V6Nb/Ti catalyst shows no remarkable difference from that of 6V/Ti; moreover there is no evidence of a contribution at 990 cm 1 related to Nb=O bonds. However, in the spectrum of 1V6Nb/Ti catalyst a broad signal at ca. 990 cm -1 indicates the presence of Nb=O bond. Its intensity is negligible if compared to the corresponding signal in the spectrum of 6Nb/Ti. The disappearence of the Nb=O signal in the ternary catalysts could suggest an interaction between the two oxide phases. The formation of V-O-Nb-O-V bridges can be suggested, or the grafting of vanadium onto niobium oxide phase can be hypothesized.

3.4 Catalytic activity tests The possible occurrence of homogeneous reactions was tested by performing experiments in the absence of catalyst under the same reaction conditions of the catalytic tests. No ethane conversion was observed up to 700~ In the activity tests the oxygen conversion was kept always <100% and the temperature increase, due to the exothermal reactions, was negligible. All the catalysts produce C2H4, CO and CO2.

291

6NbtTi

1VITi 6V/Ti - - _

.

_

~

~ 12~)o

11()0

1000

6V6Nb/Ti

901)

80'0

Wavenumber (cm -1) Figure 2. Laser-Raman spectra of 6Nb, 1V/Ti, 6V/Ti, 1V6Nb/Ti and 6V6Nb/Ti catalysts.

In Table 2 the results of catalytic tests are reported for ethane conversions of 10-15%. It is noteworthing that the support is able to activate the reaction with a quite good selectivity to ethylene (64%). The addition of niobium oxide improves the selectivity slightly depressing the original activity of TiO2. On the other hand, the addition of vanadium oxide strongly enhances the activity of TiO2 in the whole range of compositions investigated as shown by the much lower contact time required to obtain the same conversion levels. At the higher vanadium loading a decreased selectivity to ethylene (37%) is observed. In the ternary catalysts the presence of niobium increases the selectivity to ethylene only at low vanadium content, maintaining a comparable activity with respect to 1V/Ti. At higher vanadium content both activity and selectivity to ethylene are only slightly affected by the presence of niobium. In Figure 3 selectivity to ethylene is reported as a function of ethane conversion over TiO2, 6Nb/Ti, xV/Ti and xV6Nb/Ti catalysts. The most selective catalyst is 6Nb/Ti; the high selectivity to C2H4 exhibited by TiO2 is lowered by the increasing addition of vanadia. Moreover, the addition of niobium oxide to 1V/Ti has a promoting effect on the selectivity to ethylene, whereas the addition to 6V/Ti does not result in any marked change of selectivity.

292 Table 2 Results of catalytic activity tests of ethane ODH (T-550_~ Catalyst

W/F

X02

(g s cm -3) (%)

Xc2H6

SC2H4

(%)

(%) .

.

.

.

.

.

.

.

Sco

8c02

(%)

(%)

.

.

.

.

.

TiO2

0.34

24

11

64

33

3

6Nb/Ti

0.60

20

12

73

25

8

1V/Ti

0.22

45

13

44

46

10

1V6Nb/Ti

0.14

24

10

56

41

3

1V6Nb/~ (I)

0.30

33

12

51

44

5

1V6Nb/Ti (II)

0.20

35

12

53

45

2

6V/Ti

0.01

35

12

37

60

3

6V6Nb/Ti

0.01

42

12

35

61

4

6V6Nb/Ti (I)

0.01

41

12

37

59

4

6V6Nb/Ti ~II~

0.01

42

14

39

.... 59

2

It must be remembered that the results of catalyst characterization give evidence for a different effect of niobium depending on the vanadium content. In fact (Table 1) at low vanadium content the redox properties are much more affected by the presence of niobium with respect to the catalysts with high vanadium content. In contrast, a lower effect on the acidic properties has been found at low vanadium with respect to high vanadium content. The Raman spectrum of 1V6Nb/Ti seems to indicate that Nb interacts with V which is dispersed as isolated tetrahedral species at low loading. In contrast, negligible effects on the catalytic performance are caused by the addition of niobium to 6V/Ti, in agreement with the smaller effect observed with respect to the redox properties. These results should indicate that the effect of niobium on the catalytic properties of vanadia-titania catalysts depends on the nature of the VOx surface species. In Figure 4 the effect of the preparation method on the selectivity to ethylene is shown. The selectivity of 1V6Nb/Ti is only slightly higher than that of the corresponding samples prepared in two steps, confirming the results of Table 2. On the contrary, at high vanadium loading the catalysts of type (I) and (II) are more selective than 6V6Nb/Ti, but less selective than the 1V6Nb/Ti series. This supports the hypothesis that the nature of vanadium oxide species is a key factor to provide high selectivity to ethylene in ethane ODH.

293

Figure 3. C2H4 selectivity as a function of C2H6 conversion at 550~ 6Nb/Ti (0), 1V/Ti (A), 6V/Ti (El), 1V6Nb/Ti (A) and 6V6Nb/Ti (u).

on TiO2 (o),

Figure 4. C2H4 selectivity as a function of C2H~ conversion at 550~ on 1V6Nb/Ti (o), 1V6Nb/Ti (I) (E]),IV6Nb/Ti (II) (A), 6V6Nb/Ti (o), 6V6Nb/Ti (I) (u) and 6V6Nb/Ti (II) (A).

294 3. CONCLUSIONS Catalytic performances of VOx/TiO2 systems in ethane ODH are improved by the addition of niobium. When TiO2 is coimpregnated by vanadium and niobium oxides, the presence of niobium enhances the selectivity to ethylene at low vanadium content, whereas it slightly depresses the activity without enhancing the selectivity at high vanadium content. This should be due to the effect of niobium on vanadium reducibility, especially affected at low vanadium content. The interaction between the two supported oxides can be modified by changing the preparation technique. By reversing the order of addition of vanadium and niobium at high V20~ loading, catalysts having better performances can be obtained, while at low loading no effect has been found. This difference is likely due to the different nature of vanadium oxide supported species. The catalytic performances are also changed as a result of catalyst acidity modifications induced by the presence of niobium. REFERENCES

1. 2. 3 4. 5 6. 7. 8 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21.

F. Cavani and F. Trifir6, Catal. Today., 24 (1995) 307. E.A. Mamedov and V. Cort6s Corber~n, Appl. Catal., 127 (1995) 1. E. Morales and J.H. Lunsford, J. Catal., 118 (1989) 255. R. Burch and S.C. Tsang, Appl. Catal., 65 (1990) 259. E.M. Kennedy and N.W. Cant, Appl. Catal., 75 (1991) 321. K. Tanabe and S. Okazaki, Appl. Catal. A: General, 133 (1995) 191. R. Butch and R. Swarnakar, Appl. Catal., 70 (1991) 129. J. Le Bars, J.C. V6drine, A. Auroux, S. Trautmann and M. Baerns, Appl. Catal. A: General, 88 (1992) 179. R.H.H. Smits, K.Seshan and J.R.H. Ross, Studies Surf. Sci. Catal., 72 (1992) 221. R.H.H. Smits, K. Seshan, H. Leemreize and J.R.H. Ross, Catal. Today, 16 (1993) 513. J. Le Bars, A. Auroux, M. Forissier and J.C. V6drine, J. Catal., 162 (1996) 250. A. Erdohely and F. Solymosi, J. Catal., 123 (1990) 31. J.G. Eon, R. Olier and J.C. Volta, J. Catal., 145 (1994) 318. M.Roy, M. Gubelmann-Bonneau, H. Ponceblanc and J.C. Volta, Catal. Lett., 42 (1996) 93. J.-M. Jehng and I.E Wachs, J. Phys. Chem., 95 (1991) 7373. N.-Y. Topsoe, H. Topsoe and J.A. Dumesic, J. Catal., 151 (1995) 226. G.T. Went, L.-J. Leu and A.T. Bell, J. Catal., 134 (1992) 479. T. Blasco, J.M. L6pez-Nieto, A. Dejoz and M.I. V~squez, J. Catal., 157 (1995) 271. L. Lietti and P. Forzatti, J. Catal., 147 (1994) 241. G. Ramis, G. Busca, V. Lorenzelli and P. Forzatti, Appl. Catal., 64 (1990) 243. R.M. Pittman and A.T. Bell, Catal. Lett., 24 (1994) 1.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

P a r t i a l O x i d a t i o n of E t h a n e over M o n o l a y e r s of V a n a d i u m Effect of t h e S u p p o r t a n d S u r f a c e C o v e r a g e .

295

Oxide.

Miguel A. Bafiares a, Xingtao Gaob, Jos~ L. G. Fierro a, and Israel E. Wachs b aInstituto de Cat~lisis y Petroleoquimica, CSIC. Campus Cantoblanco, E-28049, Spain bZettlemoyer Center for Surface Studies, Departments of Chemistry and Chemical Engineering, Lehigh University, Bethlehem, PA-18015, USA

The nature of supported oxides and of the support plays a critical role in the partial oxidation of hydrocarbons since the support is not only providing a high surface area, but also dispersing the oxide. The interaction between the metal oxide overlayer and the underlying support similarly determines the performance of the catalyst, which may also be affected by the exposed sites of the support. To fully understand these effects, a series of supported vanadium oxide catalysts at monolayer and submonolayer coverage have been prepared. The monolayer coverage was determined by Raman spectroscopy and X -ray photoelectron spectroscopy. The activity of the supported vanadium oxide catalysts is determined by the specific support and surface vanadia coverage.

1. INTRODUCTION Supported metal oxides are currently being used in a large number of industrial applications. The oxidation of alkanes is a very interesting field, however, only until recently very little attention has been paid to the oxidation of ethane, the second most abundant paraffin (1). The production of ethylene or acetaldehyde from this feed stock is a challenging option. Vanadium oxide is an important element in the formulation of catalysts for selective catalytic reactions (e. g. oxidation of o-xylene, 1-3, butadiene, methanol, CO, ammoxidation of hydrocarbons, selective catalytic reduction of NO and the partial oxidation of methane) (2-4). Many of the reactions involving vanadium oxide focus on the selective oxidation of hydrocarbons, and some studies have also examined the oxidation of ethane over vanadium oxide based catalysts (5-7) or reviewed the activity of vanadium oxide for the oxidation of lower alkanes (1). Our work focuses on determining the relevance of the specific oxide support and of the surface vanadia coverage on the nature and activity of the supported vanadia species for the oxidation of ethane.

296

2. EXPERIMENTAL 2.1. S y n t h e s i s The oxide supports employed in the present study were: SiO2 (Cabot), A1203 (Engelhard), CeO2 (Engelhard), TiO2 (Degussa), ZrO2 (Degussa), and Nb205 (Niobium Products Co.). All supports were pretreated at 773 K overnight. The silica support was also treated with water vapor, this support will be referred to as SiO2-H20. The catalysts were prepared by incipient wetness impregnation with V-isopropanol in a glove box with nitrogen flow. The impregnated samples were kept at room temperature overnight in the glove box. Then, the samples were dried at 393 K for 1 h and at 573 K another 1 h with nitrogen flow. Finally, the samples were calcined at 573 K for 1 h and 723 K for 2 h in flowing oxygen. The catalysts were prepared with several vanadium oxide loading ranging from very low surface vanadium coverage to the presence of crystalline V205. The catalysts are referred to as "xVS", where "x ~ represents the weight percent of V205 and "S" stands for the element of the specific oxide support. 2.2. Characterization The vanadium content in the catalysts was determined with a PerkinElmer Mod. 3030 Atomic Absorption Spectrometer. The surface areas of the catalysts and of the corresponding supports were determined by nitrogen adsorption/desorption isotherms on a Micromeritics Mod. 2000 ASAP. X Ray photoelectron (XPS) spectra were acquired with a Fisons ESCALAB 200R electron spectrometer equipped with a hemispherical electron analyzer and an MgKa X-ray source (h.v = 1253.6 eV) powered at 120 watts. A PDP 11/05 computer from DEC was used for collecting and analyzing the spectra. The samples were placed in small copper cylinders and mounted in a transfer rod placed in the pretreatment chamber of the instrument. The base pressure in the ion-pumped analysis chamber was maintained below 5 x 10-9 Torr during data acquisition. The spectra were collected for 30 to 100 min at a pass energy of 10 eV (1 eV = 1.602 x 10 -19 J), which is typical of high resolution conditions. The intensities were estimated by calculating the integral of each peak after smoothing and subtraction of the "S-shaped" background and fitting the experimental curve to a combination of Gaussian and Lorentzian distributions, the G/L proportion of which varied in the range 5-27%. All binding energies (BE) were referenced to the support cation, giving values with an accuracy of + 0.2 eV. The molecular structures of the surface vanadium oxide species on the different supports were examined with Raman spectroscopy. The Raman spectrometer system possessed a Spectra-Physics Ar + laser (model 2020-05) tuned to the exciting line at 514.5 nm. The radiation intensity at the samples was varied from 10 to 70 mW. The scattered radiation was passed through a Spex Triplemate spectrometer (Model 1877) coupled to a Princeton Applied Research OMA III optical multichannel analyzer (Model 1463) with an intensified photo diode array cooled to 233 K. Slit widths ranged from 60 to 550]~m. The overall resolution was better than 2 cm -1. For the in situ Raman spectra of dehydrated samples, a pressed wafer was placed into a stationary sample holder that was installed in an in situ cell. Spectra were recorded in flowing oxygen at room temperature after the samples were dehydrated in flowing oxygen at 573 K.

297 2.3. E t h a n e o x i d a t i o n The catalysts (20 mg) were tested for the partial oxidation of ethane with oxygen at atmospheric pressure in the temperature range 760-880 K. The reactor consisted of a quartz tube of 6 mm o.d. (4 mm i.d.), where no void volume was permitted to avoid homogenous reaction from the gas phase. The O2/C2H6 molar ratio was 2 and He/O2 molar ratio was 4. The gas feed was controlled by means of mass flow controllers (Brooks). The total flow range was 15-60 mL/min. The reactor effluent was analyzed by an on-line Hewlett-Packard Gas Chromatograph 5890 Series-II fitted with a thermal conductivity detector. Chromosorb 107 and Molecular Sieve 5A packed columns were used with a column isolation analysis system. The TOF (turnover frequencies, number of ethane molecules converted per surface vanadia species per second) were calculated assuming that all the supported vanadium oxide is active, in agreement with the absence of tridimensional aggregates of vanadium oxide (100 % dispersion). 3. R E S U L T S 3.1. C h a r a c t e r i z a t i o n . The oxide supports used have the surface areas reported in Table 1 and a wide range of vanadium oxide loading on the supports has been prepared. The values presented in Table 1 correspond to monolayer coverage of vanadium oxide. The monolayer coverage can not be determined by theoretical calculation based on the coverage per VOx unit (2), since the dispersion requires an interaction with the support and the monolayer coverage does not only rely on the surface area of the support, but also on its chemical nature. The monolayer coverage of surface vanadium oxide has been determined by Raman spectroscopy and XPS. The low monolayer coverage for the silica support is due to the low surface density of hydroxyl groups. The highly dispersed surface vanadium oxide species are characterized by a Raman band at ca. 1030 cm -1 (Table 1), characteristic of a highly dispersed surface vanadium oxide species. At higher vanadia loadings, crystalline V205 species dominate (strong Raman band at 994 cm-1), and present a significantly different spectrum. The supported vanadium oxide sample with the highest vanadium oxide loading before the onset of crystalline vanadia corresponds to the monolayer catalyst. A parallel characterization has also been performed by XPS since the V/Support atomic ratio determined by XPS is very

Table 1 Characterization of the supports and of the monolaser catalysts Support Surface Area Catalyst %V205 Surface Area Vatoms V=O band (cm -1) BET (m2/g) (wt %) BET (m2/g) per nm 2 1039 SiO2 337 12VSi 11,0 247 2,1 1039 SiO2-H20 332 12VSi-H20 11,7 254 2,3 1026 A1203 222 25VA1 29,9 169 13,4 1028 CeO2 36 4VCe 4,8 23 7,6 1030 TiO2 45 6VTi 6,7 47 9,4 1030 ZrO2 34 4VZr 3,0 31 8,2 1031 Nl~O5 57 5VNb 6,1 35 6,1 1036 5TiSi 280 10V5TiSi ca. 10 . . . .

298

0,4

225, A

o

0,3

I

~A

200

9 VA1

I

mi D

, A VTi A', ,' monolayer

'N

0,2

,

A1

,

!

I |

r

175

~

150~,,

~

152

ca

48

!

0,1

9

l 0 _~' 9149 ' 0 20

% V205

I I I I I I I I

$ !

B

l

!

o

monolayer

~

II

i.v l

I

A VTi

, monolayer 4A I

l

44 40

I

0

I

I

20

40

% V20 5

Figure 1. XPS V/Support atomic ratio (A) and Surface area (BET) oxide loading of the representative series VTi and VA1.

vs.

vanadium

sensitive to the dispersion degree of the surface vanadium oxide species. Figure 1 presents the XPS V/Support atomic ratio determined for dehydrated samples on two representative series (v205friO2 and V205/A1203). The V/Support atomic ratio increases linearly with the vanadium oxide loading and then levels off. This plateau corresponds to the formation of tridimensional aggregates of vanadium oxide, which Raman spectroscopy identifies as crystalline V205. The addition of vanadium oxide to the support continuously decreases the surface area. Close to the monolayer coverage, the surface vanadium oxide species show some polymerization (3,4) as evidenced by the Raman features observed at 920, 800, 600, and 550 cm -1 (8). 3.2. Ethane

oxidation

The oxidation of ethane over the supported vanadium oxide catalysts yields ethylene, CO, CO2, and minor amounts of acetaldehyde and formaldehyde. Methane production was not observed. The initial ethane conversions and TOF's for the catalysts with monolayer coverage of surface vanadium oxide are presented in Figure 2. There are significant differences in the activity of the different monolayer catalysts, where vanadium loading corresponds to the monolayer coverage. The treatment with water appears to increase the ethane conversion on silica-supported vanadium oxide monolayer catalysts. The TiO2 and ZrO2 supports result in the highest TOF. The catalysts with high activity deactivate also after few hours on stream, and the non-selective oxidation products are dominant. The selectivity-conversion plots are presented in Figure 3. Formaldehyde is observed at low conversion for all the catalysts but acetaldehyde can only be observed at low conversions on silica-supported catalysts.

299 0,10

8O 60 0

*,'=4

~J

005 ~

~" 40 o

~

9 20 0,00

0 12VSi- 12VSi H20

4VCe

5VNb

25VA1 4VZr

6VTi

Figure 2. Conversion of ethane, absolute (columns) and T.O.F. numbers (circles) of the catalyst. Total flow 30 mL/min. Reaction temperature 823 K. W = 20 mg. O2/C2H6 = 2 molar and He/O2 = 4 molar. The most active catalysts (6VTi, 4VZr, and 25VA1) show very high selectivity to deep oxidation, mainly CO, which increases at the expense of ethylene with ethane conversion. CO2 is also present and its selectivity increases with conversion too. The catalyst 4VCe and 5VNb catalysts show an important trend to non-selective oxidation at low conversion levels. The selectivity to ethylene decreases markedly with conversion for 4VCe, and this trend is much smoother for 5VNb. The TOF values for the alumina series evidences, that at the surface monolayer coverage of vanadium oxide, the vanadia sites behave differently: nonselective oxidation is dominant and the TOF of ethane and ethylene decreases (Figure 4). The continuous increase in oxygen TOF corresponds to the lower selectivity at vanadium oxide monolayer coverage on alumina. The titanium oxide supported vanadia, on the contrary, shows a more selective activity than alumina. The TOF's of ethylene are higher at submonolayer coverage (1VTi and 3VTi) and decrease at vanadia monolayer coverage. The TOF profiles show increase of CO and CO2 with vanadia coverage but a significant decrease is observed for ethylene. These trends suggest a change in the environment of the vanadia sites as the surface coverage increases. The high activity of titania-supported vanadium oxide and the selectivity of silica-supported vanadia suggests that a ternary catalyst (V205-TiO2-SiO2) may possess positive characteristics. Vanadium oxide tends to preferentially coordinate to titania sites for the titania-silica supports (9). For this reason, the titania-silica support has been prepared with a highly dispersed titanium oxide surface species, strongly interacting with silica support, as determined by Raman spectroscopy and XPS (10). The activity of vanadium oxide on a highly dispersed titanium oxide surface species on silica is compared with the monolayers of

300 80

80

60

1 VSi

60

40

40

20

20

~"

~

9

0 80 60

A

9

-

=

0 80

~

25VA1

40

60

4VCe

4O

0 80

'

0

'

~'--F

m

....

mm

40

40

c~ 20

20

0 8O

0

I

20

40

9 CO A CO2

20

o C2H4 20

I

!

40

60

40

6O

80

% C2H6 Conversion

5VNb

0

I

4VZr

0

0

'

60

6VTi

60

'

80

x

60 ~J

.

O HCHO n CH3CHO

80

% C2H 6 Conversion Figure 3. Selectivity conversion plots for the monolayers of vanadium oxide on the different supports. Reaction conditions: Total flow 30 mL/min. Reaction temperature 740-883 K. W = 20 mg. O2/C2H6 = 2 molar and He/O2 = 4 molar. vanadium oxide on silica and on titania in Table 2. As previously mentioned, the 6VTi catalyst is the most active system (highest TOF), the temperature to reach 15 % conversion of ethane is very low (750 K), but CO is the major oxidation product. The activities of the 12VSi and 10V5TiSi catalysts are similar but the selectivity to ethylene is higher for 10V5TiSi, and CO2 is not produced for this catalyst.

301

Figure 4. TOF numbers for the oxidation of ethane. Reaction conditions as in Figure 2.

Table 2 Selectivity of vanadium monolaser on silica and titania and titania-silica Temp. Conversion % Selectivity Catalyst (K) (mole%) CO CO2 C2H4 HCHO CH3CHO 25VA1 730 15.0 53.2 2.2 44.1 0.2 0.0 4VZr 794 15.0 55.4 5.4 39.1 0.1 0.0 4VCe 891 15.0 54.6 10.0 35.3 0.2 0.0 5VNb 908 15.0 51.1 8.9 39.7 0.3 0.0 12VSi-H20 838 15.0 32.8 7.6 54.6 2.9 2.1 12VSi 847 15.0 31.7 8.3 55.9 2.9 1.1 6VTi 750 15.0 60.5 5.6 33.8 0.1 0.0 10V5TiSi 843 15.0 41.3 0.0 58.2 0.4 0.0 Reaction conditions as in Figure 2.

302 4. D I S C U S S I O N

The dehydrated surface vanadium oxide species on the different supports possess the same structure as previous studies have already shown by Raman and 51V-NMR studies (3,11-13). Surface vanadium oxide species are present as isolated VO units containing one terminal V=O bond and three bridging V-OSupport bonds. Polymeric surface vanadium oxides are also present and their concentration increases with surface vanadia coverage. These species have one terminal V=O bond and possess bridging V-O-V and V-O-Support bonds. The linear increase of the XPS Vfri atomic ratio is consistent with the dispersed nature of the surface vanadium oxide species at low coverage. The alumina supported series, exhibits a lower V/A1 atomic XPS ratio at low vanadium oxide loadings, due to the higher surface area and porosity of the alumina support. The V/A1 atomic XPS ratio is also linear with vanadia loading and supports the highly dispersed nature of surface vanadium oxide species on alumina at coverages below the monolayer. Water treatments of the silica support did not show appreciable differences between the two silica supported vanadium oxide catalysts since both catalysts perform very similar during the oxidation of ethane. Only higher activity is observed for the 12VSi-H20. The selectivity conversion trends suggests that ethane is initially oxidized to ethylene and that ethylene is further oxidized to CO. CO2 could be a primary product, since its selectivity at zero conversion limit does no appear to be zero. The most important differences observed in catalytic behavior results from an interaction of the support with the active phase. For the same reaction conditions, the TOF's differ by more than an order of magnitude for the different catalysts. The changes in TOF's do not correspond with the changes in the terminal V=O Raman bands. A similar result has also been observed for the oxidation of methanol (14) and butane (8). Consequently, the active oxygen must be the bridging oxygen. At monolayer coverage, both V-O-V and V-O-Support bonds are present. Both may play a role in the reaction. The more reducible oxide titania and zirconia yield the most active catalysts (higher TOF's). Ceria is also a reducible oxide, which makes the supported vanadium oxide species more reducible than on alumina or silica, like titania and ceria (1,15), but TOF's on 4VCe are very low. Acidic supports, alumina and niobia, show some moderate activity and the non-acidic, non-reducible silica yields the lowest TOF's. Consequently, the activity of supported vanadium oxide for the oxidation of ethane follows the trends TiO2 N ZrO2 > A1203 > ND205 > SiO2. The low activity of 4VCe and 5VNb catalysts, despite their reducibility (4VCe) and acidity (5VNb) may be due to structural transformations by reaction of vanadia with the underlying oxide at the high temperatures required for ethane oxidation. Concerning selectivity, the more reducible oxide support systems show a high selectivity to deep oxidation (CO). 4VCe shows high selectivity to CO at low ethane conversion. The acidic supports, alumina and niobia, also yield CO as the main oxidation products. Only silica-supported vanadium oxide shows higher selectivites for ethylene. Acetaldehyde and formaldehyde are also produced son 12VSi and 12VSi-H20. The relevance of V-O-V bonds can be evaluated for the performance of V205/A1203 and V205friO2 at different surface coverages. Alumina supported vanadium oxide shows increasing TOF numbers for oxygen, CO and CO2 as

303 vanadium oxide loading increases up to monolayer coverage. At monolayer coverage, where the (V-O-V) / (V-O-Support) ratio is expected to be highest, the TOF's of ethane and ethylene decrease, but TOF of oxygen, CO and CO2 increase. This could be indicative of the higher reducibility of surface polymeric vanadium oxide species with respect to isolated surface vanadium oxide species (4,8), which appears to lead to a less active and selective catalyst. A similar trend is observed for VTi series: at monolayer coverage, ethane and ethylene TOF numbers decrease. For the titania-supported vanadium oxide catalysts, the TOF's for oxygen, CO and CO2 do not increase at vanadia monolayer coverage as in the case of the VA1 series. On the contrary, they decrease slightly, but isolated surface vanadium oxide species on titania are more reducible than isolated surface vanadium oxide species on alumina. This may account for the higher TOF(oxygen)/TOF(ethane) ratio observed on the VTi series. This ratio becomes closer for VTi and VA1 series at monolayer coverage, where both series are expected to show s higher reducibility of the surface vanadium oxide species. The ternary V205/TiO2-SiO2 catalyst shows interesting structural and catalytic properties. Surface vanadium oxide species preferentially coordinate to titania sites in the TiO2/SiO2 supports (8). However, the use of a titania-silica support prepared so that titanium oxide is highly dispersed and strongly interacting with silica support results in titania with different characteristics to pure titania. The titania-silica support used here has 20% of the titanium atoms in tetrahedral coordination as determined by XPS and no crystalline aggregates of titania are formed, as determined by Raman spectroscopy (10). The V=O mode observed for the dehydrated 10V5TiSi catalyst is at 1036 cm -1, much closer to that of silica-supported vanadium oxide than to that of titanium-supported vanadium oxide (Table 1). The surface vanadium oxide species are isolated (100 % dispersion) and must also have a different coordination environment (probably anchored on both, titania and silica sites) that yields an activity similar to that on 12VSi but more selective, since no CO2 is formed and the selectivity of ethylene increases. The lower selectivity of oxygen -containing products suggest that vanadia species on the highly dispersed titania-on-silica supports may be less reducible than on the pure constituting oxide supports. 5. C O N C L U S I O N S The surface vanadium oxide species on silica, water-treated silica, alumina, ceria, titania, zirconia, niobia and titania-silica have been characterized and studied for the selective oxidation of ethane. The terminal V=O bond does not appear to be directly involved in the reaction (no correlation with TOF). However, the bridging V-O-V or V-O-Support bonds appear to critical for the oxidation of ethane. The nature of the V-OSupport bond is determined by the specific support. Bonding to a reducible support metal ion yields active catalysts (e.g. 6VTi and 4VZr). Acidic supports show some activity, but much lower than the reducible ones. The silica support is not reducible and does not possess acidic sites and shows the lowest TOF numbers. However, silica-supported vanadium oxide catalysts possess the highest selectivity. The very low activity of 4VCe and 5VNb could originate from a reaction of vanadia with the underlying support. The surface coverage increases

304 the (V-O-V) / (V-O-Support) ratio. Polymeric surface vanadium oxide species are more reducible than isolated surface vanadium oxide species in the presence of butane (15). If we assume a similar trend of reducibility with ethane than with butane, if turns out that more reducible surface vanadium oxide species are less active and selective. This effect is more evident for the VA1 series than for the VTi series, since the isolated surface vanadium oxide species on alumina are much less reducible than on titania. All the catalysts that show higher reducibility, either due to its interaction with the support or due to its surface polymerisation show lower selectivity. The surface vanadium oxide species have a different environment for 10V5TiSi catalyst, which yields an activity similar to that of 12VSi but is more selective. Further research is going on to fully understand the environments of vanadia sites in this catalyst.

ACKNOWLEDGEMENTS This research has been partially funded by the Fundaci6n Caja de Madrid (Spain).

REFERENCES "

.

3. "

5. 6. .

"

.

10. 11. 12. 13. 14. 15.

E. A. Mamedov, and C. Cortes Corberfin, Appl Catal A : General, 127, 1 (1995) G. Bond, and S. Flamerz Tahir, Appl. Catal., 1 (1991) G. Deo, I. E. Wachs, and J. Haber, Critical Reviews in Surface Chemistry 4 (3/4), 141 (1994) I. E. Wachs, and B. M. Wechkhuysen, Appl. Catal. in press (1997) S. T. Oyama, and G. A. Somorjai J. Phys. Chem., 94, 5022 (1990) J. Le Bars, J. C. Vedrine, and A. Auroux, S. Trautmann, and M. Baerns, Appl. Catal tk" General 88, 179 (1992) M. Merzouki, B. Taouk, L. Tessier, E. Bordes, and P. Courtine, in "New Frontiers in Catalysis" (Guczi et al., Eds.), p. 753. Elsevier, Amsterdam, 1993 I. E. Wachs, J.-M. Jehng, G. Deo, B. M. Weckhuysen, V. V. Guliants and J. B. Benziger, Catal. Today, 32, 47 (1996) J. -Mirn Jehng, and I. E. Wachs, Catal. Letter, 13, 9 (1992) X. Gao, M. A. Bafiares, J. L. G. Fierro. and I. E. Wachs, unpublished results G. Busca, Mater. Chem. Phys., 19, 157 (1988) H. Eckerdt, and I. E. Wachs, J. Phys. Chem., 93, 6796 (1989) J. Hanuza, B. Jezowska-Trzebiatowska and W. Oganowski, J. Mol. Catal., 29, 109 (1985) G. Deo, and I. E. Wachs, J. Catal., 146,323 (1994) J. Haber, A. Kozlowska, and R. Kozlowski, J. Catal, 102, 52 (1986)

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

305

The ethane o x i d a t i v e c h l o r i n a t i o n process and efficient catalyst for it M.R. Flid, I.I. Kurlyandskaya, Yu.A. Treger and T.D. Guzhnovskaya Scientific Research Institute "Syntez", 2 ,Ugreshskaya str., P.O. Box 56, Moscow, 109432 Russia Formation of the mixed cement-containing systems within the range of low copper concentrations with addition of alkali metal dopants as well as catalytical properties of these systems in the ethane oxidative chlorination process have been investigated. Based on the obtained data the efficient and stable copper-cement catalyst has been worked out. This catalyst will assist in the development of a new technology of the vinyl chloride production from ethane. The basic parameters of the ethane oxychlorination process have been determined : at 623-673K, time-on-stream 3-5s and reactant ratio of C2H6: HCI: :02 = 1:2:1 the conversion of ethane is more than 90% and the total selectivity to ethylene and vinyl chloride is 85-90%.

1.1ntroduction The gas-phase catalytic process for oxidative chlorination of ethane to vinyl chloride according to overall equation C2H6 + HC1 + 02 = C2H3C1 + 2H20,

(I)

proceeds in two consecutive kinetically independent reactions: (1) the oxidation of hydrogen chloride to chlorine and (2) the chlorination of ethane. This process is promising for developing a rational technology of vinyl chloride production, because ethane utilized in it is a cheap hydrocarbon raw material [ 1,2 ]. The process is conducted at high temperatures, and ethane converts to vinyl chloride due to a combination of consecutive and parallel radical-chain and heterogeneously catalyzed reactions: oxidation, chlorination, and dehydrochlorination. The contributions of homogeneous and heterogeneous reactions to the overall rate of chlorinated hydrocarbon conversion depends on the temperature ranges at which the reaction proceeds. The process as the whole may be represented by the following schematic diagram [3]: CO + C02

C2H6 - -

t

t

C2H5CI~2H4CI2 C2H4 ~

C2H3C1

t

~2H3C13

t

~2H2C14

--~ C2H2C12 --~ C2HC13

CO + CO2

t

~2HCl5 --~ C2C14

--

C2C16 (II)

306 from which it follows that the major products of ethane oxidative chlorination are ethyl chloride, 1,2- and 1,1-dichloroethanes, 1,1,2-trichloroethane, chloroethylenes, and carbon oxides as the products of deep oxidation. At relatively low temperatures (623m723K), the reaction mixture consists mainly of chloroorganic saturated compounds [3-6]. The situation changes dramatically with raising the temperature. Figure 1 demonstrates the effect of the temperature on the oxidative chlorination of ethane over the well-known conventional salt CuC12--KC1/silica gel copper-containing catalyst. 80

1

70

.

'

~

-

"

60,

-~

3

~,9 50

L 0

,'.o

40

2

30 "0

o

5

20 10

4

"'

0 723

773

823

Temperature,K

Figure 1. The effect of temperature on the ethane oxidative chlorination process (silica gel as the support, copper content of 6.0 wt %, potassium content of 4.0 wt %, reactant ratio C2H6 : HC1 : O2= 1 : 1 : 1, x = 3 s). 1 is the conversion of ethane; 2 is the yield of oxidation chlorination products; 3.4, and 5 are the yields of ethylene, deep oxidation products, and vinyl chloride, respectively ( x is time- on-stream ). Thus, in the presence of traditional catalytic systems, the yield of vinyl chloride to converted ethane does not exceed 35%. The total yield of vinyl chloride and ethylene ranges up to 80%. It was shown [3,5,6] for saturated compounds ethane, ethyl chloride, 1,2dichloroethane, and 1,1,1-trichloroethane that the observed conversion rates are satisfactory described by the equation r, = k,. Pi" Pci2 0"5

(1)

The observed rate constant in equation (1) in this case decreases in the order C2H6 > > C2H5C1 > C2H4C12. The activation energies for the transformations of saturated (130 kJ/mol) and unsaturated compounds (40--90 kJ/mol) differ dramatically; as a consequence, the yield of chloroalkenes increases with temperature. Oxidative chlorination of ethane gives rise to considerable amounts of carbon oxides. The overall rate of these side reactions is described by the empirical equation rco + c02 = ki. pi" Po_~"Pcl2~

(2)

Unsaturated compounds make the dominant contribution to formation of carbon oxides. Whereas the introduction of one chlorine atom into ethylene molecule results in a 7m 8-fold increase in the observed rate constant of deep oxidation, the further increase of chlorine content in molecule diminishes the oxidation of chloroalkenes.

307 It is essential that the reactions of saturated compounds exhibit zero orders with respect to both oxygen and hydrogen chloride and proceed kinetically independently of one another. For the unsaturated compounds, the conversion rates represent complex functions of the reaction mixture composition. Under the conditions when the reaction exhibits zero order with respect to hydrogen chloride, the kinetics of unsaturated compounds oxidative chlorination is described by the equation: 2ko2 po2" ki'pi ri

(3)

--

2ko2 "po2+ L-k,.p, where index i relates to unsaturated compounds. The process of ethane oxidative chlorination imposes heavy demands on the catalysts. The conventional salt supported catalysts are composed of Cu, K, Ca, Mn, Co, Fe, Mg, and other metal chlorides containing various additives; these salts are precipitated on alumina, zeolites, silica gel, and other supports. Catalytic systems that represent solid solutions of iron cations in the lattice of the o~-A1203 and a-Cr203 phases doped with cations, such as K, Ba, Ce, and Ag are also known [7]. The activity of the known catalytic systems and, especially, their selectivity to vinyl chloride are insufficient. In addition, the known catalytic systems tend to rapid deactivation because of gumming and carbonization of their surfaces. The main problem that determines the possibility for industrial utilization of the process is the creation of highly efficient, stable, and selective catalytic systems performing at relatively low temperatures. This problem was alleviated due to the development of a new generation of heterogeneous catalysts based on high-alumina cements and intended for the synthesis of chloroorganic compounds, l These catalysts fortunately combine the properties required in industry and genetically intrinsic to cements thermal stability, high mechanical strength, and basicity of the surface, which prevents its carbonization with the possibility of imparting the system special properties desired in a particular process [8]. The mechanism of the ethane oxidative chlorination process is distinguished by the fact that the catalyst accelerates primarily the reactions of hydrogen chloride oxidation and dichloroethane dehydrochlorination. This necessitates the modeling of cement catalytic system with the surface carrying active sites capable of catalyzing both reactions mentioned. The analysis of the known and our own experimental data indicated that the properties required may be offered by a copper-containing cement-based catalytic system modified with alkali metals. In this catalyst, copper-containing active sites catalyze the oxidation of hydrogen chloride, whereas the activity of the catalyst in the dehydrochlorination reaction is determined by the acid--base surface properties, which are inherent to cements with different phase compositions. The development of this catalytic system made it necessary to investigate the formation process of mixed cement systems within the range of low copper concentrations and with addition of alkali dopants and determination of the correlation between properties of the obtained catalytical systems and their activity in the ethane oxychlorination process.

I 'Fhe catalysts based on high-alumina cements were developed in collaboration with Prof. V.I. Yakerson (Zelinsky Institute of Organic Chemistry, Russian Academy of Sciences, Moscow, Russia) and Prof. E.Z. Golozman (Institute of Nitrogen Industry, Novomoskovsk, Russia).

308

2.Experimental. The catalysts were prepared by chemical mixing of high-alumina cements ( technical calcium aluminate- talyum) or cement-based supports (calcium aluminates with the developed surface area and various CaO/A1203 ratio - galyumin or galyumin C) [9] with the sources of copper and alkali metals in water--ammonia or ammonia ---carbonate media; the mixing was followed by the drying and thermal treatment of the samples obtained. A comprehensive study on the formation of cement catalytic systems was performed by X-ray diffraction, thermal analysis, electronic diffuse-reflectance spectroscopy and IRspectroscopy. Table 1 presents characteristics of some of the investigated catalytic systems. Table 1 Characteristics of copper--cement catalysts No

Sample

Support

Preparation

Phase composition

Ssoec,

conditions

without thermal with thermal m2/g treatment treatment at 673K 1 CuO-K20- Galyumin 348K water-gibbsite, C3AH6, KC1, CaCO3, CuO, 130 CaO-A1203 ammonium CaCO3, CHA, CuO, ,/-A1203,C12A7 solution KC1

2 CuO-Cs20- Galyumin 348K CaO-Al203

C

water-- CsC1, CaCO3, CuO, CsC1, CuO, ammonium C3AH6 CaCO3(calcite), solution C12A7, CaCO3 (aragonite)

15

Kinetic measurements were made at 623 - 773K using a circulatory flow installation. Reactions were studied in the fixed bed catalyst. Time - on- stream was varied within the range 1,5 - 10s at a reactant ratio of C2H6: HCI: O 2 = 1:1 +3,3, : 1 +1,4. Air was used as a source of oxygen. The grains of the catalyst were 0,25 - 0,5 mm in size. The gas was fed at a volumetric flow rate of 600 h - ~ . The catalytic systems were preactivated with a hydrogen chloride nitrogen mixture at 573-623 K. The analyses were based on the chemical methods (determination of hydrogen chloride and chlorine) and the gas chromatography.

3.Discussion and results. The stage of chemical mixing of the catalysts preparation involves the hydration of cements with forming C3AH6 (C is CaO, A is A1203, and H is H20), gibbsite, and calcium carboaluminate as well as the exchange processes with forming CaCO3 and copper hydroxoaluminate (CHA). The depth and the rate of hydration as well as the distinctions in the exchange processes are determined by the type of cement-containing agent. The stage of thermal treatment involves the formation of C12A7, ,/-alumina, and solid solution of aluminum and copper oxides, which is followed by the precipitation of the excess of highly dispersed copper oxide and by the formation of copper aluminate spinels with various degrees of disorder.

309 Thus, copper-containing phases can occur both as free oxide and as the forms bound with the matrix of the support; the concentrations of bound forms increase with temperature and with the duration of chemical mixing. The estimation of the depth of interaction revealed that not only the implantation of the Cu 2. ions into the matrix lattice with forming isolated ions is possible, but the formation of small surface clusters (CuO)• with highly covalent Cu--O bonds. The distribution of catalytically active component between the free oxide, clusters, and ions implanted into the matrix lattice depends both on the conditions of formation and on the composition of the catalytic system as well as on the type of cement-containing agent. As it was shown in [ 10], cement-containing matrix exerts a strong modifying effect on the active copper-containing sites. At equal concentrations of the active component, the activity of copper-containing sites incorporated into the copper--cement catalyst is higher than that of the supported salt catalysts. When the concentration of copper and surface concentration of copper-containing sites are decreased, specific catalytic activity of coppercontaining centers sharply increases. So, at an extended specific surface of the copper-cement catalyst, high catalytic activity to the oxidation of hydrogen chloride can be accomplished even at a low concentration of active copper-containing component provided that the latter is bound with the matrix of the support. The surface area of cement catalysts, which carries aluminum- and calcium-containing oxide fragments, exhibits pronounced acid--base properties. These properties can manifest itself as a catalytic activity to the reactions of dehydrochlorination, which proceed via the formation of donor--acceptor complexes between the substrate and acid or base sites at the catalyst surface. The existence of different calcium aluminate phases in the aluminum-calcium catalysts was proved by diffuse-reflectance IR spectroscopy. The presence of these phases is responsible for the complex structure of the catalyst surface. At the surfaces of these catalysts, calcium ions with lower coordination numbers can occur together with the ions octahedrally surrounded by oxygen anions. These ions can act as balance cations in the structure of C12A7, being responsible for the existence of specific terminal hydroxyls and Lewis acid sites bound to calcium. At the surfaces of galyumins, bridging hydroxyls exhibiting somewhat stronger acid properties are present along with terminal hydroxyl groups. The hydration of galyumin surface can supposedly be attended with the weakening of the A1--O--M bond (M = A1 and Ca) resulting in the appearance of additional strong adsorption sites[8]. The enrichment of surface layer in galyumin C with Ca )-+ ions at the increase of CaO/ AL203 ratio is essential for reducing the yield of deep oxidation products and preventing the carbonization of the surface. The data on the state of copper-containing phases and acid--base properties of active sites occurring at the surface of mixed cement systems, which were presented above, enable us to conclude that these catalysts can be employed in the oxidative chlorination of ethane. It ~is known that the chlorination of ethane with chlorine formed in the oxidation of hydrogen chloride proceeds by a heterogeneous--homogeneous mechanism [3]. This is why the efficiency of cement catalysts was studied separately by the examples of Deacon reaction and dichloroethane dehydrochlorination reaction. It was found that for galyumin-based cement system, the variation of copper content within 8--25% (in terms of CuO) virtually does not affect the rate of chlorine formation. For the oxidation of HC1, the rate constant is 1.2.10 -3 mol HC1/g cat.h. This value is comparable with the rate constant of HC1 oxidation in the presence of copper-containing salt catalysts. The

310 introduction of potassium chloride into a copper--cement system results in a 1.5-fold rise of the rate constant for the HC1 oxidation. Thus, the activity of copper--cement catalysts in Deacon reaction is comparable with that of commonly used salt catalysts. Systematic investigations on the performance of cement-containing catalytic systems with various chemical and phase compositions in the reaction of 1,2-dichloroethane dehydrochlorination with forming vinyl chloride C2H4CI 2 -4

C2H3C1 +

HC1

(III)

revealed that the catalytic activity of these catalysts in the process under consideration is high. At the constant composition of the reaction mixture, the maximum reaction rate was accomplished with using a cement system whose specific surface is 130 m2/g. Thus, at 623K and time-on-stream of 3.8 s, the reaction rate was 0.42--0.46 mol of vinyl chloride per litre.hour. This value is more than two times higher than the reaction rate accomplished with using a well-known supported salt catalyst CsC1--SiO2. It was also shown that the presence of copper in cement catalytic systems does not affect the activity of the catalyst in the dehydrochlorination reaction (see Fig. 2).

70 60

"6 ,..

0

50 40

P,

~

~ L

30

C 0

0

20

0

[

523

.

.

.

.

.

r

. . . . . . . . . . . . . . . . . . . . . . .

573

' ....

623

" ........

673

Temperature,K

Fig. 2. The conversion of 1,2-dichloroethane in the dehydrochlorination reaction at various catalysts as a function of temperature, z = 8 s; 1- galyumin (Sspec = 130 m2/g), 2- galyumin with a dopant of copper (8 wt % in terms of CuO); 3- CsC1/silica gel. Thus, cement-containing systems provide the conversion of dichloroethane to be increased to more than 70% even at 673K. An important positive factor is that vinyl chloride molecule is stable at this temperature. At 673K, the side reaction of vinyl chloride dehydrochlorination with forming acetylene proceeds slowly, acetylene does not form, and the reaction is not complicated by the formation of a number of by-products, for example, of perchloroethylene. Thus, the above-made supposition about bifunctional character of copper--cement catalytic systems was confirmed in the investigations of their activity in the above-mentioned reactions.

31l The oxidative chlorination of ethane as a whole was studied by using of the cementcontaining catalysts with a specific surface of 130 m2/g (sample 1) and 15 m2/g (sample 2). Copper concentration was kept constant and equal to 8 wt % in terms of CuO (see Table 1). It was found during the investigations that when the temperature was raised from 623 to 773 K, the conversion of ethane somewhat increased, and sample 1 exhibited better activity in comparison with that of sample 2. At the moderate temperatures (623--673K), an extended specific surface of sample 1 was favorable for increasing the yield of target unsaturated compounds: ethylene and vinyl chloride. The further temperature increase led to a decrease in the process selectivity because of a noticeable increase in the yield of deep oxidation products, CO• The effect is more pronounced for sample 1 (see Fig. 3). 100 F

...........................................................................................................................................................................................................................

...,,-

90

-,--I'"

.

.

.

.

I-'-

""

""

-6

.~

__,a

-----4

80 7O

.__

..... -1 b

......._-

I

6o 5O

t-

40

o

(..)

30 20 ~C

1~ L 0 ,"F:

623

-.

-

--T

,

673

'"'-

. w

.

723

.

.

.

.

_

.

.~C

._~

773

Temperature,K

Fig. 3. The conversion of ethane and the yields of reaction products for catalysts 1 and 2 as functions of temperature. Time-on-stream of 3 s; 1; the reactant ratio of C2H6 : HC1 : 02 - 1 : 2 : 1;-- - catalyst i ; catalyst 2; a is the conversion of ethane, b is the total yield of ethylene and vinyl chloride to converted ethane, c is the yield of deep oxidation products COx. The dependences shown in Fig. 3 reveal that employing a catalyst with a larger specific surface area with rising temperature would, probably, lead to the deep oxidation of vinyl chloride and, to a lesser extent, of ethylene, resulting in a decrease in the total yield of ethylene and vinyl chloride. A certain increase in the overall yield of CO• products, which was observed for catalyst 2, is accompanied with an increase in the total yield of ethylene and vinyl chloride. This suggests that saturated chlorinated h y d r o c a r b o n s - ethyl chloride and 1,2-dichloroethane m are oxidized predominantly and that the rate of oxidation is lower rate compared to that of the dehydrochlorination of these compounds. Thus, the decrease in specific surface of the catalyst involves a noticeable drop of the yield of deep oxidation products, whereas the yields of vinyl chloride and ethylene remain high. We see little reason in the further cut of the specific surface, because the rate of catalytic dehydrochlorination therewith decreases.

312 The results obtained circumstantially testify that the dehydrochlorination and oxidation reactions proceed at different active sites. It is likely that the oxidation of chlorinated hydrocarbons proceeds at the copper-containing sites. This agrees with the data we obtained in the oxidative chlorination of ethylene [ 11 ]. Taking into account the fact that the value of specific surface is a crucial factor in the choice of catalyst, the further investigations we conducted with using catalyst 2. Both the time-on-stream and the reactant ratio are important chemical engineering parameters affecting the characteristics of the process. It was found that the increase in the time-on-stream at T = 673K can improve both the conversion of ethane and the yield of ethylene. The total yield of chloroorganic products therewith decreases, but the concentration of vinyl chloride passes through a maximum. We also observed an increase in the yield of deep oxidation products COx (see Table 2). Table 2 The effect of time-on-stream on the oxidative chlorination of ethane Catalyst- 8 wt % CuO/cement; T = 673K; reactant ratio C2H6 : HC1 : O2 = 1 : 2 : 1. No.

~, s

Reactant conversion, % C2H6

HC1

O2

Yields scaled to converted ethane, % C2H4C12

C2H3C1

C2H4

COx

1

1.5

80.6

36.0

91.2

25.0

34.1

34.6

2.6

2

3.2

87.5

31.7

89.5

19.4

36.8

40.1

3.2

3

5.6

89.1

30.4

88.6

11.3

38.2

43.5

4.0

4

7.9

90.9

30.0

87.9

10.1

35.6

46.8

6.5

5

10.0

92.7

29.1

86.0

8.2

33.0

47.6

9.7

We can suppose on the strength of the data listed in Table 2 that at the short times-onstream, the major contribution to the formation of deep oxidation products is made by saturated chlorinated hydrocarbons: 1,2 dichloroethane and ethyl chloride. On increasing timeon-stream to more than 6 s, we observed a sharp increase in the yield of deep oxidation products together with the decrease in the yield of vinyl chloride. It is likely that at the longer times-on-stream, the rate of deep oxidation of vinyl chloride would increase and become higher than the rate of dichloroethane dehydrochlorination. Taking into account this fact, we believe that the optimum time-on-stream assuring the best total yield of ethylene and vinyl chloride would be 3--5 s. It was shown in the investigations that the ratio of initial reactants also essentially affects the process. It was found that the excess of hydrogen chloride is favorable for improving the selectivity of the process with reducing the yield of deep oxidation products. At 673K and the reactant ratio of C2H6 : HC1 = 1 : 1, the yield of COx ranges from 6 to 7%; at the reactant ratio of C2H6 : HC1 = 1 : 2, the corresponding yield is 3-----4% (see Table 2). A positive factor is that the carbonization of the catalyst therewith decreases. On the other hand, the increase in the excess of HC1 to ethane up to 3 : 1 involves the decrease in the yield of unsaturated hydrocarbons due to the inhibition of the dehydrochlorination of 1,2dichloroethane and ethyl chloride with hydrogen chloride. The excess of oxygen increases the conversion of ethane mainly due to its oxidation: the yield of carbon oxides increases by 1.8-2 times. Thus, the optimum reactant ratio to provide the best yields of the target products is C2H6 : HC1 : O2 = 1 : 2 : 1.

313 Perfect stability of copper-containing cement catalysts in the oxidative chlorination of ethane was confirmed by their performance for 1500 hours without any decrease in the catalytic activity.

4.Conclusions The results obtained substantiate that the utilization of copper---cement catalysts offers promise for the synthesis of vinyl chloride from ethane at law temperatures in a single step. The proposed efficient and stable copper-cement catalyst will assist in the development of a new technology for the production of vinyl chloride from ethane. This technology is lowwaste and balanced in raw materials with meeting modem requirements of ecological safety. It would be appropriate to conduct the process of vinyl chloride production from ethane, hydrogen chloride, and oxygen in a fixed bed of copper---cement catalyst modified with alkali metals, for example, at 623--673K, time-on-stream of 3--5 s, and reactant ratio of C2H6 : HC1 : 02 - 1 : 2 : 1. Under these conditions, the conversion of ethane is more than 90%, and the total selectivity to ethylene and vinyl chloride is 85-90% at the yield of deep oxidation products COx no more than 3--4%.

REFERENCES

1. Yu.A. Treger, V.N. Rozanov, M.R. Flid, L.M.Kartaschov, Usp. Khim., 57,No 4(1988) 577 2. H.Rigel, H.D.Schindler, M.C.Sze. Chem.Engng.Progr.,.69, Nol0, (1973) 89 3. E.I. Gel'perin, Yu.M. Bakshi, A.K.Avetisov, A.I. Gel'bschtein, Kinet. Katal., 19, No 6 (1978) 527. 4. A.J.Magistro, P.P.Nicholas, R.T.Carrol, J.Organ. Chem., 34 (1969) 271 5. E.I. Gel'perin, Yu.M. Bakshi, A.K.Avetisov, A.I. Gel'bschtein, Kinet. Katal., 20, No 1 (1979) 129. 6. E.I. Gel'perin, Yu.M. Bakshi, A.K.Avetisov, A.I. Gel'bschtein, Kinet. Katal., 24, No 3 (1983) 633. 7. M.M. Mallikarjunan and S. Zahed Hussain, J. Sci., Ind. Res., 42 (1983) 209. 8. V.I. Yakerson, E.Z. Golosman. React. Kinet. Catal. Let., 55, No2 (1995) 455 9. V.I. Yakerson, E.Z. Golosman. Scientific Bases for the Preparation of Heterogeneous Catalysts. VI Intern. Symp.Preprint. 3 Poster Session II Louvain-la-Neuve (Belgium), (1994) 105 10. I.I. Kurlyandskaya, I.G. Solomonik, E.D.Glazunova, E.A.Boevskaya, Yu.M.Bakshi, E.Z. Golosman,V.I. Yakerson, Khim. Prom-st, Moscow, No. 6 (1996) 368. 11. M.R. Flid, I.I. Kurlyandskaya, I.G. Solomonik, M.V.Babotina, Khim. Prom-st, Moscow, No. 6 (1996) 364.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

315

Oxidative Conversion of LPG to olefins with Mixed Oxide catalysts: Surface Chemistry and Reactions Network M.V.Landau a, M.L.Kaliya a, A.Gutman a, L.O.Kogan a, M.Herskowitz a and P.F. van den Oosterkamp b aBlechner Center for Industrial Catalysis and Process Development, Ben-Gurion University of the Negev POB 653, Beer-Sheva 84105, Israel Tel. (972-7)-6472141, Fax.(972-7)-6472902 bKinetics Technology International (KTI) B.V., POB 86, 2700 AB Zoetermeer,The Netherlands,Tel.31 (79)-3531453, Fax.31 (79)-3513561 The catalyic performance of three mixed oxide catalytic systems V-Mo-, V-Mg and RE-LiHalogen (RLH) in LPG oxidative conversion was measured at different O2/LPG ratios, temperatures and WHSV. At high LPG conversions V-Mo-based catalysts yielded low olefins selectivity and high LPG combustion (CB), V-Mg - medium olefins selectivity by oxidative dehydrogenation (ODH) route and medium LPG CB selectivity, while RLH catalysts displayed high olefins selectivity by ODH and cracking (CR) routes at low CB. TP-reaction experiments and the effects of oxygen partial pressure on catalytic performance indicated a dynamic interaction of surface oxygen in the ODH, CB and CR routes. ESCA and TPD measurements detected three types of surface oxygen with different nucleophility and bonding strength. Their distribution correlated with LPG conversion selectivities. A correlation between catalysts acidity, the surface exposed metal cations concentration and the productivity by the CR route was derived. The surface basicity was also significant in olefins productivity by the ODH and CR routes. The selectivity of LPG oxidative reactions were attributed to different intermediates formed on the surface as a result of interaction of C3-C4 paraffins with oxygen atoms of different nucleophility. Both the redox balance of surface metal cations and the acidity-basicity balance are proposed to be significant. 1. I N T R O D U C T I O N Catalytic oxidative conversion of low paraffins into olefins, a potential alternative to steam cracking, is one of the attractive optiopns that could decrease the process temperature, minimize the coke deposition at the reactors walls and increase the olefins productivity. Various catalytic processes for oxidative production of ethylene, propylene and butylenes have been published. A review of the published results measured with individual C2-C4 paraffins [1] allowed to select three most efficient oxide catalyst systems for the study: V-Mo- [2], V-Mg- [3] and Mg-RE-LiHalogen (Mg-RLH) [4]. Comparison of their performance in LPG oxidation showed that V-Mocatalyzed mainly the full paraffins CB, V-Mg- displayed average olefins selectivity producing a large amount of butadiene while the RLH - containing oxide systems showed the highest olefins selectivity at high LPG conversions producing substantial amounts of C2-C3 olefins by CR and ODH routes [ 1] The purpose of this work was to study the states of surface oxygen and relate them to the catalytic performance of selected catalysts: V-Mo, V-Mg and RLH.

2. EXPERIMENTAL

Preparation of Catalysts. V-Mo-catalysts were prepared according to procedure described in [2]. Ammonium metavanadate and paramolybdate were dissolved separately at 70~ third solution containing all the other metal components in form of nitrate salts was mixed with the first two evaporated by mixing. The catalyst material was crushed, sieved, dried at 120~ and calcined at 350oc for 5 h. V-Mg samples were prepared by mixing the MgO obtained by decomposition of Mg(NO3)2 or Mg(OH)2 (with addition of SiO2 or TiO2 powders in some cases) with water solution of ammonium metavanadate (containing metal nitrates in some cases), evaporation the suspension to dryness, dried at 120~ and calcined at 550~ for 6 h. The RLH- catalysts were prepared via an aqueous slurry containing LiNO3, NI-I4-halogen salt, Dy-oxide and the second

316

metal oxide (MgO,Ce-oxide or transition metal oxide). The water was evaporated, the paste dried at 130~ resulting solid was crushed,sieved and calcined at 500oc for 2h and at 750oc for 16h. Catalysts testing. A tubular titanium reactor 17 mm ID and 250 mm length supplied with the central thermowell was designed to test the catalysts over wide range of temperature and various feed compositions. Hydrocarbons - 25wt.% n-C4H10- 25wt.% i-C4H10- 50wt.%C3H8 (LPG artificial mixture) or its components, oxygen and nitrogen were fed separately by mass flow controllers (Brooks Instrument) and mixed in preheater at 450oc. The reactor was inserted into Carbolate tubulat oven, uniformly heated over a length about 50 mm. 1-5 g catalyst diluted with quartz pellets at 1:3 ratio was loaded between layers of quartz pellets. Axial temperature gradient in the catalyst layer during the tests was less than 5~ Homogeneous LPG oxidation in titanium reactor filled with quartz pellets at temperatures lower than 600oc was less than 5 wt.% conversion. The analysis of the reaction products excluding water was performed on line with GC HP-5890 that contained four columns - 45/60 Molecular Sieve 13X, 10 ft x 1/8"; 50 m x 0.53 mm Plot A1203; 80/100 Hysep Q 4 ft x 1/8" and 1 ft x 1/8", with internal switching valves and two detectors TCD and FID controlled by ChemStation analytical software. Selectivity was defined as wt of olefins in product divided by the wt of converted LPG feed. Catalysts characterizations. The catalysts composition was measured by energy -dispersive Xray (EDAX) - JEM-35, JEOL Co., link system AN-1000, Si-Li detector. The surface area was determined using BET method (ASTM 3663-84). Phase composition was measured by XRD in conventional, automated Philips PW 1050/70 diffractometer equipped with a long, fine focus Cu anode tube, 40 kW, 28 mA, a scintillation detector and a diffracted beam monochromator. The phase identification was carried out according to JCPDS-ICDD powder diffraction cards. PHI 549 SAM/AES/XPS apparatus with double CMA and Mg Ka X-ray source has been used for X-ray Photoelectron Spectroscopy (XPS) measurements of the catalysts. After recording general survey spectra, high resolution scans were taken at pass energy (25 eV) for the O ls peaks. The spectral components of O signals were found by fitting a sum of single component lines to the experimental data by means of non-linear least-square c.urve fitting to Gauss-Lorentz shape function using software provided by instruments manufacturer for peaks deconvolution. Care was taken to protect the calcined fresh samples from the contact with atmosphere by pressing them into 10 mm disks and transfering to the ESCA analytical chamber. The quantitative distribution of oxygen atoms with different O ls characteristics as well as total atomic surface concentrations of oxygen were calculated by conversion the peak areas into atomic compositions taking in account the sensitivity factors of all detected elements. Binding energies were referenced to the carbon ls line at 284.5 eV. The TPD and TP-reaction measurements were carried out in AMI-100 Catalyst Characterization System (Zeton-Altamira) equipped with quadrupol mass-spectrometer (Ametek1000). 3. RESULTS AND D I S C U S S I O N

3.1. Phenomenological description of observed catalytic effects Table 1 presents the olefins selectivitiy and productivity measured catalysts at about 30% LPG conversion. The measurements were temperature and O2/LPG ratio, keeping the LPG conversion constant by olefins selectivity is determined by a few basic components increasing in

with all the tested oxide carried out at constant varying the WHSV. The following sequence:

V-Mo- (5.1-8.4%) < V-Mg- (39.2-55.0%) < RLH (67.0-79.0%). The nature of promoters or components in RLH catalysts affected mainly the olefins productivity. The significance of different reaction routes is apparent in Table 2 that compares the CR and CB selectivities measured with selected representatives of the three catalyst groups: V-Mo-Nb-SbCa (Cat.A), 0.07V2Os-Mg(Cat.B) and Mg-Dy-Li-C1 (Cat.C). It also includes the results obtained with a catalyst that yielded a higher olefins productivity where the RLH composition was supported on a transition metal oxide (TM-RE-Li-C1, Cat.D). LPG was almost fully combusted on the V-Mo-catalyst. V-Mg-catalyst converted LPG mainly by ODH and CB routes with about equal efficiency. RLH catalysts enhance the ODH and CR routes with relatively low CB. Table 3 comoares the catalvtic oerformance of M~-Dv-Li-C1 catalyst in oxidation of individual LPG

317

components. All the hydrocarbons were converted mostly by ODH and CR routes, with CR selectivity increasing in the sequence: propane < n-butane < i-butane, so that the contribution of cracking products to total olefins yield was 55-65%. Figure 1 presents the olefins selectivity as a function of LPG conversion. Such plots are commonly used for comparison of low paraffins oxidation catalysts [5,6]. The V-based catalysts showed strong decrease in olefins selectivity with increasing conversion ( more expressed with V-Mo-) normally found with ODH catalysts [5.6], while the selectivity of RLH catalysts was almost independent on LPG conversion. Table 1 Compositions and performance of the tested catalyst belonging to the three selected groups Catalyst group

V-Mo

V-Mg

Catalyst composition

S.A., m2/g

0.09V205 0.74MOO3 0.02Nb205 0.02Sb203 0.13CaO 0.05V205 0.83MOO3 0.12CaO 0.17V205 0.83MOO3

Phase composition Olefins Olefins sel.*),% product. g/gCat., h

14

Sb204, Nb205, 8.4 SbNbO4 [Mo4011]O 5.1 MoO3,[Mo4011]O, 7.5 VMoO14

6 10

0.07V205 0.93MgO 0.07V205 0.93MgO 0.05V20 5 0.79MgO 0.16SIO2

60 100 90

0.05V205 0.94MGO 0.006TIO2 0.004Cr203 0.07V205 0.88MGO 0.05Li20 0.06V205 0.79MGO 0.05Li20 0.1C1

55

MgO, Mg3V208 MgO, Mg3V208 MgO, Mg3V208, Mg2SiO4 MgO, Mg3V208

57 52

Mg-RLH 0.8MgO 0.09Li20 0.002Dy203 0.1C1 0.7MgO 0.09Li20 0.002Ce203 0.21C1 0.39MgO 0.43Ce203 0.003Dy203 0.08Li20 0.1 C1 0.88MgO0.01Li20 0.001Dy203 0.1I 0.82MgO 0.1Li20 0.004Dy203 0.08Br 0.77MgO 0.09Li20 0.005Dy203 0.14F *) T = 585~

V-Mo catalysts T = 500~

0.03 0.028 0.027

44.9 44.3 43.5

0.15 0.15 0.14

39.2

0.13

MgO, Mg3V208 --

55.0 54.5

0.18 0.18

20 18 19

MgO,LiDyO2,Li20 -MgO, CeO2

77.3 79.0 78.5

0.08 0.1 0.1

-15 20

MgO,LiDyO2,Dy203 MgO, DyOBr,Dy203 MgO,LiDyO2,Li20

82.0 70.0 77.3

0.25 0.16 0.02

O2/LPG = 1; LPG conversion -- 30%

Table 2 Performance of selected representatives of the three catalyst groups in oxidation of LPG *) Catalyst

A

B d

a

b

C c

d

a

b

D

a

b

c

c

d

a

b

c

d

8.4

91.6

-- 0.03 44.9 55.1 3.1 0.15 77.3 22.0 39.00.08 74.7 28.0 36.2 1.03

*) Testing conditions as in Table 1; LPG conversion -30%; a-olefins selectivity,%, b -combustion selectivity,%, c -cracking selectivity (C1+C2),%, d - olefins productivity, g/g Cat.h A scheme of LPG reactions is proposed in Fig.2 to show the possible low paraffins transformations according to main three routes. It is based on measured products distributions and

318 90

.....

e

9 O'ql~

r

:

9 4~

9

4P,

41,~

9

;>

O r

,,, 30

0

.

.

.

.

I~

I-

0

A

I. . . . . . .

20

40 LPG c o n v e r t ; I o n ,

A

,.A

~.. 60

%

Figure 1. Olefins selectivity vs. LPG conversion plots for all the testcd catalysts

"~

C2H6

CzH4~

C3H 8 ~ - - . E . _ ~ - C " H ? ~

9 i-C4Hzo

~

8. ~ ,.,,

~

n-C4HI0 . ~ ' " 7.

~ *oa

~.

~

, O ~ , ,...~6 ~._,0.

~.~" .n__,_

.

tg.1

~

L ,-,-

7-- co~_

n-C4H~ ~~ cis,trans-C4H8 t6. H24-O2 - - ~ H20 ~7.CH4 + O 2 ~ CO + H20 ~ i-C4H8 l,.C H 4+O 2 ~ CO 2 + H 2

Cn_xH2(n.x),CH4,Cn_xH2(n.x)+2,H20

CRI 1.3,6,8,15,19 ODH CnH2n+2 CB

CO

,.~ ,o2

r.- CnH2n,H20

2,4,7,14 5,9,10,11,12,13,16,17,18

CO,CO/,H20,Hz Figure 2. Reactions network in oxidative conversion of LPG

8O

319

Table 3 Performance of Mg-RLH catalyst C in oxidation of LPG components *) Paraffin

n-Butane

i-butane

a

b

c

d

a

b

72.2

26.2

50.3

0.13

76.1

20.1

c

propane d

a

58.4 0.19

78.2

b

c

d

21.5

60.1

0.18

*) Testing conditions as in Table 1" Hydrocarbons conversion -- 30%. a - olefins selectivity,%, b combustion selectivity,%, c - cracking selectivity,%, d - olefins productivity, g/g Cat.h kinetic studies. The molar amount of hydrogen detected in products was higher than the amount of olefins could produce without a change in the number of carbon atoms while the amount of consumed oxygen was lower than needed for combustion of hydrogen stochiometrically. Therefore reactions like 5, 10, 13 and 19 in Fig.2 were included in the reactions network assuming production of hydrogen as a result of partial combustion. 3.2. Surface oxygen role

in

oxidative conversion of light

alkanes

Lattice oxygen in metal oxides reacted in catalytic cycles is replenished by reoxidation [7-9]. The effect of O2/LPG ratio on the catalytic performance of three selected catalysts shown in Fig.3 indicates that oxygen from gas phase is a reactant in all the three routes of catalytic conversion. Molecular oxygen could react with adsorbed hydrocarbons or oxygen bonding and activation at the catalysts surface could be nesessary.

90

/ ?

60

Catalyst A

~

80

3

90

atalyst B

1

2 ~

3

60

40

3O

30

0 0

0.5

1

1.5

2

OxygenlLPG molar rallo

2.5

--

0

I

0.5

--I .....

1

t-

1.5

OxygenlLPG molar ratio

t---

2

----t--

0

I

0.4

0.8

1.2

OxygenlLPG molar rallo

Figure 3. Effect of O2/LPG ratio on performance of selected catalysts in LPG oxidation at 585~ 1-LPG conversion, 2 - olefins selectivity, 3 - oxygen conversion, 4 - cracking selectivity Three consecutive runs of n-C4H 10-TP-reaction experiments were carried out with selected catalysts A,B,C and D. 25 cm3/min mixture 9%.vol. n-C4H10-He was fed to the reactor of the AMI-100 Catalysts Characterization System containing 3 g catalyst after heating to 200oc in He flow. Then the temperature was gradually increased at 5OC/min up to 600oc (Cat.A), 750oc (Cat.B,C) and 800oc (Cat.D). After reaching the required temperature, the gas flow was switched to He, catalysts were purged for 1 h,cooled to 200oc. Then the procedure of the first run was repeated. Before the third run, performed at the same conditions, the catalysts were reoxidized in 5%vol Oa-He flow at 550oc for 2 hours with subsequent cooling to 200oc. During the n-C4H10-TP-reaction runs the concentration of n-C4H10 in effluent gas as well as

320

concentrations of C4H8 (ODH product), C2H4,CH4 (CR products), CO2, H20 and H2 (CB productg) were monitored by MS. TP-reaction spectra for butane consumption (similar in shape for all catalysts) is shown in Fig.4a. It could be divided into three parts reflecting different catalysts performance as the temperature increases: I - no butane consumption at low temperature, II increasing butane consumption by fresh and reoxidized catalyst and no consumption with reduced catalyst, III- increasing butane consumption in all the runs that could be a result of other reaction routes (e.g. homogeneous reactions with oxygen evoluted by oxides decomposition). In the second series of TP-reaction experiments, the temperature during butane flow was changed in a ramp mode: it was increased in the same way as in previous series up to value a little higher than it corresponded to the end of the part II and kept constant for 1 hour. In this case (Fig.4b) the butane concentration spectra with fresh and reoxidized catalysts showed a minimum as a result of gradual conversion of surface oxygen while the reduced catalyst did not display defined peaks. The concentrations of all the other compounds monitored by MS displayed a -

(a)

F~gure 4. n-C4H10-TP-reaction spectra recorded with Mg.RLHcatalyst B

321

maximum over the same time pcricxt. The normalized MS peaks intensities lot butane, butylene, ethylene, methane, water, carlx~n dioxide and hydrogen measured at maximum butane consumption lot catalysts A,B,C and D are presented in Fig.5. The peaks normalization was done separately for every experiment, so their relative intensities shown in Fig.5 for different catalysts could not be compared. In all cases the products distribution with fresh and reduced catalysts were close to those measurcd in steady-state experiments, excluding high CO2 evolution with the fresh RLH catalysts. Reducing the V-Mg and RLH catalysts in butane flow almost fully depressed their ODH and CB activity shifting the products distribution in the direction of CR and dehydrogenation while the V-Mo- catalyst in reduced form produced the same CB products as fresh and reoxidized form with lower efficiency. These results are evident for the need for adsorbed oxygen species in the reaction cycles producing products according to the three main conversion routes detected in steady state experiments. Then the differences in performance of the three selected catalysts groups in LPG o~dation is probably caused by different states and concentrations of the surface lattice oxygen atoms. It is widely accepted 18-101 that the ability, of surface oxygens Os to react with hydrocarbons and the type of reaction depend on the distribution of Os among the different species: O2(gas) ~ O2(ads) w," O2-(ads) ~-*'20-(~s) w-~'202-(lattice). The performance of the most strongly bonded lattice oxygen that could be removed at high temperatures by the reaction with hydrocarbons in catalytic cycles is governed by their nucleophilicity being directly related to the effective negative charge and bonding strength [8-11]. Those characteristics together with surface concentrations of different oxygen forms for selected catalysts A-D were measured by TPD and ESCA.

Fi~zure 5. n-C4Hlo-TP-reaction products distribution with catalysts A-D at butane consumption

322

3.3. Surface chemistry characterizations The TPD experiments were carried out with 3 g catalysts A,B,C and D in He flow 25 cm3/min monitoring by MS the evolution of 02, CO2 and H20 over a temperature range 200-800oc ( for VMo- catalyst 200-600oc), heating at 5~ The results presented in Table 4 showed that only V-Mo- catalyst contains comparatively weakly bonded oxygen that could be partially desorbed at the temperatures used in steady-state catalytic tests. The oxygen bonding strength corresponding to Table 4 He-TPD of fresh catalyst Catalysts

Desorbed species:

02

A B C D

H20

CO 2

a

b

c

a

b

c

a

b

c

>600 680 705 ND

>20 40 50 ND

480 560 650 ND

ND 700 720 480

ND 500 80 8

ND 510 640 420

ND 710 720 580

ND 30 300 90

ND 580 600 550

a - Temperature of peaks maximum,~ : b - normalized MS peaks intensity c - Temperature of initial product desorption, oC; ND - not detected the temperatures of initial oxygen desorption and its maxima, increased in catalysts sequence: A
A

B

C

D

3 1 8 220 100 22500

1 ND ND 2 1 100

1 ND ND ND ND 30

1 ND ND 1 ND 130

323

absent in reoxidized catalysts (Fig.5). The hydroxyls are formed during the reaction of butane with fresh or reoxidized catalysts. The amount of evoluted water in butane-TP-reaction experiments was comparable with the amounts of other products (Fig.5). After switching the butane flow to He substantial amount of water was desorbed at the purging stage at a much higher concentration compared to the other products (Table 5). Taking into account that no water evolution was detected at the purging stage as well as during butane-TP-reaction with reduced catalysts (except V-Mo- ) it appears that at least part of surface lattice oxygen reacts with hydrocarbons forming hydroxyl groups that are removed at the purging stage as a result of nonreductive dehydroxylation. The ESCA measurements of the O ls electrons BE carried out with fresh catalysts A-D detected one, two or three bands in RFE-spectra depending on catalysts origin corresponding to the O ls electrons BE range 528.4-529.3 eV (OI), 529.9-530.3 eV (OII) and 531.0-531.8 eV (OIII). The values of total oxygen surface concentrations, O ls electrons BE corresponding to different oxygen states and the relative amounts of those oxygen species exposed at the catalysts surface are shown in Table 6. Decreasing the O ls electrons BE in the sequence OIII >OII>OI reflects increasing of electron density or effective negative charge on oxygen atoms. This corresponds to increasing of oxygen atoms nucleophility (basicity or ability for proton abstraction from hydrocarbon molecule). Other observations showed that: Ar sputtering with increased duration removes from the RFES spectra of RLH catalysts the peaks corresponding to OIII species leaving the OI species unchanged; in case of Vcontaining catalysts Ar sputtering does not affect the shape of the spectra and the O ls characteristic was very close to 530.0 eV observed with pure V205. The ESCA measurements with separate individual oxides, hydroxides and carbonates of the elements building the catalysts compositions showed that the O ls electrons BE values of OIII oxygen species correspond to carbonates, hydroxyls, magnesia or lithia. It could be concluded that OIII oxygen atoms with low nucleophility being included mainly in subsurface species cannot play significant role in catalytic cycles. According to ESCA measurements carried out with vanadium oxide, consistent with the data presented in [ 12], the O ls characteristic of OII oxygen species is very close to the V==O doubly bonded oxygen atoms exposed at the surface of (010) planes of V205 crystals. The oxygen species with O ls characteristic of OI do not exist at the surface of individual main components of catalysts A,B vanadia, molybdena, or exist in small concentration at the surface of magnesia or TM. They -

-

-

Table 6 Characteristics of surface oxygen species in selected catalysts according to ESCA Catalysts A Oxygen surface concentration, % at. 80 Metal cations surface concentration, % at. 20 O Is characteristics of oxygen species,eV: OI -OII 530.3 OIII -Normalized oxygen species concentrations: OI 0 OII 100 OIII 0 1- (OII/total) 0

B 70 30

C 55 48

D 36 55

529.2 530.3 531.8

529.3 -531.0

528.8 530.2 531.5

47 40 13 0.6

48 0 52 1

58 20 22 0.8

form mainly as a result of interaction of those components with additives: V-Mg, Mg-RLH or TM-RLH.This is illustrated in Table 7 for Mg-RLH catalyst. From the results of catalytic tests in butane oxidation it is evident that existence of all the components is essential for the performance of this catalytic system. In case of Mg-RLH system magnesia displayed two RFES peaks of O III oxygen at 530,5 and 531.5 eV corresponding to MgO and Mg(OH) 2 in agreement with [13,14] and less than 10% of oxygen in form of O I (529.7 eV) which could be attributed to cationic

324

vacancies at the MgO surface. Introduction of lithia creates additional O I oxygen species (Table 7) while subsequent introduction of CI and Dy strongly shifts the distribution of surface oxygens into OI direction with increasing the effective negative charge at those atoms. The formation of highly nucleophilic oxygens is a result of changes in coordination and chemical bonds polarity of lattice oxygen atoms caused, for example, by formation of new phases like Mg3V208 where oxygen ions became bridged between V and Mg ions [15], substituting Li into MgO lattice [ 16] or formation of LiC1 crystals covered by thin lithia layer[17]. Decreasing the total oxygen surface concentration in the catalysts sequence A --> D ( Table 6) expose more metal cations and chlorine that behave as electron acceptors. Thus increasing the amount of highly nucleophilic oxygen atoms in the same row as electron donors should be accompanied by substantial changes in catalysts acidity-basicity. Those characteristics were measured by NH3- and CO2-TPD after saturation the catalysts samples with corresponding gases at 40oc. The results are shown in Table 8. V-Mo- catalyst displayed the lowest acidity corresponding to the lowest metal cations concentration but about 50% of the acid sites were strong desorbing ammonia at >250~ The other catalysts contain few strong acid sites but the total acidity strongly increased in the sequence B
Catalysts

529.7 MgO 529.2 0.94MgO-Li20 . . . . 0.994MgO-Dy203 -0.93MgO-0.006Dy203-Li20 529.2 0.87MgO-0.06Li20-C1 0.86MgO-0.006Dy203529.3 0.06Li20-C1 *) T= 550~

--530.3 ---

530.5; 531.5 530.5; 531,5 531.2 531.3 530.6;531.5 531.0

11.4 4.5 25.0 5.2 18.0

55 74 56 70 69

38.0

70

WHSV = 0.33 h -1, O2/C4H10 = 1

Table 8 Acis-base characteristics of selected catalysts Catalysts Acidity: total, ~M NH 3/g - > 250oC/total Basicity: total, l.tM CO 2/g >250oC/total -

-

-

A

B

C

D

30 0.5

81 0.002

140 0.04

340 0.03

0.5 0

7 0.3

4 0.4

4 0.6

to the absence of highly nucleophilic oxygen species. The basicity of other catalysts was about one order of magnitude higher: V-Mg and Mg-RLH catalysts displayed about equal distribution between strong and weak basic sites while at the surface of catalyst D the relative amount of strong basic sites was more than twice higher. It corresponds to apperance of highly nucleophilic oxygen species and increasing their nucleophility from C to D (Table 6).

325

3.4. R o l e of d i f f e r e n t

s u r f a c e s p e c i e s in catalytic cycles

Comparison the surface characteristics of selected representatives of the three catalysts groups with their catalytic performance in LPG oxidation show: i - at temperatures less than 600oc all three LPG oxidative conversion cycles - ODH, CR and CB, are controlled by interaction of hydrocarbons with surface lattice oxygen atoms OI and OII, that form surface OH-groups being removed by dehydroxylation before reoxidation, as indicated from the results of TP-reaction and TPD experiments discussed in the part 3.2. ii - combination of OII oxygen species with low nucleophilicity (basicity) bonded to easy reducible metal cations (V,Mo) with acid sites leads to CB increasing with increased acid sites strength; it was indicated by direct correlation between olefins selectivity measured with catalysts A-D (Table 2) and parameter [ 1-OII/Ototal] (Table 6)reflecting decrease of CB selectivity with decrease of the fraction of OII in all the surface oxygen atoms and furthermore by substantial increase of the strong acid sites concentration from V-Mg to V- Mo (Table 8). iii - combination of OI oxygen species with high nucleophility (basicity) bonded to hardly reducible cations (Mg,RE) with weak acid sites leads to ODH and CR increasing with increased basicity of OI atoms, as indicated by comparing changes in the fraction of OI (Table 6) and their basic strength (Table 8) from catalyst A to catalyst D with CR and olefins selectivities of those catalysts presented in Table 2. iiii - the efficiency of CR conversion route increases with increased weak acidity of the catalyst as indicated from the direct correlation between CR productivity of A-D catalysts that could be easily estimated from the data of Table 2 and surface concentrations of metal cations and chlorine given in Table 6. Based on this information two different modes of paraffins activation are assumed, leading to CB or ODH-CR products depending on catalysts surface chemistry that are consistent with generally accepted models [8-11]. V-Mo- catalyst containing strong acid (electron-acceptor) sites, easy reducible cations andweak nucleophilic (proton-acceptor) oxygen atoms could adsorb the hydrocarbon molecule as a result of hydride-ion abstraction by acid sites. Reduction of metal cation with splitting of one of metal-oxygen bonds and stabilizing the proton and carbanion in form of OH and alkoxy species: CnH2n+2 ?-

O H!1

/z,,

CnH2n+ 1 +

OH OCnH2n 1 (n-l)+ M e - O_ MIe (m--l; . _

0

I] m+

n+Me 0 lvle !

! !

| i

!

RLH catalysts do not contain strong acid sites and easy reducible metal cations but have strongly nucleophilic (proton-acceptor) oxygen atoms and weak acid sites. The hydrocarbon molecule could be adsorbed as a result of proton abstraction by strongly nucleophilic lattice oxygen without splitting the metal-oxygen bond and stabilization of proton and carbcation in form of OH and alkyl species: CnH2n+2 O~ CnH2n+ I O - M e (n'l)+ - 0 - M e n+ - 0 ~ ! i

! i

CnH2n+l O - M e (n-l)+ i i

H O - M e n+ - 0 | i

The subsequent transformations of alkoxy radicals containing strong C-O bonds at the surface of V-Mo- catalyst with weakly bonded oxygen atoms yields preferentially formation of full CB products with some hydrogen evolution, while alkyl radicals stabilized on acid sites at the surface of RLH catalysts as a result of C-H bonds polarization in the strong field of metal-

326

oxygen ion pairs should be preferentially transformed to olefins as a result of further hydrogen abstraction (ODH) or CR. The fraction of CR products in olefins depends on catalysts acidity increasing the lifetime of alkyl radicals on the catalyst surface.The V-Mg- catalyst contained the both types of surface oxygen OI and OII in about equal amounts (Table 6) displaying average acidity and basicity (Table 8) and including easy reducible (V) as well as hardly reducible (Mg) metal cations. As a sequence it showed an average catalytic performance. In both cases the catalytic reaction cycle became closed as a result of dehydroxyation of catalysts surface and further oxygen adsorption-insertion in the oxide lattice that in case of V- or V-Mo-containing catalysts is accompanied by increasing of metals oxidation extent. In addition to further reacting of alkyl and alkoxy intermediates at the catalysts surface with dynamic lattice oxygen they could be desorbed into gas phase and react there homogeneously with gas oxygen as it was demonstrated in [ 13] for V-Mg-catalyst. Testing the RLH catalysts in fixed-bed reactor with void fraction of catalysts layer varied from 28 to 43% showed that this route became significant at temperatures higher than 590oc but no substantial changes in products distribution were observed. 4.

SUMMARY

The RLH-based catalysts display high olefins selectivity at high LPG conversions producing olefins by oxydative dehydrogenation and oxidative cracking. The last charactristics allow them to produce ethylene from LPG that is the main product of steam cracking. Supporting the RLH system at different carriers affects mostly the catalysts productivity. The RLH-based catalysts display about 50% olefins yield with productivity per reaction volume close to steam cracking. The high selectivity of RLH-catalysts to olefins is a result of a definite combination of surface oxygen state, oxygen / metal cations ratio, redox properties of metal cations and acidity-basicity balance. Further studies are needed in order to understand the role of the support and the proper functioning of RE-Alkali-Halogen systems in oxidation of low paraffins. REFERENSES

1. M.V.Landau, M.L.Kaliya, M.Herskowitz, P.F.van den Oosterkamp and P.S.G.Bocqu6, CHEMTECH, 26, No.2 (1996) 24. 2. J.H.McCain, US Patent No. 4 524 236 (1985). 3. H.H.Kung and M.A.Chaar, US Patent No. 4 777 319 (1988). 4. C.J.Conway, D.J.Wang and J.H.Lunsford, Appl.Catal., 79 (1991) L 1. 5. F.Cavani and F.Trifiro, Catal.Today, 24 (1995) 307. 6. S.Albonetti, F.Cavani and F.Trifiro, Catal.Rev.-Sci.Eng., (1996), 413. 7. P.Mars and D.W.van Krevelen, Chem.Eng.Sci.(Special Suppl.), 3 (1954) 41. 8. A.Bielanski and J.Haber, "Oxygen in Catalysis", Marcel Dekker, Ink., New York, 1991. 9. V.D.Sokolovskii, Catal.Rev.-Sci.Eng., 32, No. l&2 (1990) 1. 10. G.Centi, F.Trifiro, J.R.Ebner and V.M.Franchetti,Chem.Rev., 88 (1988) 55. 11. H.H.Kung, Ind.Eng.Chem.Prod.Res.Dev., 25 (1986) 171. 12. J.Zi61kowski and J.Janas, J.Catal., 81 (1983) 298. 13. X.D.Peng, D.A.Richards and P.C.Stair, J.Catal., 121 (1990) 99. 14. J.C.Fuggle, L.M.Watson and D.J.Fabian, Surf.Sci.,49 (1975) 61. 15. H.H.Kung, Adv. in Catal., 40 (1994) 1. 16. T.Ito, J.X.Wang, C.H.Lin and J.H.Lunsford, J.Amer.Chem.Soc., 107 (1985) 5062. 17. D.Wang, M.P.Rosynek and J.H.Lunsford, J.Catal., 151 (1995) 155.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

327

Free Radicals as I n t e r m e d i a t e s in Oxidative T r a n s f o r m a t i o n s of L o w e r Alkanes. M. Yu. Sinev, L. Ya. Margolis, V. Yu. Bychkov, and V. N. Korchak Semenov Institute of Chemical Physics, Russian Academy of Sciences 4 Kosygin street, Moscow 117334, Russia Catalytic oxidative transformations of lower alkanes attract the attention as possible ways to transfer these substances into more suitable chemicals - olefins and oxygenates (alcohols, aldehydes, acids, etc.) - and to involve them into the industrial use as raw materials for chemical and petrochemical synthesis. However, the yields of desirable products reached up to date are not sufficiently high. The progress in the studies of intrinsic mechanism of catalytic partial oxidation of lower alkanes is not sustainable either. We believe that these two facts are correlated and that the analysis we performed in the present work can brighten up some important details of the mechanism of catalytic oxidation of lower alkanes*. Experimental facts and theoretical concepts existing in the literature indicate that the formation of free radicals plays an important role in a number of catalytic oxidation reactions [1-5]. In the present paper we analyze the contribution of free radicals to several oxidative transformations of lower alkanes over oxide catalysts. Based on the thermochemical data and on the results of kinetic simulations it is shown that the observed reaction kinetics and product compositions in the mentioned above processes are determined by a set of interdependent heterogeneous and homogeneous reactions of free radicals, i.e. they should not be considered as "spectators" taking part in side reactions, but as principal intermediates causing the main features of lower alkanes oxidation and design of catalysts. 1. ACTIVATION OF ALKANE MOLECULES Alkane molecules do not have any specific "reactive centers", like functional groups or multiple bonds. This means that their activation can be carried out only by the bond dissociation or charge transfer processes. It is evident that the more energetically favorable is the first step of activation, the higher is the probability of its contribution to the overall reaction. In other words, using the value of the energy expenditures in different elementary steps ( Eex = AH - Est, where AH is the overall enthalpy change and Est is the energy of stabilization of the activated molecule or its fragment on the catalyst surface) one can estimate which one of them is more feasible. Such an analysis may be performed on the basis of the available thermochemical data (see, for example, [6]). The possible processes and corresponding values of Eex which we have considered are given below: * This study was carried out under the financial support of the Russian Foundation for Basic Research (research grant No. 96-03-32440)

328 (i) homolytic dissociation of C-H bonds accompanied by the formation of surface OH-group and free radicals [O] + RI-I => [OH] + R

(1)

Eex = DR_H where [O] - strong oxidizing surface center having a high affinity to the hydrogen atom; D i - energy of corresponding bond dissociation; (ii) heterolytic C-H bond dissociation with a proton abstraction on a strong basic center [ 0 2-] + RH => [ 0 2- ... H +] + R-

(2)

Eex = DR_H + IH - IRwhere I i - ionization potentials of corresponding particles; (iii) heterolytic C-H bond dissociation with a hydride-ion abstraction on a Lewis acidic site [M n+] + RH => [M n+ ... H-] + R +

(3)

E e x - DR_H + I R - I H(iv) ionization of alkane molecule [h+] + RI-I => [h+ ... e-] + RH +

(4)

Eex = IRH where [h+]

- hole center;

[h+ ... e-] - trapped electron. It is easy to demonstrate that the sign and the magnitude of energy changes in all these processes depend on the compensation of Eex by binding the fragments of the activated molecule to the surface centers. The values of Eex given in Table 1 show that the energy which has to be compensated is minimal for the process (1). On the other hand, one may assume that the energy of the O-H bond formed in this process is comparable to Eex, i.e. the second fragment (free radical R) should not be bound to the surface in order to compensate the energy expenditure. This assumption is in a good agreement with calorimetric measurements, according to which in the case of oxide catalysts active in oxidative coupling of methane (OCM) and oxidative dehydrogenation of ethane [7,8], as well as in total oxidation of alkanes [9] the O-H bond strength ranges from 250 to 470 kJ mo1-1. On the contrary, in the case of heterolytic C-H bond dissociation the energy expenditure is so large that its compensation requires the binding of both fragments, which must occur in a

329

Table 1 Energy expenditure on the activation of lower alkane molecules Energy expenditure ( kJ mol "1 )

Molecule

reaction (1)

reaction (2)

reaction (3)

reaction (4) .

CH4

431

1630

1308

1250

C2H6

410

1615

1183

1120

C3H8

398

1609

1162

1078

n-C4HI0

393

1605

1154

1037

iso-C4H10

389

1601

1120

1016

single reaction step. This requires the presence of paired centers with specific configurations and energy relations. Such a mechanism including the simultaneous abstraction of H + and H ions from n-butane as a first step of maleic anhydride formation was suggested by Trifiro et al. [ 10] to explain the unique properties of V-P-O catalysts. One may assume, however, that due to steric restrictions the smaller an alkane molecule, the lower is the probability of such a synchronic mechanism. The process (4) seems to be improbable because it requires the existence of the hole centers with an electron affinity comparable with ionization potentials of alkane molecules. The above analysis shows that the formation of free radicals in the interaction of alkane molecules with the surface of oxides may prove to be energetically preferable as compared to any other mechanisms of their activation. Furthermore, this process requires only one type of single active centers and it proceeds in a single step. The combination of these factors may render this process the most favorable. This conclusion is experimentally confirmed by the correlation between the concentration of strong oxidative sites and the catalytic properties of the oxide catalysts (see, for instance, [ 11,12]); the Polanyi-type relation between the activation energy for the oxidative dehydrogenation of alkanes in their interaction with oxides and the energy of O-H bonds formed simultaneously [13-15]. The difference in the reactivities of C1 - C4 alkanes is mainly caused by the difference in the Eex values. Estimations based on the data of Table 1 show that the difference in rate constants at 700 - 1000 K between methane and butanes over the same catalyst can exceed 103. The H-atom affinity in the case of efficient catalysts for methane activation should be the highest. As a result, if the O-H bond strength is high enough to compensate the energy expenditure in the reaction (1), the process of active sites regeneration (reoxidation) becomes more impeded and the difference in the optimal reaction temperatures for different alkanes can reach 100 K or more. ff the catalytic oxidation of alkane molecule starts with the formation of a free radical on the surface of an active catalyst particle and its escape to the gas phase, the complete reaction network includes both homogeneous and heterogeneous steps of the transformation of primary (CnH2n+l) and secondary radicals. Since all these processes are sufficient for the formation of the final products, the analysis of the influence of different factors on the -

-

330 selectivity of a complex heterogeneous-homogeneous process can be carried out by considering the elementary reactions of free radicals. 2. ELEMENTARY REACTIONS OF FREE RADICALS 2.1 H o m o g e n e o u s

reactions

The main types of primary gas-phase transformations of free radicals formed in the reaction (1) are - recombination CnH2n+l + CnH2n+l (+ M)=> C2nH4n+2 (+ M*)

(5)

- H-transfer and elimination (ifn >__2) CnH2n+l + CnH2n+l => CnH2n+2 + CnH2n

(6)

or CnH2n+l + 0 2 => CnH2n + HO2

(7)

or

CnI-I2n+l => CnH2n + H

(8)

- oxygen molecule capture CnH2n+l + 0 2 <=> CnH2n+102

(9)

- oxidation Cnn2n+l + 0 2 => CnH2nO + O H

(lO)

or CnH2n+l + 0 2 => CnH2n+lO + O

(ll)

Reactions (5)-(8) and (10) lead to the formation of stable molecules (hydrocarbons and aldehydes). Subsequent reactions of peroxy- (CnH2n+lO2) and oxy-radicals (CnH2n+lO) formed in reactions (9) and (11) lead to the formation of oxygenates (alcohols, aldehydes, etc.), carbon oxides, and/or olefins. The fractions of radicals transformed into different fmal products depend on the reaction conditions (temperature, oxygen pressure) and on the number of carbon atoms in the alkane molecule. For example, the stability of peroxy radicals decreases with increase of the number of carbon atoms in the alkyl fragment, that is why the probability of total oxidation via their subsequent transformations decreases from methane to

331 butanes. The higher temperature also decreases this probability, due to the shift of the equilibrium (9) towards alkyl radicals, increasing the fraction of radicals transformed into the products of coupling and dehydrogenation. However, if the temperature increases beyond some certain value, the fraction of oxygen containing products increases again because of more sufficient contribution of reactions (10) and (11). In particular, the low efficiency of reactions of metyl radicals with 02 molecules likely causes the existence of a temperature "window" for the OCM process at 900-1100 K. The development of chains in the gas phase leads to the acceleration of the secondary radicals formation, as well as to the additional conversion of the initial reactants and to the shifts of product selectivities. 2.2. H e t e r o g e n e o u s r e a c t i o n s

As we have mentioned above, if the catalyst pores are sufficiently narrow, i.e. the species diffuse through them in Knudsen or transitional re~mes, the contribution of heterogeneous reactions of free radicals to the overall reaction rate and selectivity may be predominant. The main types of elementary reactions between radicals and surface sites proposed elsewhere [ 15] are the following: - H-atom transfer, for example [O] + CnH2n+l => [OH] + CnH2n

(12)

- O-atom transfer, for example [O] + CnH2n+l => [ ] + CnH2n+~O

(13)

- radical capture [0] + CnH2n+~ => [OCnH2n+q

(14)

Let us consider the possible role of these reactions in the formation of final products. If n > 2, the successive reactions (1) and (12) lead to the formation of the desirable product in the case of oxidative dehydrogenation processes. The possible contribution of alkoxy radicals to the formation of reaction products is already mentioned above. In this section we should emphasize that the relative probability of reactions (13) and (14) depends on the properties of the catalyst (oxygen binding energy E[ol) and on the reaction temperature: the higher the temperature and the lower the E[o], the more probable is the reaction (13). In this case one may expect an increase of selectivity of partial oxidation to oxygenates. The fate of radicals captured by the surface sites with the formation of the alkoxy groups depends on the number of carbon atoms in the alkane molecule, as well as on the properties of the catalyst surface. According to the data obtained by Aika and Lunsford with the use of IR spectroscopy and TPD [ 16], in the case of MgO (an oxide with very high E[ol) the methoxy groups decompose forming CO and H2, but in the case of higher alkoxides the formation of corresponding olefins takes place.

332 Taking into account the whole set of homogeneous and heterogeneous reactions, one may conclude that depending on the target product the requirements to the catalyst and to the reaction conditions should be different: if we wish to increase the yield of oxidative dehydrogenation products, we have to increase the temperature and to use the catalysts with higher E/ol. The rigidity of these requirements increases as the number of carbon atoms in alkane molecule decreases due to the increasing strength of C-H bonds and stability of peroxy-radicals. On the contrary, the lower the temperature and oxygen binding energy Eiol, the higher is the probability of the oxygenate production. We have to notice, however, that these requirements are contradictory if the olefm is an intermediate for the further formation of oxygenates. In this case it is substantial that the catalyst contains the active sites of different types: one of them (strongly-bound oxygen with high H-atom affinity) is responsible for the formation of free radicals, and the second one (with lower Etol) supplies O-atoms for the insertion into the organic molecule. The efficiency of oxides containing vanadium, molybdenum, tungsten, and similar cations as catalysts for partial oxidation of hydrocarbons to oxygenates is likely due to the division of functions between oxygen species of different types (terminal M n+ = O and bridge Mn+-O-M n+) which these cations are able to form. An analogous co-operation of oxygen species is likely to take place in the case of multicomponent catalysts: if one oxide phase actively produces free radicals which react subsequently with weakly-bound oxygen species on the surface of another component, the total rate of the final product formation will be higher as compared to that measured over each individual oxide. This explanation is alternative (or complementary) to the so-called "oxygen remote-control mechanism" discussed by Weng and Delmon [17], according to which a synergy in catalytic action may be caused by the transfer of active oxygen species between two oxide phases with different donor-acceptor properties. 3. SIMULATIONS OF SURFACE-ASSISTED FREE-RADICAL PROCESS The preliminary analysis of elementary reactions of free radicals in the presence of an active catalyst demonstrates that the heterogeneous generation of primary radicals initiates the homogeneous processes. In their turn, both primary and secondary free radicals affect reciprocally on the surface active sites. A kinetic model which considers the heterogeneous and homogeneous transformations as interdependent and presumes that all the particles (stable molecules and free radicals) present in the gas phase oxidation undergo both homogeneous transformations and interactions with active sites of the catalyst surface was discussed elsewhere [18]. This model was previously used to simulate the OCM reaction in a quasihomogeneous system. Taking into account that the subsequent fate of the radicals formed in the reaction (1) depends, in the general case, on the relation between the number of collisions with other particles in the gas phase and with the surface and also on the nature, concentration and reactivity of the surface centers, we utilized the approach proposed in [18] to simulate the heterogeneous-homogeneous oxidation of methane in combination with mass-transfer in the gas-solid system. The reaction space was considered as a gas volume of a varied thickness (L) exposed to a flat surface with a varied concentration of active sites (C). The results of simulations of the reaction accompanied by the one-dimensional masstransfer directed normally to the surface are given in Fig. 1-2. If C = 0, a self-acceleration

333

d[CH4]/dt, -

100

C

nmol/s.

5"10

-

-

16

m-2

5 " 1 0 15 hi-2

10-

5 " 1 0 14

-2

m

i

0.1

I

I I

I I

I

I

i

i

I

1 I

-3

t

I

-2

I

I I

I

I 1 I

I

I

i

-1

I

I

0

I I

I

I

I

I

I

i

I

log t (s.)

Figure 1. Methane conversion rate as the fimction of time at different concentrations of active sites on the surface (1000 K, 1 atm., CH4 902 = 10 91, L - 1xl0 -4 cm)

W(s)/W(tot)

log(lO 4 to.t) gas r e a c t i o n -4

0.8

-3

0.6 -2 0.4>

0

o _

i

i

i 4

0

-6

-5

-4

-3

-2

-1

"-

0 log L ( c m )

io

Figure 2. Effect of gas volume thickness on the fraction of the rate of heterogeneous reaction in total conversion and on the time of 1 0 % oxygen conversion ( 1 0 0 0 K , 1 atm., CH4 9 0 2 1 0 " 1, C = 5 x 1 0 1 6 m 2)

334 typical for chain reactions with branching was reproduced. At C = 5x1016 m -2 a kinetic behavior of that kind disappears, and the process becomes "linear", i.e. its rate reaches the maximum at t = 0 and then declines due to the consumption of the reactants. At intermediate C values the gradual transformation of kinetic behavior from a self-acceleration type to a "linear" type takes place. The effect of the gas volume thickness on the contribution of the surface reaction to the overall kinetics is presented in Fig.2. At L = 10 nm the time of 10% conversion characterizing the rate of reaction is -104 times less than in the case of a homogeneous gas reaction and the fraction of the rate of heterogeneous reaction Ws in the total conversion rate Wtot is nearly 1. At increasing thickness of the gas volume the fraction of the heterogeneous reaction and the rate of overall process both decline. However, even at L = 1 cm, the reaction occurring in the gas volume still experiences the influence of the surface taking part in the radical reactions. Taking into account that the specific surface areas of the oxide catalysts usually used for partial oxidation range between-~1 and 20 m2g1, the characteristic size of solid crystallites is -~10-5-104 cm and the role of heterogeneous reactions of radical species in the catalyst pores is likely predominant. According to the results of kinetic simulations, if the specific surface area of the catalyst is more than-~1 m2g-1, the most of radicals formed in the reaction (1) undergo the reverse transformation into the initial alkane molecules: [OH] + CnH2n+l--> [O] + CnH2n+2

(-1)

This means that, although the surface of pores is much larger than the outer surface of the grains, the contribution of the latter to the formation of the final products can be sufficient due to the lower probability of the reaction (-1) for the radicals formed outside the pores. The grain size which usually ranges f r o m - 0.1 mm to few centimeters, makes it possible for the homogeneous chain reaction to develop in the free volume of the catalyst bed (see Fig.2). Recently we observed the effect which supports the conclusion about the substantial role of the radical reaction outside of the catalyst grains. When a very efficient OCM oxide catalyst (10% Nd/MgO) was placed into the reactor together with an inactive metal filament (Ni-based alloy) the sharp increase of conversion accompanied by the selectivity shift from oxidative coupling to the formation of CO and 1-12 was observed [19]. Since the metal component has a low activity also with respect to ethane oxidation, this behavior is not due to successive oxidation or decomposition of C2 hydrocarbons on the metal surface, but should be attributed to the reactions of methane oxidation intermediates. Almost total disappearance of ethane (which is a product of CI-I3 radicals recombination) and acceleration of the apparent reaction rate by the addition of an "inert" material indicate that the efficiency of methane oxidative transformations can be substantially increased if the radicals have a chance to react outside the zone where they formed and the role of reaction (-1) decreases. Although the second (metal) surface is not active enough to conduct the reaction of saturated hydrocarbon molecules (methane and ethane), the radicals generated by the oxide can react further on the metal surface. As a result, the fraction of the products formed from methane activated in the reaction (1) increases, and the formation of the final reaction mixture of different composition takes place.

335 4. CONCLUSIONS 1. The most energetically favorable process of lower alkanes activation over oxide catalysts is a homolytic C-H bond dissociation with the formation of free radicals. The difference in energy expenditures for the formation of free alkyl radicals cause the difference in reactivities between C1-C4 alkanes. 2. The main factors determining the efficiency of different oxides as catalysts for lower alkanes oxidation are the H-atom affinity of strong oxidizing surface sites and the oxygen binding energy. These thermochemical factors cause the rates and directions of free-radical reactions and, as a result, the catalytic activity and selectivity to certain products. 3. The total rate of reaction and the selectivity to different products (olefins, oxygenates, carbon oxides) depend on relative efficiencies of different transformations of free radicals in the gas phase and in the heterogeneous steps, as well as on the transport phenomena. REFERENCES

1. 2. 3. 4. 5.

V.M. Polyakov, Usp. Khim., 17 (1948) 351. D. J. Driscoll, K. D. Campbell, and J. H. Lunsford, Adv. Catal., 35 (1987) 139. J.H. Lunsford, Langmuir, 5 (1989) 12. T.A. Garibyan and L. Ya. Margolis, Catal. Rev., Sci. Eng., 31 (1989-1990) 35. M. Yu. Sinev, L. Ya. Margolis and V. N. Korchak, Usp. Khim. (Russ. Chem. Rev.), 64 (1995) 373. 6. V.N. Kondratiev (ed.), Chemical Bond Dissociation Energies, Ionization Potentials and Electron Affinities, Handbook, Moscow, Nauka, 1962 (in Russian). 7. V. Yu. Bychkov, M. Yu. Sinev, V. N. Korchak, E. L. Aptekar' and O. V. Krylov, Russ. Kinet. Catal., 30 (1989) 1137. 8. M. Yu. Sinev, V. Yu. Bychkov, V. N. Korchak, and O. V. Krylov, Catal. Today, 6 (1990) 543. 9. V. Yu. Bychkov, M. Yu. Sinev, Z. T. Fattakhova, and V. N. Korchak, Russ. Kinet.Catal., 37 (1996) 366. 10. G. Centi, F. Trifiro, J. R. Ebner, and V. M. Franchetti, Chem. Rev., 88 (1988) 55. 11. D. J. DriscoU, W. Martir, J.-X. Wang, and J. H. Lunsford, J. Am. Chem. Soc., 107 (1985) 5062. 12. M. Yu. Sinev, V. Yu. Bychkov, Yu. P. Tulenin, B. V. Rozentuller, and A. M. Rajput, 9th Soviet- Japanese Seminar on Catalysis, Novosibirsk, Nauka, 1990, p. 75. 13. A. A. Bobyshev, V. A. Radtsig, Russ. Chem. Physics, 7 (1988) 950. 14. A. A. Bobyshev, V. A. Radtsig, Russ. Kinet. Catal., 29 (1989) 638. 15. M. Yu. Sinev, Catal. Today, 13 (1992) 561. 16. K.-I. Aika and J. H. Lunsford, J. Phys. Chem., 81 (1977) 1393. 17. L. T. Weng and B. Delmon, Appl. Catal., 81 (1992) 141. 18. M. Yu. Sinev, Catal. Today, 24 (1995) 389. 19. Yu. P. Tulenin, M. Yu. Sinev, and V. N. Korchak, 1 lth Int. Congress on Catalysis, June 30 - July 5, 1996, Baltimore, ML, USA, Programme and Book of Abstracts, P-275.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

337

Alternative m e t h o d s to prepare and m o d i f y v a n a d i u m - p h o s p h o r u s catalysts for selective oxidation o f h y d r o c a r b o n s . V.A.Zazhigalovla, J.Haberlb, J.Stochlb, A.i.Kharlamov 2, i.V.Bacherikovala and L.V.Bogutskaya Ia Ukrainian-Polish Laboratory of Catalysis: a) Institute of Physical Chemistry, National Academy of Sciences of Ukraine, Pr.Nauki 31, Kyjiv-22, 252022 Ukraine b) Institute of Catalysis and Surface Chemistry, Polish Academy of Sciences, Niezapominajek, Krakow, 30-239 Poland 2 Institute for Materials Science Problems, National Academy of Sciences of Ukraine, Kryjanovski 3, Kyjiv, 252680, Ukraine Among numerous compounds formed in vanadium-phosphorus-oxide system, vanadyl pyrophosphate is known to be an efficient catalyst for C4-C5 paraffins partial oxidation [1 ]. Typical process of its synthesis can be represented by a following scheme: ROH, Reductant V 2 0 5 -k- H3PO4

> V O H P O 4 . 0 . 5 H 2 0 . . . . . 5> ( V O ) 2 P 2 0 7

It has been established that the properties of vanadyl pyrophospate are strongly dependent on its biography, i.e. the preparation method, presence of overstoichiometric phosphorus and additives [1-4]. Therefore, considerable effort of the researchers was directed to optimization of the synthesis technique and in recent studies also, non-traditional methods for the catalysts preparation were considered [5-10]. It has been shown by us [9,10] that mechanochemical treatment is a perspective method to modify the properties of the precursor VOHPOa.0.5H20 and thus, to influence the catalyst prehistory. The present paper deals with the possibilities of mechanochemical and barothermal treatments applied at different stages of the catalyst synthesis: the initial reactants, the precursor and the final catalyst. 1. E X P E R I M E N T A L

V205 (purefor analysis) and H3PO4 (pure) were used as initial reagents. The synthesis of the precursor of VPO catalysts was carried out according to the procedure described in [11 ], starting from V205 and H3PO4 in butanol medium in the presence of organic reducing agent. The solid product, after filtration, was heated stepwise up to 300 ~ in vacuum (total time was 60 h). The activation* of the precursor i.e. its transformation into the (VO)2P207 phase was performed in a reactor, just before the catalytic test, with a gas reaction mixture consisting of 1.7% Call10 in air. The activation was carried for 72 h at the temperature gradually rising up to 440~ Witht the exception of samples activated by means of mechanochemistry

338 Mechanochemical treatment was applied at different stages of the synthesis described by scheme 1: to the starting reagent V205, precursor VOHPO4.0.5H20, final product (VO)2P207 and mixtures of the powder VPBiO precursor + La203. It was carried out in a planetary mill at 3,000 rpm. La203 was prepared prior to the milling, by decomposition of La2(CO3)3.xH20 (Aldrich) in an inert gas flow. The solids for the treatment were either suspended in ethanol or water or used without any dispersant (dry milling). Barothermal synthesis and treatment were carried out in a stainless steel autoclave lined with internal Teflon glass (V - 20 cm3). The "barothermal" procedure included its both well known variants named hydrothermal synthesis (in the presence of water) and organothermal one (in the presence of organic compounds) as well as synthesis without any solvent. For the synthesis the powders of starting compounds and phosphoric acid with/without solvent were loaded into the Teflon glass. For modification of VPBiO, precursor grains (D = 5 mm, L = 5-6 mm) were placed in the glass mold in the reactor, and the modificator was located in the space between the glass and autoclave walls. Different temperatures and times of treatment were applied in these experiments. The method of the barothermal treatment was described in details in [ 12]. Phase composition of the samples was analyzed using DRON-3M X-ray diffractometer with Cu Ka radiation. The specific surface area (SsA) of the samples was measured by BET method on Gasochrom-1. Thermal analysis was carried out with the thermoanalytical instrument Derivatograph Q-1500 D (system F.Paulik-J.Paulik-L.Erdey) in helium atmosphere at a heating rate of 10 K/min. The surface composition was examined using VG ESCA-3 X-ray photoelectron spectrometer (A1 Kal.2). The spectra were calibrated against C ls (284.8 eV) line as the standard in the binding energy determination. Jeol-100 CX transmission electron microscope and Nanoscope scanning tunneling microscope were used for the investigation of morphology. Details of the measurements and data processing are given in [ 11 ]. Catalytic properties of the synthesized samples after activation were examined in the hydrocarbon-air reaction mixture in reactions of the oxidation of: i) n-butane (1.7 vol. % in air) to maleic anhydride, ii) butene-2 (1.6 vol. % in air) to maleic anhydride, iii) n-pentane (1.2 vol. % in air) to maleic and phthalic anhydrides, and iv) propane (1.8 vol. % in air) to acrylic acid. Catalytic tests were performed in the flow system with GC control of the reaction products. 2. RESULTS AND DISCUSSION

2.1. Mechanochemistry 2.1.1. Mechanochemical modification of the initial reagents for synthesis of VPO catalyst To prepare VPO precursor (P/V = 1.15) samples of the V205 reagent untreated (V205-R) and after wet (in ethanol, V205-E) and dry (V2Os-D) milling were used. Some properties of these solids are given in Table 1. It has been established that mechanochemical treatment increases the specific surface area of V205 and produces V 4+ ions. The latter phenomenon is indicated by an appearance of the low-energy contribution in the XPS spectrum of V 2p electrons. The STM study [13,14] showed that after the mechanochemical treatment in ethanol the change of V205 texture took place due to an anisotropic plastic shift deformation along the planes parallel to (001). This leads to an increase of the relative exposure

339

Table 1. Properties of V2_Qs_before and after its mechanochemical treatment Sample

V205-R V2Os-D V2Os-E

SSA

Treatment Medium Time, min.

m2/g

Air Ethanol

3.8 13.8 8.8

XRD R*

12 30

1.33 0.85 4.33

W*

3.5 7.5 3.5

XPS Binding energy O 1s V2p(1) V2p(2)

v(1)/

531.0 531.1 530.6

0 0.09 0.37

516.3 516.1

517.8 517.6 517.6

V(2) .

* R - the ratio of I(001)/I(~10) indexes intensity, W - WHPM - width at the half of the (001) reflection of the latter over the surface of V205 (see Table 1) while an average size of the particles remains almost unchanged [13]. On the contrary, dry milling of V205 by chaotic destruction of the crystals produced smaller particles, which is reflected by the increase of the XRD peaks width (Table 1). The degree of surface reduction of vanadium pentoxide (given by the content ratio V(1)/V(2) in Table 1) was much higher after the treatment in ethanol as compared to dry milling. No special features of the synthesis of VPO-D from VzOs-D were observed as compared to traditional VPO-R compound synthesis, both lasting 12-16 h. But synthesis of VPO-E precursor using VzOs-E proceeded at a much higher rate and in the presence of organic reducing agent was completed in 1h (sample VPO-E 1). Another preparation using V2Os-E was carried out without reducing agent and with smaller amount of the solvent (sample VPO-E2). In this case the time needed for the full formation of the precursor phase was about 2.5-3.0 h. Some properties of the prepared VPO precursors are listed in Table 2. Table 2 Properties of VPO precursors prepared from V2Q5 treated mechonochemically Sample

SsA

DTA

XRD*

Tendo Texo I0.570/I0.329 VPO-R VPO-D VPO-E1 VPO-E2

mZ/g

~

~

20.2 19.5 4.6 12.1

448 435 468 448

500 485 505 495

/I0.293 75/46/100 100/40/90 73/45/100 100/35/76

XPS _Binding energy, eV O 1S V 2p P 2P

(P/V)s

532.3 532.2 532.4 532.3

2.08 2.00 3.30 2.05

517.5 517.4 517.5 517.5

133.6 133.6 133.7 133.7

*Intensity ratio for reflections at d = 0.570, 0.329 and 0.293 nm It follows from the data in Tab.2, that reduction of the V205 particles size results in VPOD precursor in increased intensity of the 0.570 nm peak attributed to the exposure of (001) plane containing the vanadyl groups. Also some decrease of the temperatures, at which the amorphous (Tendo)and the crystal (Texo) phases are formed in the course of vanadyl pyrophosphate preparation, was observed (see [10, 15] for details on phase transformations). The change of V205 texture during its mechanochemical treatment in ethanol leads to the synthesis of VPO-E 1 precursor with low specific surface area and unchanged texture. It can be

340 assumed that freshly formed microcrystalls of the precursor, during the fast synthesis, rapidly grow into agglomerates. This sample shows also an increased temperature of amorphization and of crystallization during the formation of vanadyl pyrophosphate and the increased surface P/V ratio. Modification of the conditions of synthesis, consisting in some deceleration of the process, allows the preparation of the sample (VPO-E2) with larger specific surface area. Moreover, it has favourable morphology with high exposure of (001) crystallographic plane. Table 3 shows the catalytic properties of these samples. One can see from the data, that all catalysts synthesized on the basis of the mechanochemically treated V205 show an increased selectivity towards maleic anhydride and higher specific rate of n-butane and npentane oxidation as compared to those obtained in traditional synthesis. The best effect in the improvement of selectivity can be reached by increase of the relative exposure of (001) plane at the VOHPO4.0.5.H20 surface which is known to be transformed into (200) plane of (VO)2PzO7. The low paraffins conversion over VPO-E samples at the given reaction conditions can be directly connected with their low specific surface area. The comparison made between samples VPO-E1 and VPO-E2 shows that the precursor synthesis using VzOs-E needs to be optimized in order to improve the catalytic performance. Nevertheless, the results clearly demonstrate that mechanochemistry is obviously a promising method for the pretreatment of initial reagents in order to synthesize efficient VPO catalysts of paraffins oxidation. Table 3 Properties of VPO catalysts prepared using mechanochemically treated V205 in reactions of paraffins oxidation. Sample

n-Butane oxidation X, % SMA,% W. 104

n-Pentane oxidation X,% SMA,% SPA,% W. 104

VPO-R VPO-D VPO-E1 VPO-E2

73 79 25 63

52 59 17 43

58.5 68.5 60.0 69.0

1.13 1.45 1.60 1.59

21 35 33 30

12 10 4 6

0.92 1.17 1.32 1.29

Note: X - paraffin conversion, SMA, SpA- selectivity to maleic and phthalic anhydride, respectively, W - specific rate of the paraffin oxidation, mol/h m 2. For n-butane T = 425 ~ GHSV = 3000 h -1 and for n-pentane T = 320 ~ GHSV = 1800 h -~ 2.1.2. Mechanochemical modification of the VPO precursor Previously we showed [ 13,16] that the nature of dispersant had a quite remarkable effect on the properties of VPBiO precursor which was subjected to a short-time mechanochemical treatment (up to 10 rain.) and the best result was obtained when ethanol was used. Here we will describe the influence of the time of treatment on the properties of VPO precursor (P/V = 1.07). As can be seen from Table 4, the longer is the mechanochemical treatment (up to 20 min) the higher are both specific surface area of the precursor and relative intensity of its (001) crystallographic plane reflection. The latter observation can be explained by anisotropic plastic deformation of the crystals arising from their layered structure. The observed increase of the surface P/V ratio is in a good agreement with the mechanism of the VOHPO4.0.5H20 phase transformation [ 15].

341 Table 4. Properties of VPO precursor after its mechanochemical treatment in ethanol Sample Time SsA treatm.,

min

mZ/g

XRD

I0.570/ I0.293

XPS Binding energy, eV O 1s V 2p P2p

n-butane oxidation* (P/V)s

SSA1

X

S

VPO 6.0 74/100 531.7 517.5 133.9 1.43 9.4 62 61 VPO10 10 8.8 95/100 531.6 517.4 133.8 1.62 12.3 68 65 VPO20 20 14.2 100/78 531.8 517.4 133.7 1.80 17.2 73 70 VPO30 30 8.0 ** 532.1 517.3 133.7 1.92 8.5 77 74 VPO60 60 6.4 *** 532.2 517.3 133.7 1.84 6.5 70 68 *T=440~ GHSV=3200h 1, S S A 1 - specific surface area after catalysis (m2/g), X - n-butane conversion (%), S- selectivity to maleic anhydride (%); the samples VPO30 and VPO60 did not need activation prior to catalysis, **Amorphization of the sample, very weak peaks at d = 0.328, 0.305 and 0.285 nm ***All reflections ofvanadyl pyrophosphate are present, the most intense one is at d = 0.313 nm Continuation of the mechanochemical treatment leads to amorphization of the sample followed by the formation of vanadyl pyrophosphate phase. It should be however noted that even after 60 min. of the treatment this compound is not well crystallized and the specific surface area of the final catalyst is much smaller than that of a sample obtained by in situ activation of a precursor after 20 min of its mechanochemical treatment. It follows from the results presented in Table 4, that the samples obtained by mechanochemical treatment of the precursor become more active in the reaction of n-butane partial oxidation so that the hydrocarbon conversion and selectivity to maleic aL!hydride increase. The sample converted into vanadyl pyrophosphate by means of the mechanochemical treatment turned out to be more efficient than that activated with the reaction mixture. The most interesting is the sample after 30 min. treatment which is "half-activated" and consists of the amorphous phase. The active component forming directly in the catalytic mixture without long activation procedure gives rise to the most active and selective catalyst for n-butane oxidation.

2.1.3. Mechanochemical promotion of VPBiO precursor Recently [ 13, 16] we have shown the possibility of efficient promotion of VPO precursor with bismuth compounds. The present paper reports new results on promotion of VPBiO precursor with lanthanum compounds (previously a similar catalyst was shown to be active in tetrahydrofuran formation [17]). Table 5 compares the properties of traditionally prepared VPBiLaO sample (by simultaneous introduction of bismuth and lanthanum additives in the course of the synthesis of VPO precursor) with that (VPBiO-La-M) produced by the mechanochemical treatment of VPBiO precursor and lanthanum oxide powders. The treatment in the latter case was carried out for 10 min. in ethanol medium. It can be seen that introduction of lanthanum by both methods leads to an increase of the catalytic activity. At the same time, its introduction in the course of the synthesis of VPBiO precursor causes a decrease of the selectivity to partial oxidation products in both investigated reactions: oxidation of butane and propane. This negative effect on selectivity was also observed when the catalyst was prepared by means of mechanochemistry but to a much lesser degree. As a result, the latter catalysts were more efficient and produced higher yield of the desired product (see Table 5 data). It can be noted that for the traditional sample a higher value of V 2p-electrons binding energy is observed suggesting that the oxidation degree of vanadium can be in this case higher than in the traditional samples.

342 Table 5 Properties of VPO precursor with additives bismuth and lanthanum. Sample*

VPBiO VPBiLaO3 VPBiLaO5 VPBiO-La3M VPBiO-La5M

XPS Binding energy, eV** P/V Bi/V La/V O 1 s V 2p P 2p

Oxidation n-Butane*** Propane**** X SMA Y X SAg Y

531.5 531.8 531.7 531.6 531.5

48 56 61 55 58

517.4 517.9 517.9 517.5 517.5

133.7 1.58 0.12 134.0 1.61 0.17 0.027 133.9 1.88 0.14 0.063 133.8 1.73 0.09 0.023 133.6 1.95 0.08 0.058

69 61 54 68 66

33 34 33 37 38

40 49 60 47 59

32 21 17 30 28

13 10 10 14 16

*Number in the sample name represents the atomic ratio (La/V).100 ** B.E. Bi 4f = 159.9 160.2 eV, La 3d = 836.5-836.7 eV *** T = 420 ~ GHSV = 3000 h "l, X - n-butane conversion (%), SMA - selectivity to maleic anhydride (%), Y - yield of maleic anhydride (mol. %); **** T = 435 ~ GHSV = 1500 h -1, X - propane conversion, %, SAA - selectivity to acrylic acid, %, Y - yield of acrylic acid, mol. %

2.1.4. Mechanochemical modification of vanadyl pyrophosphate Vanadyl pyrophosphate was prepared by heating the VPO precursor in an inert gas flow at 550 ~ for 24 h. Its characteristics are listed in Table 6. One can see from TSM pictures (Figure 1a) that the catalyst prepared in this way is composed of quite large aggregates including crystals of geometrically-regular shape. Mechanochemical treatment disintegrates them to produce much smaller particles of different shape (Figures 1 b,c,d). It is noteworthy, that in the case of the treatment of (VO)2P207, there is no dependence of the morphology on the dispersant nature at variance with the case of the VPO precursor treated similarly [13]. An increase of the intensity of the reflection at d = 0.387 nm corresponding to (200) plane of vanadyl pyrophosphate was observed by XRD to be the only structural change occurring at the treatment in ethanol. As a result, the catalyst after treatment shows an increase of both activity and selectivity in n-butane partial oxidation. The specific rate of the hydrocarbon conversion decreases after the treatment which is believed to be due to non-proportional growth of the number of active centers and the specific surface area. Table 6 The properties of vanadyl pyrophosphate after mechanochemical treatment Treatment

SSA

Solvent Time, min. m2/g Water Ethanol

10 10 10

* T = 440 ~

4.6 9.2 5.4 7.1

XRD

I0.387/ I0.313 82/100 95/100 92/100 100/73

XPS Binding energy, eV O lS V 2p P 2p 531.5 531.6 531.8 .

517.5 517.4 517.5 . .

133.8 133.6 133.8 .

n-Butane oxidation* SMA, W-10 4 %

P/V

X, %

1.22 1.34 1.42

68 76 69 77

GHSV = 1500 h l , W - rate of n-butane oxidation, mol/h.m 2

60 58 62 69

1.44 0.86 1.19 1.08

343

a.

b.

Figure 1. Transmission Scanning Microscope pictures of vanadyl phyrophosphate, a) initial, and after mechanochemical treatment: b) in water, c) in ethanol and d) dry milling. 2.2. Barothermal synthesis 2.2.1. Barothermal synthesis of VPO catalysts Soon after publications of J.Johnson and A.Jacobson [18,19] hydrothermal synthesis has begun to be applied to different vanadium phosphates synthesis. In the present work an attempt has been undertaken to use organothermal (with n-butanol addition) synthesis and that without any solvent. It has been established that in hydrothermal synthesis starting from V205 and H3PO4 it is possible "depending on the synthesis temperature and duration" to obtain VOPO.2H20 and VOPO4.H20. When VO2 is used the mixture of vanadium (IV) and (V) compounds such as 13VOPO4 and VOHPO4.0.5H20 can be formed. Their catalytic performance in n-butane and butene-2 oxidation (better for samples HS-1) is worse than that of the catalysts prepared by other methods.

344 Organothermal synthesis with the use V205 leads to the formation of [3-VOPO4 (OS-1). The latter compound shows low activity in paraffin oxidation, but is a quite efficient catalyst for olefin oxidation (Table 7). When VO2 was used for organothermal synthesis, an unknown compound was obtained at low temperature and/or short time of the synthesis. Continuation of the synthesis led to the formation of VOHPO4.0.5H20 (OS-2). In the synthesis using V205 in solvent-free conditions, 13-VOPO4 was found to be the only product (AS-1) but quite a high temperature and long time were needed for the reaction to be completed. In the case of the use of VO2, a new compound was formed. Table 7 Properties of the barothermally synthesized samples Butene-2 oxidation*** SMA,% W 104

n-Butane oxidation** W 104

Sample*

X, % SMA,% TS-1 TS-2 HS-1 OS-1 OS-2 AS-1

68 63 35 75 28

60 60 20 64 21

X, o~

1.44 1.11 1.02 1.76 0.48

76 78 74 71 70

71 55 69 51 72

3.22 2.86 3.31 3.14 3.06

* T S - 1 , -2 - catalysts prepared following traditional methods for n-butane and butene-2 oxidation, respectively **T = 440 ~ GHSV = 1500 h -l', ***T = 380 ~ GHSV = 3600 h -1', X hydrocarbon conversion, SMA - selectivity to maleic anhydride, W - rate of hydrocarbon oxidation, mol/h.m 2 2.2.2. Barothermal

modification

of VPBiO

precursor

The VPBiO precursor treatment was performed with n-butanol and phosphoric acid vapours as the perspective media for the treatment of VPO [ 12]. The obtained results, sortie of which are listed in Table 8, show that the higher is the temperature and the longer is the treat ment the lower becomes the specific surface area. At the same time, it should be noted that Table 8 An influence of VPBiO precursor barothermal treatment on its properties Sample*

VPBiO VPBiOnbl VPBiOnb2 VPBiOnb3 VPBiOpal VPBiOpa2 VPBiOpa3

Treatment Time T, h ~ -

6 10 6 14 10 6

-

250 250 300 200 250 300

SSA

XRD I0570/ m Z / g I0293 12.5 11.2 10.0 9.3 12.0 10.2 8.7

100/98 100/78 100/76 100/39 100/76 100/49 100/45

XPS Binding energy, eV O ls V 2p P-2p Bi 4f 531.5 531.7 531.6 531.4 531.1 531.4 531.3

517.4 517.5 517.6 517.1 517.1 517.3 517.3

133.6 133.9 134.0 133.5 133.4 133.5 133.4

159.9 160.2 160.1 159.7 159.7 159.8 159.7

*nb, pa- treatment with n-butanol and phosphoric acid, respectively

(P/V)s (Bi/V)s 1.47 1.87 2.08 2.07 2.23 2.21 1.98

0.08 0.12 0.10 0.10 0.10 0.11 0.11

345 such treatment leads to some change of the structure of VPBiO precursor as it can be seen from XRD results. An increase of the relative intensity of the reflection at d = 0.570 nm attributed to crystallographic plane having vanadyl groups is namely observed. Enrichment with phosphorus is found at the surface of the particles after treatment. Similar effect could be expected at the treatment with phosphoric acid but looks unusual in the case of the treatment with alcohol. The catalytic properties of the samples are presented in Table 9. The data indicate that the barothermal treatment favours an increase of the selectivity in paraffins oxidation. Moreover, the treatment in n-butanol also leads to the growth of catalytic activity of the samples. In the case of the treatment with phosphoric acid vapours, the catalytic activity remains almost unchanged, due to effect of water steam as described in [ 12]. Table 9 Catalytic properties VPBiO precursor after barothermal treatment Sample

SSA 1

m2/g VPBiO VPBinbl VPBinb2 VPBinb3 VPBipal VPBipa2 VPBipa3

14.6 12.0 9.8 9.0 11.6 9.7 8.5

n-Butane oxidation* X, % SMA,~ 49 50 57 57 48 50 45

68 72 75 77 73 81 76

n-Pentane oxidation** X,% SMA,% SPhA,% 67 -

32 -

17 -

61 65 65

39 42 36

14 18 21

SSA1 - specific surface area after catalysis * T = 400 ~ GHSV = 2400 h -1 *** T = 420 ~ GHSV = 1500 h -1

Propane oxidation*** X, o~ SAA,o~ 26 28 29 31

48 55 59 58

-

GHSV = 3000 h -1 **T = 350 ~

Concluding, it should be emphasized that mechanochemical and barothermal methods have been shown to be advantageous as alternative technologies for preparation and modification of VPO catalysts for partial oxidation of saturated hydrocarbons. ACKNOWLEDGMENT This study was supported in a part by ISF (Grants UBI000 and UBI200) and in a part by Scientific Research Committee (Poland) Grant No 3T09A 08010. The authors thank Prof. V.G.Iljin, Dr. G.A.Komashko and V.E.Yaremenko for assistance in some experimental work.

REFERENCES 1. 2. 3.

G. Centi (editor), Vanadyl Pyrophosphate Catalysts, Catal. Today, 16 (1993) 1-147. G.J. Hutchings, Appl. Catal., 72 (1991) 1. F. Cavani and F. Trifiro, Preparation of Catalysts VI, Stud. Surf. Sci. Catal., Elsevier, Amsterdam, 91 (1995) 1.

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V.A. Zazhigalov, V.M. Belousov, G.A. Komashko, A.I. Pyatnitskaya, Yu.N. Merkureva, A.L. Poznyakevich, J. Stoch and J. Haber, Proc. 9th Int. Congr. Catal., Chem. Inst. of Canada, Ottawa, 4 (1988) 1546. 5. P.F. Miquel and J.L. Katz, Preparation of Catalysts VI, Stud. Surf. Sci. Catal., Elsevier, Amsterdam, 91 (1995) 207. 6. P.M. Michalakos, H.E. Bellis, P. Brusky, H.H. Kung, H.Q. Li, W.R. Moser, W. Partenheimer and L.C. Satek, Ind. Eng. Chem. Res., 34 (1995) 1994. 7. V.V. Guliants, J.B. Benziger and S. Sundaresan, J.Catal., 156 (1995) 298. 8. P.F. Miquel, E. Bordes and J.L. Katz, J.Solid State Chem, 124 (1996) 95. 9. L. Bogutskaya, V. Zazhigalov, M. Misono and T. Okuhara, Japan-FSU Catal. Seminar'94, Catal. Sci and Techn. 21 Century Life, Tsukuba, Japan, (1994) 202. 10. V.A. Zazhigalov, J. Haber, J. Stoch, L.V. Bogutskaya and I.V. Bacherikova, Appl. Catal. A, 135 (1996) 155. 11. V.A. Zazhigalov, J. Haber, J. Stoch, A.I. Pyatnitskaya, G.A. Komashko and V.M. Belousov, Appl. Catal. A, 96 (1993) 135. 12. V.A. Zazhigalov, I.V. Bacherikova, V.E. Yaremenko, I.M. Astrelin and J. Stoch, Teoret. Eksperim. Chem., 31 (1995) 206. 13. V.A. Zazhigalov, J. Haber, J. Stoch, L.V. Bogutskaya and I.V. Bacherikova, 1 l th Int. Congr. on Catal. - 40th Anniversary, Stud. Surf. Sci. Catal., Elsevier, Amsterdam, 101B (1996) 1039. 14. V.A. Zazhigalov, J. Haber, J. Stoch, A.I. Kharlamov, L.V. Bogutskaya, I.V. Bacherikova and A. Kowal, Solid State Ionics (in press). 15. C.C. Torardi, Z.G. Li and H.S. Horowitz, J.Solid State Chem., 119 (1995) 349. 16. J. Haber, V.A. Zazhigalov, J. Stoch, L.V. Bogutskaya and I.V. Bacherikova, Catal. Today (in press). 17. V.A. Zazhigalov, J. Haber, J. Stoch, G.A. Komashko, Ai. Pyatnitskaya and I.V. Bacherikova, New Develop. Select. Oxid. II, Stud. Surf. Sci. Catal., Elsevier, Amsterdam, 82 (1994) 265. 18. J.W. Johnson, A. Jacobson, J.F. Brody and S.M. Rich, Inorg. Chem., 21 (1982) 3820. J.W. Johnson and A. Jacobson, Angew. Chem., 95 (1983) 442. D

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

Active species and working mechanism of silica supported MoO3 and catalysts in the selective oxidation of light alkanes

347

V205

A. Parmaliana% F. Arena% F. Frusteri b, G. Martra c, S. Coluccia r and V. Sokolovskii d aDipartimento di Chimica Industriale, Universitb, di Messina, Salita Sperone 29, 98166 S. Agata (Messina), Italy blstituto CNR-TAE, Salita S. Lucia 39, 98126 S. Lucia (Messina), Italy CDipartimento di Chimica Inorganica, Chimica Fisica e Chimica dei Materiali, Universit/l di Torino, Via P. Giuria 7, 10125 Torino, Italy aDepartment of Chemistry, University ofWitwatersrand, Johannesburgh, P.O. Box 106, South Africa The catalytic performance of a series of silica supported (2-7 wt%) MoO3 and (2-20 wt%) V205 catalysts has been comparatively evaluated in both the partial oxidation of methane to formaldehyde (MPO) and the oxidative dehydrogenation of propane to propylene (POD) in the range 500-800~ and 500-525 ~ respectively. V205 acts as a promoter of the reactivity of the SiO2 for both MPO and POD, while MoO3 plays a generally negative effect on the catalytic functionality of the SiO2 surface acting however as a promoter in the MPO at T>650~ A direct relationship between the density of reduced sites of low-medium loaded silica based oxide catalysts and reaction rate in both MPO and POD strongly suggests the occurrence of a concerted reaction mechanism involving the activation of gas-phase 02 on the reduced sites of the catalyst surface. The nature of the active sites has been defined on the basis of a complete characterization of the surface and redox features of MoO3/SiOz and V205/SIO2catalysts. 1. INTRODUCTION The disclosure of the mechanism of a catalytic reaction leads to the identification of the active sites being then the basis for the design of more effective catalytic systems. Generally, a mechanistic model is adequate for describing the pathway of a class of reactions on a class of catalysts, however this rule is not completely valid for the partial oxidation of light alkanes on bulk and/or supported oxide catalysts [ 1]. Since the excellent performance of supported MoO3 and V205 based catalysts in the selective oxidation of light (C~-C3) alkanes, during the last decade a considerable research effort has been directed to ascertain the working mechanism of such oxide systems as well as the nature of the active sites and the origin of the oxygen involved in the formation of reaction products. However, no definitive conclusions have been still provided about these issues and therefore a much deeper investigation of the reaction dynamic mechanism and the correlation with the nature of the surface is necessary [2]. Several factors, such as the nature of the support, the oxide loading and the reaction conditions control the formation, the coordination and the stabilization as well as the catalytic action of the

348 various surface oxide species. On this account, we have evaluated the catalytic behavior of silica supported MoO3 and V205 systems in both the selective oxidation of methane to formaldehyde (MPO) [3,4] and the oxidative dehydrogenation of propane to propylene (POD) [5] disclosing that V205 acts as a promoter of the reactivity of the SiO2 surface while the action of MoO3 strictly depends upon the kind of the silica support [3-5]. The aim of this paper is to provide a correlation between the catalytic pattern of differently loaded silica supported MoO3 and V205 catalysts in MPO and POD reactions with their surface and redox features in order to highlight the nature of the active surface species in the selective oxidation of light alkanes. 2. EXPERIMENTAL 2.1. Catalysts

Differently loaded (2-7 wt%) MoO3/SiO2 and (2-50 wt%) V205/SIO2 catalysts were prepared by incipient wetness impregnation of a "precipitated" silica (PS) support (Si 4-5P Grade, Akzo Product; S.A.BET, 400 m2.g~) according to the procedure described elsewhere [3]. The list of the catalysts along with their composition and BET surface area values are reported in Table 1. Table 1 List of catalysts Code Chemical composition

(wt %) PS VPS 2 VPS 5 VPS 10 VPS 20 VPS 50 V205 MPS 2 NIPS 4 NIPS 7 MoO3

SiO2 2.0% V205/SIO2 5.3% V205/5iO2 10.1% V205/5i02 20.8% V2OJSi02 50.8% V2OJSi02 V205 2.0% MoO3/SiO2 4% MoOa/SiO2 7% MoO3/SiO2 MoO3

S.A.BET

(m~.~-') 400 260 230 200 190 160 5 300 190 75 2

2.2. Catalytic measurements

Catalytic measurements in MPO were performed by Temperature Programmed Reaction (TPR) tests [4] using a conventional flow apparatus and a linear quartz microreactor connected on line with a Thermolab (Fisons Instruments) Quadrupole Mass Spectrometer (QMS) for continuous scanning of the reaction stream. TPR tests were run in the T range 400-800~ by using 0.05 g of catalyst, a heating rate (13) of 10~ "1 and a reaction mixture He/CH4/O2 in the molar ratio 7:2:1 flowing at 50 STP cma.min-1. Catalytic measurements in POD have been performed at atmospheric pressure in the range 500-525~ using 0.25 g of catalyst sample diluted with same sized SiC (1/5, l/vol) and a reaction mixture in the molar composition CaH8:O2:N2:He=2:l:l:6 flowing at the rate of 100 STP cmS-min'l[5]. All the tests have been carried out at GHSV of 1,700 h"1 (STP m3C3H89m'aeat-h'l).

349

2.3. Catalyst characterization Temperature Programmed Reduction (H2-TPR) measurements were performed using a linear quartz gradientless microreactor and a 6% H2/Ar mixture flowing at 60 STP cm3 rain~ according to the procedure described elsewhere [6]. High Temperature Oxygen Chemisorption (HTOC) measurements were performed in a pulse mode according to the procedure described in detail elsewhere [6]. Reaction Temperature Oxygen Chemisorption (RTOC) measurements in the range 500650~ were performed in a pulse mode after treatment of the samples in the C3HffO2/He or CH4/Oz/He reaction mixture. 02 pulses (Vp,l,o, 4-10.8 mol 02) were injected into the carrier gas until saturation of the sample was attained, the density of reduced sites (p, 1016 sites.gc,t~) being calculated by assuming a chemisorption stoichiometry 02:reduced site = 1:2 [3,5]. Diffuse Reflectance UV-Vis DR UV-Vis spectra of differently loaded V2OflSiO2 samples, calcined in situ in 02 at 600~ were obtained by a Perkin Elmer Lambda 19 spectrophotometer, equipped with an integrating sphere. 3. RESULTS and DISCUSSION

3.1. Catalytic activity Methane partial oxidation (MPO). The catalytic activity of differently loaded MPS and VPS samples in the range 500-800~ expressed in terms of normalized specific surface activity, NSSA (NSSA=SS&/SSAps, where SSAi and SSAps are the specific surface activity of the catalyst i and bare PS support, respectively), is compared in Figure 1A. 1.0

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350 The HCHO selectivity as a function of CH4 conversion in the range 500-800~ for MPS and VPS catalysts is shown in Figure lB. For MPS 2 and MPS 4 catalysts, the trend of HCHO selectivity with CH4 conversion (Fig. 1B) is similar to that of the unloaded PS, whereas for the MPS 7 sample a significant improvement in HCHO selectivity at conversion levels lower than 3% is observed. By contrast, at the same level of CH4 conversion, a progressive decline in HCHO selectivity with V205 loading occurs on VPS catalysts. Medium loaded MPS 7 and highly loaded VPS 20 samples are the most and the least selective systems respectively. Oxidative dehydrogenation of propane (POD). The predominant products in the POD over silica based oxide catalysts in the range 500-525~ were propylene and carbon oxides. Ethylene and acetaldehyde along with a considerable amount of acrolein and traces of propionaldehdye are formed on PS [5]. The addition of MoO3 and V205 to the SiO2 support implies a higher selectivity to propylene and correspondingly a lower production of COx, a slight cracking activity and a significant decrease in the amount of oxygenates. Table 2 shows a detailed comparison of the activity of bare and differently loaded MPS and VPS catalysts in terms of propane conversion, propylene selectivity, reaction rate and C3I-I6 productivity values. The catalytic functionality of the SiO2 surface is strongly promoted by the addition of V205, while it is depressed by MOO3, the extent of this effect rising with the MoO3 loading [9]. Table 2 Activity of bare and promoted silica catalysts in the oxidative dehydro[genation of propane Catalyst T C3Hs conv. C3I-I6 s e l . Reaction rate STYc3H6 (%) (].tmol.gc.tl.s "1) (g.kgc.t'l.h "1) (~ (%) PS 5OO 0.9 37 0.49 27.5 525 1.9 37 1.04 57.7 MPS 2 500 0.8 62 0.44 41.4 525 1.3 58 0.71 62.5 MPS 4 5OO 0.7 69 0.39 40.8 525 1.4 59 0.75 67.1 MPS 7 500 0.2 80 0.09 6.5 0.4 525 85 0.18 11.6 VPS 5 500 7.8 60 4.25 386.4 525 13.3 55 7.25 602.8 VPS 10 5O0 9.8 51 5.34 413.0 525 177 41 9.65 600.0 VPS 20 500 7.8 27 4.25 174.0

3.2. Redox properties of MoO3/SiO2 and V2Os/SiO2 catalysts under reaction conditions In previous papers we disclosed a direct relationship between catalytic activity in MPO of low-medium loaded silica based oxide catalysts and oxygen uptake under steady state reaction conditions pointing out that such property governs the catalytic behavior of MPO catalysts [3,7,8]. Then, in order to find out whether such a relationship is also valid for POD reaction, the density of reduced sites (p) of various silica based MoO3 and V205 catalysts in both MPO and POD has been evaluated and related with their catalytic activity. The direct relationships between reaction rate in MPO and POD and p, shown in Figure 2A, well account for the opposite effects exerted by MoO3 and V205 on the activity of the bare PS carrier in both MPO [3,4,6-8] and POD [5]. Indeed, V205 effectively promotes the activity of bare PS, and allows

351 for the stabilization of a higher density of reduced sites owing to its easier "reducibility" [6] whereas MOO3, being essentially unreducible under reaction conditions [6], depresses the activity of the PS carrier because of a negative physical effect linked to the partial coverage of the silica surface own active sites [3,6,8]. However, though the above relationships (Fig. 2A) explain the origin of the catalytic action of low-medium loaded (___5wt%) MPS and VPS systems, they do not account for the activity of highly loaded (>__10wt%) VPS catalysts in MPO. 15

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Indeed, the activity of VPS catalysts reaches the maximum for VPS 5 sample, while the density of reduced sites increases steadily up to a loading of 20% (VPS 20) resulting in the peculiar volcano-shaped relationship between P and reaction rate shown in Figure 2B. This singular behavior is to be related with the capability of highly loaded V2Os/SiOz systems to "loose" constitutional oxygen [6] under reaction conditions which, however, is ineffective towards the selective oxidation of light alkanes [5,8,9]. 3.3. Surface structures of MoO3/SiO2 and V2Os/SiO2 catalysts

The H2-TPR profiles of differently loaded MPS and VPS catalysts in the range 200-1200~ compared with those of bulk MoO3 and V205 respectively, are shown in Figure 3A and B, while the values of oxygen uptake (HTOC) and oxide dispersion (O/Me) are listed in Table 3. A wide and convoluted band of H2 consumption starting (To, ,~d) at T ranging between 435 (MPS 7) and 486~ (MPS 2) and spanning a T range of 500-600~ accounts for the stoichiometric reduction of MoO3 to Mo Oin all MPS catalysts. The spectrum of the low loaded MPS 2 catalyst features a low rate of H2 consumption up to ca. 800~ thereafter the reduction rate increases sharply giving rise to a main reduction peak with maximum at 934~ (Fig. 3A, a). The increase in the MoO3 loading from 2 to 4 wt% (MPS 4) causes a marked shift of To,red to lower T (436~ and a concomitant enhancement of the reduction kinetics at lower T (<800~ giving rise to a convoluted reduction profile with two unresolved peaks with maxima at 588 and 765~ respectively (Fig. 3A, b).

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A further increase in the MoO3 loading strongly enhances the reduction rate of the system at T<800~ (Fig. 3A, c), even if the To,~d value (435~ of the MPS 7 catalyst is equal to that of the MPS 4 sample (436~ Table 3 HTOC characterization data of MPS and VPS catalysts Catalyst 02 uptake (l.tmol.g.t 1) PS 0.5 MPS 2 6.1 MPS 4 66.4 MPS 7 111.5 MoO3 20.7 VPS 2 68.0 VPS 5 186.9 VPS 10 262.1 VPS 20 328.3 VPS 50 182.3 V20s 42.8

O/Me (%) 8.8 47.8 45.9 0.6 61.9 64.5 46.6 28.7 6.5 0.8

The reduction pattern consists of two sharp peaks with resolved maxima at 574 and 703~ respectively, along with a shoulder of Hz consumption in the range 800-1050~ Because of a poor dispersion (Table 3), the reduction of bulk MoO3 (Fig. 3A, d) starts at T considerably higher (To,,od, 532~ than those found for supported MPS systems [10], resulting in two

353 overlapped peaks, with maxima at 763 and 840~ which account for the step-wise (MoVL--~MorV---~Mo~ reduction of MoO3 to Mo ~ The 02 uptake of MPS catalysts increases monotonically with the loading from 6.1 (MPS 2) to 111.5 i.tmol-g1 (MPS 7). The oxide dispersion (O/Me) results very low for the MPS 2 catalyst (8.8%), while it suddenly increases for MPS 4 sample (47.8%) keeping unchanged (45.9%) at higher MoO3 loading (MPS 7). The broad band of H2 consumption featuring the reduction pattern of MPS catalysts in comparison to bulk MoO3 is diagnostic of a strong metal oxide-support interaction which markedly depresses the reduction of MoO3 promoter [10]. By using a least-square fitting program, the spectra of MPS catalysts have been resolved into the contribution of three similar discrete Gaussian-shaped peaks, assuming that the first two peaks (i.e., MI and M2) monitor the step-wise reduction (MoVL-~MorV-->Mo~ of "MOO3 crystallites" (Mc) while the third one (M3) refers to the one-step reduction (MoVI---~Mo~ of"Isolated molybdates" species (Ira) [10]. The fitting parameters of the deconvoluted TPR profiles of MPS 2, MPS 4 and MPS 7 samples (Figures not reported here), expressed as peak maximum position (M~), full width at half maximum (FWHM0, percentage peak area (Ai, %) and concentration of Mc and Im species are reported in Table 4. Such data signal that the intensity of the first two peaks, (M0 and (ME) respectively, rises monotonically with MoO3 loading; whilst the highest temperature peak (M3), monitoring the reduction oflm species, follows an opposite decreasing trend [ 10]. Table 4 Fitting parameters of TPR spectra and percentage of Mc and Im species in differently loaded MPS catalysts Sample Mc Im Mc Im M3 FWHM3 A3 Concentration M1 FWHMI A1 M2 FWHM2 A2

(oc) (oc) (%) (~

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On the whole, the surface composition of the differently loaded MPS catalysts provides evidence of the influence of the loading on the surface structures of MoO3/SiO2 catalysts, allowing also to explain the unusual trend of chemisorption data (Table 3). Indeed, the prevailing stabilization of"hardly reducible" Im species accounts for the small oxygen uptake and quite low dispersion value in MPS 2 sample (Table 3). At MoO3 loadings higher than 2 wt%, the increasing formation of Mc (Table 4) enhances the reducibility of the system as well as the chemisorption capability and dispersion values (Table 3). However, also for the MPS 4 sample the dispersion degree calculated from HTOC data (Table 3) would be underestimated, being really much higher than that of MPS 7 catalyst [10]. On the basis of the above considerations, it can be inferred that three types of "surface sites" contribute to the reactivity of MoO3/SiO2 catalysts in the selective oxidation of light alkanes: the siloxane bridges of the silica surface [6,7]; the Mo-O-Mo bridging functionalities and Mo=O terminal bonds of Mc [6]. Then, the catalytic pattern of MoO3/SiO2 system in MPO and POD depends on both concentration and activity of various surface sites. The H2-TPR profiles of differently loaded VPS catalysts in the range 200-1100~ shown in Figure 3B, account for the stoichiometric reduction of V 5+ to V 3+ in both VPS and bulk V2Os systems. The H2-TPR pattern of the low VPS 2 catalyst (Fig. 3B, a) entails a very sharp

354 reduction peak with the maximum (TM0 at 55 I~ slightly asymmetric on the high temperature side due to the presence of a shoulder of H2 consumption zeroing at T~850~ The To,~d is equal to 360~ being the lowest found in the series. An increase in the V205 loading to 5 wt% (VPS 5) does not substantially affect the reduction pattern of the system neither in terms of To,red (364~ nor in peak shape (Fig. 4B, b) even if the TM~ is slightly displaced to higher T (561~ More evident changes in the reduction pattern of the V2OJSiO2 system (Fig. 3B, c) occur at higher V205 loading (VPS 10), since the To,roavalue further shifts to higher T (372~ while a broadening Of TM~ peak along with a further rise Of TM~ value (571~ are recorded. Moreover, two new smaller peaks, bearing resolved maxima at 638 (TM2) and 730~ (TM3), are observable (Fig. 3B, d). At a loading of 20 wt% (VPS 20), the TM~ peak decreases sharply in intensity assuming an asymmetric shape, while its maximum further shifts to higher T (TM~=584~ TM2 and TM3 peaks rise in intensity, the former becoming the predominant one. Bulk V205 (Fig. 3B, f) displays two very sharp, partially overlapping, reduction peaks centered at 649~ (TM2) and 716~ (TM3) respectively and a third broader peak at higher T with maximum at 939~ (TM4). The relative intensity of such three peaks accounts for the following sequential reduction path" V2Os~VtO~3~VO2~V203. The oxygen uptake of VPS catalysts increases monotonically with the loading attaining the maximum value (~328 l.tmol.g"~) for the VPS 20 sample. The oxide dispersion reaches the maximum value (60-62%) for low-medium loaded (2-5 wt%) catalysts, thereafter at loadings higher than 5% it decreases to 46.6 and 28.7% for VPS 10 and VPS 20 samples respectively. The energy of the oxygen~vanadium charge-transfer absorption band, which is correlated with the minimum diffuse reflectance, is strongly influenced by the number of ligands surrounding the central vanadium ion [ 11] and thus it provides insights into the influence of the loading on the coordination of V S+ ions in VPS samples. DR-UV-Vis spectra of differently loaded VPS catalysts, shown in Figure 4, signal that the loading markedly affects the structure of supported V s+ ions.

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355 keeping unchanged the Td coordination of V ions 6 [ 11]. The DR-UV-Vis spectral features of the VPS 5 sample, showing a predominance of the band at ca. 34,000 cm~, suggest that the concentration of the oligomeric structures has increased with loading, even if monomeric forms are still present (Fig. 4b). The spectral pattern of VPS 10 system is characteristic of more complex structures with a higher degree of"nuclearity" (Fig. 4c) though the Td coordination of V ions is still preserved [ 11]. Finally, the bands at ca. 25,000 and 21,500 cm~, attributable to bidimensional patches of pentacoordinated vanadium ions (Fig. 4d) and tridimensional V2Os crystallites (Fig. 4e) [ 11], with V 5+ in octahedral coordination (Oh), feature the spectra of the highly loaded VPS 20 and VPS 50 catalysts. Oxygen chemisorption data match with the above spectroscopic findings, since the maximum oxide dispersion, found for low loaded VPS 2 and VPS 5 catalysts, really corresponds to the maximum development of mono/oligomeric Td structures. Thus, the TPR spectra of differently loaded VPS catalysts indicate that the reduction pattern of V205/SIO2 depends upon the nuclearity of surface V species; in fact a rise in the extent of agglomeration of V 5+ ions causes a change in their coordination symmetry [ 11] rendering them less reducible. In fact, the very sharp reduction peak featuring the TPR spectra of low-medium (2-5.3 wt%) VPS catalysts signals mostly the presence of an easily reducible "surface V205 species" characterized by a rather uniform interaction strength with the PS support. For V205 loadings higher than 5 wt%, the concomitant shift of To,tea and TM~ to higher T, the increasing intensity of TM2-TM3peaks, characteristic of the reduction of bulk V205 (Fig. 3B), and the lowering in oxide dispersion (Table 3) altogether provide unambiguous evidences of the incipient nucleation of VO43 units into Td oligomeric and polymeric structures, like "polyvanadates" (Pv), and V205 clusters (Vc). Pv and Vc exhibit a lower reducibility than Iv species because of the change in V 5+ symmetry from Td to square-pyramidal (Sp) and octahedral coordination (Oh) respectively [6,11 ]. 4. SUMMARY The direct relationship between reaction rate and p (Fig. 2A) indicates that the selective oxidation of light alkanes on low loaded (0-5 wt%) MPS and VPS catalysts proceeds mainly via the surface mechanism [4-8]. The addition of V205 to SiOz implies a simultaneous increase in P and reactivity whilst MoO3 negatively affects both 9 and reaction rate[8]. Indeed, on increasing the MoO3 loading, the decrease in BET S.A. (Table 1) and the increasing coverage of the silica surface concur to lower the concentration of active sites of the SiOz surface causing, mainly at T<_650~ a regular decrease in SSA of MPS systems with respect to PS (Fig. 1A). Yet, a corresponding increase in HCHO selectivity leads to a featureless trend of the Yield at 600~ with the O2 uptake as shown in Figure 5A. At higher T, the capability of MoO3 promoter to interact with CH4 molecules [6] improves the SSA of MPS catalysts with respect to PS (Fig. 1A). Then, the increasing oxygen uptake (Table 3), paralleling the higher concentration of Mc (Table 4), accounts for the promoting effect of MoO3 loading on both SSA and Yield of MPS catalysts in MPO at T>650~ [7]. This is well supported by the straight relationship between Yield to HCHO at 800~ and 02 uptake, shown in Figure 5A. The negative effect of MoO3 on the functionality of PS surface also in POD results in a fiat trend of the Yield to C3H6 at 500~ vs. the 02 uptake (Fig. 5A). Such results likely reflect the particular functionality of"Mo=O '' terminal bonds of MoO3 crystallites towards the formation of selective oxidation products [6,12].

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(~ " 5'0 100 1,50 (3 2'0 4'0 ' 6'0 ' 02 uptake (nmol.mcat "2) V205 dispersion (%) Figure 5. Yield to HCHO (11, 0 ) and to C 3 H 6 (rl)ofMPS (A) and VPS (B) catalysts in MPO and POD reactions vs. O2 uptake and V205 dispersion respectively. The sudden growth in activity along with the moderate increase in p (Fig. 2B) observed for the VPS 5 catalyst suggest that at medium loading vanadia exerts a positive influence on the functionality of the PS support because of the stabilization of dispersed surface species enabling the formation of very active surface reduced sites [7]. Namely, on the basis of characterization data we infer that the highest performance of low-medium loaded (2-10 wt%) VPS catalysts stems from Iv species which ensure the formation of high extents of "active" oxygen species [5,6,8,13]. The role of Iv species in the formation of selective oxidation products is strongly supported by the exponential-like relationships between the specific oxide productivity to HCHO and C3H6 in MPO and POD respectively and V2Os dispersion, shown in Figure 5B. By contrast, the less effective redox mechanism becomes predominant for highly loaded (>5 wt%) VPS catalysts and bulk V system [8] as confirmed by the volcano-shaped relationship between reaction rate and p (Fig. 2B). Then, on highly loaded VPS catalysts, the high density of reduced sites favors a rapid incorporation of activated oxygen species into oxide lattice [8] allowing the occurrence of a classical redox mechanism which is less effective and selective than the surface concerted mechanism [6-8]. REFERENCES 1. V.D. Sokolovskii, Catal. Rev.-Sci. Eng., 32 (1990) 1 2. E.A. Mamedov and V. Cort6s Corberb.n, Appl. Catal. A, 127 (1995) 1 3. A. Parmaliana, V. Sokolovskii, D. Miceli, F. Arena and N. Giordano, J. Catal., 148 (1994) 514 4. F.Arena, F. Frusteri, A. Parmaliana and N. Giordano, Appl. Catal. A, 125 (1995) 39 5. A. Parmaliana, V. Sokolovskii, F. Arena, F. Frusteri and D. Miceli, Catal. Lea., 40 (1996) 105 6. F. Arena, N. Giordano and A. Parmaliana, J. Catal., 167 (1997) 66 7. A. Parmaliana and F. Arena, J. Catal., 167 (1997) 57 8. A. Parmaliana, F. Arena, V. Sokolovskii, F. Frusteri and N. Giordano, Catal. Today, 28 (1996) 363 9. M. Puglisi, F. Arena, F. Frusteri, V. Sokolovskii, A. Parmaliana, Catal. Lea., 41 (1996) 41 10. F. Arena and A. Parmaliana, J. Phys. Chem., 100 (1996) 19995 11. M. Schraml-Marth, A. Wokaun, M Pohl and H.-L. Krauss, J. Chem. Soc., Faraday Trans., 87 (1991)2635 12. M.R. Smith and U. S. Ozkan, J. Catal., 141(1993)124 13. L. Owens and H. H. Kung, J. Catal., 144(1993)202

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

357

M e c h a n i s t i c Aspects of P r o p a n e O x i d a t i o n over N i - C o - M o l y b d a t e Catalysts David L.Stern a and Robert K. Grasselli b aMobil Technology Company, Strategic Research Center, 600 Billingsport Road, Paulsboro, NJ 08066-0480, USA bUniversity of Delaware, Department of Chemical Engineering, Newark, DE 19716-3116, USA

Mechanistic aspects of propane oxidation over Nio.5Coo.5MoO4 supported on SiO2 reveal, that the homolytic C-H bond breaking of the methylene hydrogen abstraction is the rate limiting step of the reaction, involving most probably an electron rich oxygen of a molybdenum located adjacent to the catalyst's cation (A---O-Mo), with the second hydrogen abstraction involving an oxygen of an adjacent molybdenum to the moiety of the first abstraction (A---O-Mo-O-Mo-O), where Ni comprises a more efficient A atom than Co, and where the transition state is relatively symmetric in nature. The reaction is first order in propane and zero order in oxygen, consistent with a Mars-van-Krevelen mechanism involving lattice oxygen of the catalyst. Because this catalyst produces propylene as the exclusive first formed product from propane at low conversions, and since propane and propylene compete for the metal oxide catalytic sites with similar effectiveness, the named paraffin activating catalyst or a modification therefrom show promise, when combined with known multicomponent Bi-molybdates efficient in olefin oxidation, for the direct conversion of propane to acrylic acid or in presence of ammonia to acrylonitrile.

1. INTRODUCTION Functionalization of paraffins to the corresponding olefins, dienes, anhydrides, unsaturated acids and nitriles is of industrial interest [1 ] since paraffins are much more abundant and less expensive than the corresponding olefins. For this reason research interests have recently shifted away from olefins and towards paraffins as starting feeds for the production of petrochemicals and industrial intermediates. Oxydehydrogenation of paraffins is a viable approach towards this goal. The most studied systems for oxidative propane upgrading are vanadium [2], vanadiumantimony [3], vanadium-molybdenum [4], and vanadium-phosphorus [5] based catalysts. Another family of light paraffin oxidation catalysts are molybdenum based systems, e.g. nickel-molybdates [6], cobalt-molybdates [7] and various metal-molybdates [8-9]. Recently, we investigated binary molybdates of the formula AMoO4 where A = Ni, Co, Mg, Mn, and/or Zn; and some ternary Ni-Co-molybdates promoted with P, Bi, Fe, Cr, V, Ce, K or Cs [10-11 ]. A good representative of these systems is the composition Nio.5Coo 5M004 which was recently selected for an in depth kinetic study [12] and whose mechanistic aspects are now further illuminated here.

358 2. EXPERIMENTAL The Ni-Co-molybdate catalysts were prepared by refluxing the metal nitrates with ammonium heptamolybdate in the presence of silica sol (Ludox AS-40). After drying, pulverizing and sizing, they were air calcined for four hours in air at 290 C ~ and for four hours at 600 C ~ The catalyst contained 80 wt% active phase and 20 wt% silica support. Additional data on catalyst preparation, catalyst evaluation, and analytical methodology were described already earlier [ 10-12].

3. RESULTS AND DISCUSSION The system Ni0.sCoo.sMoO4 is effective for the oxidative activation of n-butane, i-butane, propane and ethane, and the conversion to their respective olefins (Figure 1). The selectivity to useful products is similar for propane and n-butane, but is lower for ethane. Methane is rather refractive, and only trace conversions to formaldehyde were achieved under the conditions studied. As is typical of oxidation catalysts, the selectivity to useful products declines with conversion, as seen for the conversion of propane to propylene (Figure 2) and the conversion of propylene to acrolein (Figure 3), respectively. From these results, Figure 1. Oxidation of Light Alkanes. measured at low conversions under C1 at 770~ C~ C~ at 560~ n-C~ at 500~ essentially differential reactor Feed: 15 Hydrocarbon / 1502 / 70 N2. conditions, it is apparent that propylene is the sole primary product of propane oxidation over this catalyst, since extrapolation to zero propane conversion results in 100 percent propylene selectivity, while the conversion to COx (i.e. CO and CO2) waste products at zero propane conversion extrapolates to zero COx selectivity.

1 ~

~

~

2

,9 "'-, 4"',~ 9

~ H I

I

3~

5 .....

0

~

~1110H 0 ./" i

~,

i

,,ss S

CO, CO 2

Scheme 1. Reaction Network of Propane Oxidation

359 When propylene is used as the feed, acrolein is observed as the major primary oxidation product. These are significant mechanistic observations, since it follows therefrom that all higher oxidized products, when starting with propane as feed must derive from the first formed propylene or a subsequently formed intermediate. A reaction network (Scheme 1) consistent with these observations portends that propane is first oxidized to propylene, which further oxidizes to acrolein, which in turn can be further oxidized to acrylic acid, but more preferentially to COx waste products.

Figure 2. Selectivity vs. Propane Conversion [12].

Figure 3. Selectivity vs. Propylene Conversion [12].

In a separate study, selectivities to useful products at higher conversions were also investigated, for propane and propylene oxidation, respectively (Figure 4). It was observed, as is common for sequential reactions, that propylene yield from propane, and acrolein yield from propylene, go over a maximum as hydrocarbon conversion is increased. The maximum propylene yield from propane oxidation is 16% at 34 propane conversion, while the maximum yield of acrolein from propylene is 6% at 9% propylene conversion. By assuming as a first approximation that direct oxidation of propylene to COx is negligible, consistent with the observations above at low conversions, and that only insignificant amounts of acrylic acid are formed, also consistent with our experiments, then a simplified scheme of series reactions can be written

as follows: A (propane)

kl k2 k3 ---> B (propylene) ---> C (acrolein) ---> D (COx)

Reactions 1 and 2 are both first order in hydrocarbon disappearance as will be seen below. By taking into account the above described maximum yields and utilizing the equations from [13], the relative reaction rates can be calculated as being kz]kl = 3.5, k3/k2 = 13 and k3/kl = 46. From these values it is apparent that high selectivities to propylene can be achieved only at relatively low propane conversions using Ni0.sCo0.sMoO4 as catalyst. Conversely, it might be possible to achieve high yields of oxygenates such as acrolein/acrylic acid or nitriles such as acrylonitrile, by combining the studied catalyst or a modification therefrom with a known

360 multicomponent olefin conversion catalyst as a physical mixture or a conglomerate in a single reactor.

Figure 4. Propylene Yields from Propane Oxidation and Acrolein Yields from Propylene Oxidation vs. Conversion [12].

Reaction orders for propylene and acrolein formation were studied by investigating the rate of propylene formation from propane as a function of oxygen concentration (Figure 5a) and as a function of propane concentration (Figure 5b). The results reveal, that the conversion of propane is zero order in oxygen and first order in propane. This is consistent with a Mars van Krevelen mechanism [14], whereby lattice oxygen of the catalyst is the oxidizing agent, and with methylene hydrogen abstraction being the rate limiting step in the activation of the paraffin. A similar dependence is found for acrolein formation from propylene [ 12]. Reaction orders for COx formation from propane are 1/2 order in oxygen and first order in propane [12] . For propylene oxidation they are 1/2 order in oxygen (Figure 6a), but exhibit a Langmuir-Hinshelwood dependence on propylene (Figure 6b).The mechanistic implication of these data is that at low propane conversions, and hence low propylene concentrations in the gas phase over the catalyst, the in situ first formed propylene product readily desorbes from the catalyst surface, and that the readsorbed propylene converts to acrolein, which is readily converted to COx waste

Figure 5. Reaction Orders for Propylene Formation from products. At higher conversions of Propane; a. Oxygen Dependence, b. Propane Dependence propane, under conditions where the [12]. concentration of propylene is rather high, the acrolein formed therefrom appears to react with another catalytic center, possibly a highly acidic site (Mo-O-H § or a reduced site ( M o O ) leading to waste products, or it reacts without desorption with an

361 adjacent site giving waste products. This explanation is consistent with the observed Langmuir-Hinshelwood dependence for waste formation from propylene. The studied catalyst exhibits a significant primary isotope effect for both propane disappearance (kn/ko = 1.7) with theoretical maximum being 1.95) and propylene disappearance (kH/kD = 1 . 9 ) and a theoretical maximum of 2.05). These large, observed kinetic isotope effects are consistent with a homolytic C-H bond breaking as the rate determining step, and a relatively symmetric transition state; hence, methylene hydrogen abstraction for propane activation and t~-hydrogen abstraction for propylene activation. Competition experiments of propane and propylene [ 12] reveal that propane and propylene compete with similar effectiveness for the catalytic Figure 6. Reaction Orders of CO2 Formation from metal oxide sites, albeit as expected Propylene; a. Oxygen Dependence, b. Propylene propylene is favored by a factor of 2.3. Dependence [12]. Since the operating temperature is rather high, the results also imply that the thermal contribution to the respective C-H bond breaking is significant, diminishing the customary importance of the o~-hydrogen bond weakening in propylene due to the rt-bond interaction of the olefin with the catalyst surface.

4. MECHANISTIC MODEL Under the differential reaction conditions used in this study [ 12], the concentrations of all products in the gas phase are small and therefore their respective surface coverages are small. Under these constraints, the following kinetic expressions apply for the partial oxidation of propane to propylene, and propylene to acrolein, respectively: rc3- = k c3= xc3 ~ (02) ~

(1)

racr. = kacr. xc3

(2)

(02) ~

The deep oxidation rates of propylene to COx ( i.e. CO2 as well as CO ) are described by the rate expression:

362

rcox =

112

kcox_._~Xc3 = xo~ ( 1 + Kc3-xc3-)

(3)

Where XC3 o, XC3=, and x02 are the mole fractions of propane, propylene, and oxygen in the gas phase, kc3--, kacr. and kcox are the rate constants of propylene, acrolein, and COx production, Kc3-- is a propylene adsorption constant, and (02) denotes the zero order oxygen dependence. From our study reported here, from our earlier studies of divalent metal molybdates [ 1011 ], and oxidation literature in general [ 1], we can postulate a reaction mechanism (Scheme 2) which takes all of these factors into account and is consistent with them. Accordingly, methylene bond breaking is the rate determining step in the activation of propane and it proceeds by interaction with the NiA .--(~ OI Co-molybdate catalyst surface. The site /MOo/MO responsible for the methylene abstraction is most likely the electron rich oxygen of a molybdenum atom adjacent to the A atom o" El .o o / I I I (A...O-Mo) on the surface of the AMoO4 /M~176 ~ /M~176 ~ catalyst. One could reason, that the activating sites are comprised of an A (~+~)--O.-MoS+-moiety; i.e., Ni3§ 5§ in Ni HO OH HO OH I I I I -molybdate; where O. denotes the partial ~M~176 ~ radical character of this specific oxygen. And the surface concentration of such paraffin activating sites is anticipated to be relatively small, yet sufficient to catalyze the Scheme 2. Proposed Mechanism of Propane reaction in question. As a consequence, Oxidation over Ni0.sCo0.sMoOx,where A= Ni,Co. Ni...O-Mo is more active than Co...O-Mo, which is much more active than Zn...O-Mo for the activation of paraffins [11]. The second hydrogen abstraction from the so formed propyl radical is performed by a Mo-O moiety adjacent to the activating A...Mo-O site, i.e., (A...Mo-O-Mo-O), as depicted in Scheme 2. Propylene desorbs, leaving a hydroxylated site, which after dehydration forms a reduced surface site, which in turn is reoxidized by lattice oxygen of the catalyst. The depleted lattice oxygen is then replenished by gaseous oxygen, thus completing the redox cycle of the catalytic process. Waste products are formed from propane on this catalyst primarily via the second formed acrolein. It is postulated that the acrolein, once formed, readily readsorbs on the catalyst surface, and presumably interacts with a second site which might be either a highly acidic surface site (Mo-O- H +) or a reduced surface site ( Mo I-i ) or simply by interaction, before desorption, with an adjacent overactive surface site. This scenario is particularly strongly implied by the observed Langmuir-Hinshelwood dependence of waste formation from propylene. In order to minimize waste formation and maximize useful product yields, several opportunities reveal themselves from the observations gained in our studies and those gleaned

363 from oxidation catalysis in general, ff propylene is to be the desired end product, then operation at low conversions using a Ni-Co-molybdate catalyst is possible. An improvement would be operation at very short contact times and particularly by employing a transferline reactor in the absence of gaseous oxygen, using external regeneration of the catalyst. Another potential improvement would be to use a redox promoted Ni-Co-MoOx system [10-11], because with such systems the reaction could be carried out at lower temperatures, and hence improved selectively. / 0\

/0/

NH3

Olefin , NH3 n ,s\ ctivating Site

-~ [] M1

/ [ ~ " ~2 ~ R e ~ 1 7 6 [1 # M2 J

Reducedsite

! HzC~-CHCN

3M [ ]

3/2 02 J

~ -~ //NH M2

..-O"

" M1

[o12-

~

1/2o2

x~.~.~ H20

/\

~ M1

JM~176

NH O" [] I I // / M ~ 1 7 6 ~ M2

AmmoxidatiOnsite H20" ~

H~

.o-

. oI

.,Mo _Mo

.~H HzC.-" ~""Z' " ~H2 ,

........... .

.

.

.

.

.

.

.

.

.

.

.

O

- d 2 ' HO

Paraffin Activating Site

~

I

O I I / M~ O/M~

oI

J

jMO~oIMO ~

.

M1 M2 AllylicSurface Complex

Olefin Activation (where M]=Bi, M2=Mo)

Paraffin Activation (where A=Ni, Co)

Scheme 3. Proposed Mechanism of Propane Ammoxidation using a Paraffin Activating Catalyst, e.g.,NiaCObMcMoOx; and a Multicomponent Mixed Metal Molybdate Olefin Ammoxidation Catalyst, e.g., CSaKbNicMgdMeBifSbgMohOx, where M = Ce, Cr, or Fe [ 15]. Still another scenario, and one which we feel has particular promise, is starting with propane to aim for and isolate a useful product more stable than propylene or acrolein. To achieve this [15], it might be possible to combine the paraffin activation prowess of the systems discussed here and earlier [10-12], with the olefin conversion efficiency of known multicomponent molybdate catalysts [16] to produce more stable compounds such as unsaturated acids (e.g., acrylic acid) or nitriles (e.g., acrylonitrile). The two catalysts could be commingled as a physical mixture in a single reactor, or sequenced and commingled in a single catalyst composition. This idea is presented in Scheme 3 for the production of acrylonitrile from propane. In this scheme the propane is activated on a Ni-Co-molybdate catalyst, preferably one doped with redox elements such as Ce or Cr, producing propylene as the major intermediate. The so produced propylene reacts then further with the olefin conversion catalyst, which is selected from one of the well known alkali doped Bi- molybdate multicomponent catalysts, preferably one containing a redox element such as Ce , e.g. CSaKbNicMgdCeeBifSbgMohOx. In the presence of ammonia, the in situ produced propylene is

364 immediately converted to acrylonitrile, a useful product of commerce, and relatively stable under the conditions of the reaction. It is important in this processing scheme that the two catalysts are chemically compatible (therefore both are selected from the molybdate family), and that they are temperature matched. The latter implies that compositions must be found which operate at a compromise optimum temperature. For this reason the paraffin activating catalyst is redox doped with say Ce or Cr, to lower its operating temperature, enhance its activity and improve its selectivity towards the production of propylene. Similarly, the olefin activating catalyst is doped with Ce instead of the customary Fe, in order to increase its operating temperature in the direction of the operating temperature of the paraffin activating catalyst. It is of course also imperative, that the two catalysts be in intimate contact with each other. This is achieved by commingling the two compositions in a single extrudate or coprecipitate, or by commingling physical mixtures of the two catalyst powders (small particle size) in a single reactor. Additional experimental work is needed to optimize the choice of the named catalyst compositions, and to optimize the reaction and reactor conditions.

5. CONCLUSIONS Ni-Co-Molybdates are viable oxidation catalysts for the activation of light paraffins such as propane and butanes to produce the respective olefins. Maximum yields are in the range of 16% at about 80% selectivity. The catalysts activate methylene C-H bonds, abstracting the hydrogen of the substrate in the rate limiting step of the reaction. With propane as feed, propylene is the only first formed product, and all higher oxidized products ensue in subsequent steps, after the propylene has been formed. Acrolein is formed from the in situ produced propylene, and acrolein is the main intermediate leading to waste products CO and CO2. The first hydrogen abstraction is deemed to involve the oxygen of the Mo adjacent to the A metal of the catalyst, i.e. A---O-Mo, with the reactivity order of the A metal being Ni > Co >> Zn. The second hydrogen of the so formed propyl radical is deemed to occur by an oxygen of a Mo located adjacent to the first abstracting site, i.e. A---O-Mo-O-Mo-O. Propylene is desorbed from the catalyst surface, and after loss of a water molecule the so formed reduced catalyst site is reoxidized through lattice oxygen to is original state, and the depleted lattice oxygen is in turn replenished by gaseous oxygen, which completes the redox catalytic cycle. The large primary kinetic isotope effects observed for propane and propylene oxidation over these catalysts suggest homolytic C-H bond breaking as the rate limiting step, and a relatively symmetric transition state. Propane and propylene compete with similar effectiveness for the metal oxide catalytic sites, with propylene favored by a factor of 2.3. The mechanistic results of this and earlier studies [ 12] suggest that in addition to producing propylene from propane, the described paraffin activating catalysts could be combined with known olefin oxidation catalysts in a single reactor to produce industrially desirable value added products which are more stable than propylene under the reaction conditions employed; such as acrylic acid or acrylonitrile.

365 ACKNOWLEDGEMENTS

The authors thank Professor D. J. Buttrey, University of Delaware, for fruitful discussions of structure-mechanism relationships, and to Mr. B. Leyer, University of Munich, for the artwork.

REFERENCES

1.a.R.K. Grasselli and J.D. Burrington, Adv. Catal., 30 (1981) 133 b. R.K. Grasselli, J. Chem. Ed., 63 (1989) 216. 2.a.M.C. Kung and H.H. Kung, J. Catal., 134 (1992) 668. b. A. Corma, J.M. Nieto Lopez and N. Paredes, ibid, 144 (1993) 425. 3.a.A.T. Guttmann, R. K. Grasselli, J.F. Brazdil and D.D. Suresh, US Patent No. 4 746 641 (1988). b. R. Catani, G. Centi, F. Trifiro and R.K. Grasselli, Ind. Eng. Chem. Res. 31 (1992) 107. c. A. Andersson, S.L.T. Andersson, G. Centi, R.K. Grasselli, M. Sanati and F. Trifiro, Appl. Catal. A, 113 (1994) 43. 4.a. Y-C. Kim, W. Ueda and Y. Moro-oka, Catal. Today, 13 (1992) 673. b. J.P. Bartek, A.M. Ebner and J.F. Brazdil, US Patent No. 5 198 580 (1993). 5.a.M. Ai, Catal. Today, 12 (1992) 679. b. M. Ai, J. Catal.,101 (1986) 389. 6.a.C. Mazzocchia, E. Tempesti and C. Aboumard, US Patent No. 5 086 032 (1992). b. C. Mazzocchia, C. Aboumard, C. Diagne, E. Tempesti, J.M. Herrmann and G. Thomas, Catal. Lett., 10 (1991) 181. 7.a.H.F. Hardman, US Patent No. 4 131 631 (1978). b. H.F. Hardman, US Patent No. 4 255 284 (1981). 8. F. Cavani and F. Trifiro, Catal. Today, 24 (1995) 307. 9. Y.S. Yoon, N. Fujikawa, N. Ueda. Y. Moro-oka and K.W. Lee, ibid, 24 (1995) 327. 10. D.L. Stern and R.K. Grasselli, 211 Nat. Mtg., ACS, Petr. Chem. Pre-Prints, (1996) 172. 11. D.L. Stern and R.K. Grasselli," Propane Oxydehydrogenation over Molybdate Catalysts" J. Catal., (1997) in press. 12. D.L. Stern and R.K. Grasselli, "Reaction Network and Kinetics of Propane Oxydehydrogenation over Nickel Cobalt Molybdate" J. Catal., (1997) inpress. 13. A.A. Frost and R.G. Pearson, "Kinetics and Mechanism" John Wiley & Sons, (1953). 14. P. Mars and D.W. van Krevelen, Chem. Eng. Sci. Suppl. 3 (1954) 41. 15. R.K. Grasselli, Handbook of Heterogeneous Catalysis, G. Ertl, H. Knoezinger and J. Weitkamp (eds), B (1997) 4.6.7. 16.a.R.K. Grasselli, D.D. Suresh and H.F. Hardman, US Patent Nos. 4 139 552 (1979), 4 162 234 (1979), 4 190 608 (1980), 4 778 930 (1988). b. R.K. Grasselli and H.F. Hardman, US Patent No.4 503 001 (1985). c. R.K. Grasselli, Appl. Catal., 15 (1985) 127. e. D.D. Suresh, M.S. Friedrich and M.J. Seely, US Patent No. 5 212 137 (1993).

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

367

Oxidative Dehydrogenation of Propane by Non-Stoichiometric Nickel Molybdates* Doron Levin and Jackie Y. Ying* Department of Chemical Engineering, Massachusetts Institute of Technology, Cambridge, MA 02139 A series of catalysts represented by the formula Nil+6Mol_~/304, where - 1/5 _<6 _< 1/3, were investigated for activity towards oxidative dehydrogenation of propane. The catalysts were synthesized from layered ammonium nickel molybdate materials having a general formula (NH4)HzANi3_xO(OH)(MoO4) 2, where 0 _< x _< 3/2. The location of the nickel atoms on a crystallographic site of variable occupancy allowed for the careful manipulation of chemical composition of these precursors. The nickel/molybdenum ratio of the precursor was inherited by the catalyst, enabling the synthesis of non-stoichiometric molybdates. These catalysts were tested for the catalytic oxidative dehydrogenation of propane to elucidate the effect of the extent of nonstoichiometry (6) on catalytic performance. The activities of the catalysts are highly sensitive to the catalyst composition and increase with decreasing Ni content. Selectivity towards propene is independent of the Ni content when the Ni/Mo ratio is less than unity, but decreases linearly with increasing amounts of Ni when Ni/Mo > 1. 1. I N T R O D U C T I O N Transition metal molybdates have been examined as multicomponent catalytic systems for a number of reactions. They are well known to be catalytically active for partial oxidation reactions, particularly for the selective oxidations of lower alkanes. There is currently increased industrial interest in the oxidative dehydrogenation (ODH) of paraffins to olefins [1], due to several limitations in thermal and catalytic dehydrogenation processes. The thermodynamics of thermal light paraffin dehydrogenation are such that under conditions of low temperature and high pressure, the equilibrium favors the paraffin. Catalytic dehydrogenation at elevated temperatures leads to undesirable side reactions such as cracking and coking of the catalyst. In addition, paraffin dehydrogenation is strongly endothermic, making the process extremely energy intensive. To supply the heat of reaction, hydrogen oxidation can be coupled with dehydrogenation. The addition of oxygen to the feed, however, exposes the synthesized olefins to oxidation conditions that could result in the formation of economically-useless carbon oxides. Desired catalysts for oxidative dehydrogenation of lower alkanes need to be sufficiently active for hydrogen atom

* This work was supported by the National Science Foundation (CTS-9257223, CTS-9411901) $ To whom correspondence should be addressed

368 abstraction from a C-H bond, yet operate at temperatures that minimize oxygenation of the desired products. Catalysts based on transition metal molybdates, typically bismuth, cobalt and nickel molybdates [2-6], have received recent attention. Of the transition metal molybdates, those based on nickel, and in particular the stoichiometric NiMoO4, have attracted the greatest interest. NiMoO 4 presents two polymorphic phases at atmospheric pressure: a low temperature a phase, and a high temperature [3 phase [2,7]. Both phases are monoclinic with space group C2/m. These phases differ primarily in the coordination of molybdenum which is distorted octahedral in the phase and distorted tetrahedral in the [3 phase. The [3 phase has been shown to be almost twice more selective in propene formation than the a phase for comparable conversion at the same temperature [2]. A similar effect has been noted for oxidative dehydrogenation of butane, with the 13phase being approximately three times more selective in butene formation than the o~phase [8]. The reason for the difference in selectivities is unknown, but the properties of the phases are known to be dependent on the precursors from which they are derived. Typically, nickel molybdates are prepared by calcination of precipitated precursors. The structure of one such precursor has recently been solved [9]. This precursor, having a nickel to molybdenum molar ratio equal to one, forms a stoichiometric NiMoO4 upon calcination at 550 ~ The solution of the crystal structure showed the chemical formula of the precursor to be (NHn)HNi2(OH)2(MoO4) 2. A detailed investigation into the structure of this material showed it to be a member of a solid solution series of (NHa)H2xNi3_xO(OH)(MoO4) 2, where 0 _<x _<3/2 [9]. The non-stoichiometry arises from variable occupancy of the octahedral site accommodating the nickel atoms. The nickel atoms are situated at a crystallographic site having a Wyckoff designation 9(e) and it has been shown that the occupancy of that site can vary from V2to 1 [9]. The molybdenum atoms have been determined to fully occupy site 6(c), thereby allowing the Ni/Mo ratio to vary from 0.75 to 1.5. The variation in the Ni content is made possible by the ability of the structure to accommodate protons required for charge balancing. The accommodation of the nickel on a crystallographic site of variable occupancy allows for the careful manipulation of chemical composition of these precursors. This structure is also not limited to this class of nickel molybdates, but can be synthesized with a variety of elements such as Co, Zn, Mg, and Cu, either alone or in combination with others [ 10]. These materials form a class of layered transition metal molybdates, termed LTM. It is these properties of this LTM precursor that make it so versatile for the preparation of engineered catalytic phases. Since the properties of a catalyst will be dependent on the precursor from which it is derived, the ability to systematically manipulate the chemical composition of the LTM precursor allowed for a detailed examination of the structure/property relationship for these materials.

2. EXPERIMENTAL 2.1. Catalyst Preparation The ammonium nickel molybdate precursors were prepared by chemical precipitation [9]. Ammonium heptamolybdate ((NH4)6MOTOz4.4H20) and nickel nitrate (Ni(NO3)2.6H20) were used

369 to prepare a solution containing Ni and Mo in a molar ratio between 0.75 and 1.5. The addition of concentrated ammonium hydroxide (28.8% NH3) precipitated a green solid that dissolved in an excess of ammonia to give a deep blue solution. This solution was heated with constant stirring for four hours, leading to the formation of a pale green precipitate. The products were isolated by vacuum filtration, washed with deionized water, and dried overnight at 110 ~ and atmospheric pressure. The catalysts of the form Ni~+sMOl_6/304, as indicated in Table 1, were prepared by in situ calcination of the LTM precursors at 550 ~ in flowing He for 1 hour. Table 1 Catalysts Derived from (NH4)Hz~Ni.~_,O(OH)(MoO4)z Precursors x Ni/Mo Precursor Formula Catalyst Formula Ratio 0.0 1.5 (NH4)Ni30(OH)(MoO4) 2 Nil.333Moo.8890 4 0.5 1.25 (NH4)Ni2.5 (OH)z(MoO4)2 Nil.|v6Moo.9410 4 1.0 1.0 (NH4)HNiz(OH)z(MoO4) 2 NiMoO4 1.25 0.875 (NHa)H l.sNi 1.75(OH)z(MoO4)2 Ni0.903MOl.03204 1.5 0.75 (NH4)HzNil.5(OH)z(MoO4) 2 Ni0.g0Mo~.o670 4

0.333 0.176 0.000 -0.097 -0.200

2.2. Catalytic Activity Measurements The catalytic activity of the nickel molybdates [ 11 ] was tested in a 1/4 inch quartz flowthrough tubular reactor operated at atmospheric pressure. The reactor was contained within an electrically heated tube furnace. The temperature of the reactor was controlled according to the temperature of the gases at the base of the catalyst bed. The composition and flow rate of the gas feed mixture was measured using MKS mass flow controllers calibrated for each specific gas. Certified gas mixtures with Grade 5 helium (99.999%) as the balance gas were used throughout. The reactor effluent was analyzed by gas chromatography using a Perkin Elmer Autosystem GC equipped with a TCD detector. A two-column setup incorporating a 15-foot Poropak QS column and a 6-foot Carbosphere column was used with a 10-port sampling valve to fully separate the gases. This setup resulted in good separation of the individual gases, allowing for closure of the carbon balance to within + 5%. Conversion has been defined as the percentage conversion of propane into all possible products. Selectivity has been defined as the number of moles of propane converted into a specific product divided by the total number of moles of propane converted into all products, expressed as a percentage. The yield of propene has been defined as the product of the propane conversion and the selectivity towards propene, and is expressed as a percentage.

3. RESULTS AND DISCUSSION The series of nickel molybdates tested for oxidative dehydrogenation of propane produced a product spectrum limited to propene, carbon dioxide, carbon monoxide, and water. No cracking

370 products such as ethane, or oxygenated products such as aldehydes were present in the exhaust from the reactor.

3.1. Effect of Non-Stoichiometry The yield of propene attainable with nickel molybdates of the form Nil+6Mol_6/304, where -1/5 _< 6 _< 1/3, is shown in Figure 1. These data were collected at 500 ~ with a feed flowrate of 30 ml/min composed of 5 mol% O2, 5 mol% C3H8, and the balance being He. The data of Figure 1 show an increase in the yield of propene with a decrease in the nickel content of the materials. The material with the highest Ni/Mo ratio of 1.5 is a better catalyst for combustion than oxidative dehydrogenation, with approximately 60% of the product spectrum being carbon oxides. The material with the lowest Ni/Mo ratio of 0.75 has the highest activity under the reaction conditions, leading to the highest propene yield.

Figure 1. Propene yield as a function of Ni/Mo ratio.

Figure 2. Propene selectivity at 20% propane conversion.

The series of nickel molybdate catalysts show a decrease in the selectivity towards propene with an increase in conversion. Figure 2 shows the selectivity towards propene at a propane conversion level of 20%. As the Ni/Mo ratio increases above 1, the selectivity towards propene decreases almost linearly as carbon oxides become the dominant products. It is interesting to note, however, that the selectivity towards propene is essentially independent of the Ni/Mo ratio at values less than and equal to 1. This suggests that it is the increase in activity with decreasing Ni/Mo values that accounts for the increase in yield noted above.

3.2. Effect of Oxygen:Propane Ratio To further improve the yield of propene attainable with the non-stoichiometric nickel molybdates, the oxygen:propane molar ratio was varied from a 1:1 ratio to an oxygen-rich 2.5:1 value. The results of this study are shown in Figure 3. These data were collected at 550 ~ with

371 a feed flowrate of 70 ml/min composed of 1 mol% C3H8, an 0 2 concentration set by the ratio under investigation, and the balance being He. In an analogous set of results to the data shown in Figure 1 collected at a lower space velocity, the series of nickel molybdates showed an increase in propene yield with a decrease in the Ni/Mo ratio. These data also show that the propene yield increases slightly as the oxygen:propane ratio increases, but with the increase becoming less significant as the Ni/Mo ratio decreases.

Figure 3. Propene yield as a function of

Figure 4. Propane conversion as a function

Ni/Mo ratio and 02:C3H 8 ratio.

of Ni/Mo ratio and 02:C3H 8 ratio.

The effect of increasing the oxygen:propane ratio is more evident when examining the propane conversion. As shown in Figure 4, the propane conversion increase with increasing oxygen concentrations is significant for all Ni/Mo ratios. However, as the Ni/Mo ratio decreases, the higher conversion resulting from a higher oxygen concentration does not lead to higher propene yields due to a shift in selectivity towards carbon oxides.

3.3. Structure/Catalytic Property Relationship Physical characterization of the phase present under catalytic reaction conditions has shown a single [3-phase of Nil+6Mo~_6/304 is present at 550 ~ following calcination of the (NH4)HzxNi3_xO(OH)(MoO4)2 precursors [12]. Analysis of the defect chemistry of these nonstoichiometric nickel molybdates identified majority point defects, leading to a correlation between the electrical conductivity and the defect structure [ 12]. This correlation is summarized below. For nickel molybdates having a Ni/Mo ratio greater than one, corresponding to an excess of NiO, the structure is proposed to have interstitial Ni atoms as the major defects that are compensated for by the presence of Mo vacancies. Given this excess of Ni atoms, it was proposed that the majority carrier arises from the oxidation of Ni u to Ni m, represented by the equation,

372 1

2 Nixi + ~ O 2 ( g ) -

2 NiNi + 0 o + Vffi

(1)

This formation of Ni In leads to the formation of electron-acceptor levels close to the valence band, leading to hole conductivity (p-type semiconductor.) For nickel molybdates having a Ni/Mo ratio less than one, corresponding to an excess of M o O 3, the structure is proposed to have Ni vacancies as the major defects that are compensated for by either interstitial Mo atoms, or Mo atoms occupying Ni sites. Given this excess of Mo atoms, it was proposed that the majority carrier arises from the reduction of Mo vx to Mo v, represented by the equation 2MOMo + 0 o

1

=,

~O2(g )

/

+

V 0

+

Z

2MoMo

(2)

This formation of MoV leads to the formation of electron-donating levels close to the conduction band, leading to electron conductivity (n-type semiconductor.) It has been reported in the literature that adsorbed oxygen species transform at the surface of an oxide according to the general scheme: 2-

02(ads)-*

O•ads ) ~ O(ads ) --' O(lattice )

(3)

in which they are gradually becoming richer in electrons [ 13]. Transition metal oxides containing cations that are capable of increasing their degree of oxidation and thereby supplying adsorbed oxygen molecules or atoms with electrons (i.e. p-type materials) tend to form electron-rich species such as O- and O 2-. Transition metal oxides that are n-type in nature have a small concentration of donor centers capable of transmitting electrons to the adsorbed oxygen and tend to form oxygen species that are less rich in electrons, such as O 2 ions. Focusing on the nickel molybdate having a Ni/Mo ratio of 1.5 (Nil.333Mo0.88904) , it can be seen that both the high number of Ni atoms capable of oxidation and the low number of Mo atoms capable of reduction leads to a high number of electrons available for the formation of electron-rich species. It is therefore proposed that the dominant surface oxygen species on this nickel molybdate is O-. Turning to the nickel molybdate having a Ni/Mo ratio of 0.75 (Ni0.80Mo~.06704), it can be seen that both the low number of Ni atoms capable of oxidation and the high number of Mo atoms capable of reduction leads to a low number of electrons available for the formation of electron-rich species. It is therefore proposed that the dominant surface oxygen species on this nickel molybdate is 02-. With these two nickel molybdates forming the bounds on the range of Ni/Mo ratios possible, it can therefore be proposed that as the Ni/Mo ratio is varied through this range, the

373 nature of the surface oxygen species changes with increasing amounts of electron-rich species as the Ni/Mo ratio increases. It must, however, be mentioned that this would only be true with 02 as the oxidant. It has been shown that conversion decreases with an increase in the Ni/Mo ratio (Figure 4). This can be attributed to a decrease in the reactivity of the surface oxygen species as the nature of the surface oxygen coverage shifts from an O2-majority to an O- majority. The catalytic data suggests that the O2- form of adsorbed oxygen is more reactive than O- for H atom abstraction from propane. When O2 is used as the oxidant, the amount of electrons available for donation to adsorbed oxygen molecules decreases as the Ni/Mo ratio decreases, resulting in a more reactive oxygen surface coverage that leads to higher propane conversion. 4. CONCLUSIONS The series of non-stoichiometric ammonium nickel molybdate precursors having the general formula (NH4)HzxNi3_xO(OH)(MoO4)2, where 0 _< x _< 3/2, is an excellent family of materials for the preparation of transition metal molybdate catalysts having well-controlled stoichiometry. This flexibility allowed for a systematic study into the relationship between catalytic performance and catalyst structure. This class of LTM precursors encompassing a Ni/Mo ratio of 0.75 to 1.5 forms pure ~-Nil+6Mol_6/304 on calcination at reaction conditions. These nickel molybdates are catalytically active for the oxidative dehydrogenation of propane into propene, with the propene yield being maximized by minimizing the Ni/Mo ratio.

REFERENCES

.

10. 11. 12. 13.

Cavani, F., and Trifirb, F., Catalysis Today 24, 307 (1995). Mazzocchia, C., Aboumrad, C., Diagne, C., Tempesti, E., Herrmann, J.M., and Thomas, G., Catal. Lett. 10, 181 (1991). Mazzocchia, C., Anouchinsky, R., Kaddouri, A., Sautel, M., and Thomas, G., J. Therm. Anal. 40, 1253 (1993). Mazzocchia, C., Kaddouri, A., Anouchinsky, R., Sautel, M., and Thomas, G., Solid State Ionics 63, 731 (1993). Minow, G., Schnabel, K., and Ohlmann, G., React. Kinet. Catal. Lett. 22, 389 (1983). Yoon, Y. S., Fujikawa, N., Ueda, W., and Moro-oka, Y., Chem. Lett. 1635 (1994). Mazzocchia, C., Di Renzo, F., Aboumrad, Ch., and Thomas, G., Solid State Ionics 32, 228 (1989). Martin-Aranda, R. M., Portela, M. F., Madeira, L. M., Freire, F., and Oliveira, M., Appl. Catal. A 127, 201 (1995). Levin, D., Soled, S. L., and Ying, J. Y., Inorg. Chem. 35, 4191 (1996). Levin, D., Soled, S. L., and Ying, J. Y., Chem. Mater. 8, 836 (1996). Levin, D., and Ying, J. Y., J. Catal., to be submitted. Levin, D., and Ying, J. Y., J. Electroceram., to be submitted. Bielmiski, A., and Haber, J., Catal. Rev. - Sci. Eng. 19, 1 (1979).

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

375

Selective oxidation of propane into oxygenated compounds over promoted nickel-molybdenum catalysts J. Barrault a, C. Batiot a, L. Magaud a and M. Ganne b a Laboratoire de Catalyse en Chimie Organique, ESIP, UMR CNRS 6503 40 avenue du Recteur Pineau, 86022 Poitiers cedex, France

b Laboratoire de Chimie des Solides, Institut des Materiaux de Nantes 2 rue de la Houssiniere, 44072 Nantes cedex 03, France Partial oxidation of propane was investigated in the presence of molybdenum oxide based catalysts. We have shown the existence of a synergetic effect between the two phases (~NiMoO4 and ~MoO3 . Indeed activity and selectivity towards acetic acid and acrylic acid were maximal with a ratio t~MoO3 / (otNiMoO4 + t~MoO3 ) close to 0.25. These results could be explained by an interaction and a mutual covering of the two phases. The addition of bismuth to these mixed systems led to a total or a partial inhibition in the production of acetic acid and an increase in the formation of acrolein and acrylic acid. 1. INTRODUCTION The selective oxidation of light alkanes has attracted much attention because it represents a route to obtain more valuable organic compounds from low cost saturated hydrocarbons. Numerous works have been carried out in the last decade on the oxidative coupling of methane (OCM) and the selective oxidation of the other light paraffins (ethane, propane, isobutane). The oxidative dehydrogenation of propane to propene and its selective oxidation to acrolein and acrylic acid were studied over various solids. With catalysts containing molybdenum Morooka et coll. claimed that a selectivity to acrolein of over 60% for a 13% conversion of propane in the presence of a silver doped bismuth vanadomolybdate [1,2]. Mazzochia et coll. proposed that the propane oxidative dehydrogenation (P.O.D.) occurred mainly over a 13NiMoO4 phase (containing tetrahedral molybdenum centers) [3]. Previously these authors found a synergetic effect on acrylic acid selectivity by adding MoO3 to otNiMoO4 [4]. Schrader et coll. also showed that nickel molybdenum catalysts contained mainly a mixture of otNiMoO4 and t~MoO3 [5,6], and that the catalytic properties resulted of <<specific interactions )) between the two phases. More recently Y.S. Yoon et coll. proposed that the P.O.D. involved surface molybdenum oxide clusters supported on metal molybdate matrix when Using cobalt and magnesium molybdates [7]. In our laboratory we showed that Bi-Mo (W,V, Ti) Aurivillius or Sillen phases and bismuth-molybdenum oxides supported on titania were more selective for the P.O.D. to propene [8,9]. In this field our objective was to improve the activity of molybdenum based materials either for the P.O.D. to propene or the selective oxidation of propane to acrolein and acrylic acid. For these reasons we have focused on the catalytic performances of NiMoO4-MoO3 phases doped with bismuth.

376 2. SYNTHESIS AND CHARACTERIZATION OF MATERIALS 2.1 Synthesis of materials and reaction procedure Nickel - molybdenum catalysts were prepared by coprecipitation to obtain a significant surface area ( 5-50 m2gl). The preparation procedure was an optimisation of the procedure proposed by Mazzocchia [3,4] and Schrader [5]. 13NiMoO4 was obtained after a thermal treatment of ct phase at 700~ in oxygen for 1 h. Nevertheless this ot<----~13 transition was reversible and occurred below 300~ [3]. In order to compare all the (Ni-Mo) samples, we used a molar ratio RA = ot M o O 3 / (r M o O 3 + ~ NiMoO4) Bismuth compounds were added to some of the previous samples. These modifications consisted of an impregnation of bismuth salt or of a precipitation of bismuth hydroxide over the [ nickel-molybdenum] solids. Prior to the reaction lg of catalyst was activated under helium at 500~ The reactions were performed in a flow reactor at a temperature of 375-425~ under atmospheric pressure. The feed gas consisted of a 60 vol % propane, 20 vol% 02 and 20 vol % He or H20. The total flow rate was 1.5 l.h 1. The outlet gases were analysed by FID and TCD gas chromatography with Cpsil5, BP 21, Porapack R and Porapack Q columns.

2.2 Characterization of materials All the X R analysis of nickel-molybdenum catalysts with 0
Chemical Analysis (%)

BET (m2.g"1)

Formula

XRD

Mo

Ni

0.01

44.95

26.50

44

Ni0.99MoOx otNiMoO4 (t~ ~MoO3)

0.25

44.90

20.65

29

N i 0 . 7 5 M o O x otNiMoO4, r

3

0.48

49.54

15.74

21

N i 0 . 5 2 M o O x otNiMoO4, r

3

0.83

56.66

5.74

11

N i 0 . 1 7 M o O x otNiMoO4, ~MoO 3

1

-

-

4

t~MoO 3

All the XRD diagrams showed the presence of the monoclinic otNiMoO4 phase and/or the orthorhombic ~MoO3 phase. There were no mixed ~) phases before and after the catalytic test. The formation however of a solid solution undetected by XRD cannot be excluded. The sample with a formula <) was composed of a mixture of two phases after calcination at 550~ and after the catalytic test. After activation at 700~ and the catalytic test the intensity of different peaks ((0k0) with k=2n) of ~ M o O 3 was considerably

377 modified, which indicates a preferential orientation of the (010) phases, which correspond to the reticular planes perpendicular to the b axis of ct M o O 3 . Tables 2 and 3 give the results of the XPS analysis obtained for NiMoO4 and NiMoO4-MoO3 samples (RA=0.83) before and after the catalytic test. Table 2 XPS analysis of NiMoO4 catalysts (1) XPS Analysisat atom. %

(2)

(3)

(1)

eV

atom. %

atom. %

eV

Bulk Analysis ( atom. %)

Mo 3d5/2

31.67

2.20

34.13

32.51

1.40

16.62

Ni

7.96

3.20

6.88

6.65

1.70

16.38

O ls

60.37

2.30

58.98

61.85

1.70

67.00

Mo/Ni+Mo

0.80

-

0.83

0.85

-

0.50

BET(m2.g -1)

44

-

31

20

2p3/2

Table 3 XPS analysis of

MoO 3 - NiMoO 4

catalysts, RA=0.83

(1)

(2)

(3)

(1)

eV

Bulk Analysis

XPS Analysis

atom. %

eV

% atom.

eV.

Mo 3d5/2

29.35

1.80

33.00

1.40

21.72

Ni

3.44

2.60

2.70

2.50

3.75

O ls

67.21

2.30

64.30

1.80

74.53

Mo/Ni+Mo

0.89

-

0.92

-

0.85

BET(m2.g -1)

12

-

10

-

2p3/2

atom. % :surface composition eV : half height width (1) sample after calcination at 550~ (2) sample after thermal treatment at 500~ (or) and catalytic test (3) sample after thermal treatment at 700~ (13) and catalytic test

3

378 We can see that the surface composition is very different from the bulk composition. After calcination and the catalytic test, XPS peaks are thinner and the molybdenum surface content increases. The oxidation degrees of molybdenum (Mo 3d5/2 = 236.7_+0.1 eV, Mo 3d3/2 = 233.6 _+0.1 eV) and the oxidation degrees of nickel (Ni 2p3/2 =856.7_+0.1 eV, Ni 2pl/2 =875.0 _+0.1 eV ) are not modified by the calcination and the catalytic test. The solids however have been in contact with air before the analysis. The significant differences between energy levels Ni2p of NiMoO4 and ofNiO (Ni 2p3/2 =854.4 eV) show clearly that there is no NiO in the samples.

3. RESULTS AND DISCUSSION 3.l.Nickel-molybdenum catalysts 3.1.1. Effect of MoO 3 on the catalytic properties of ctNiMoO4 We studied the influence of MoO 3 on the catalytic properties of aNiMoO4. The catalytic test was carried out 375~ with 0
Table 4 Propane oxidation over [ctMoO3 - otNiMoO4 ] catalysts RA

Conv. (%)

Selectivity (%)

C3H8

02

C3H6

COx

O2/C

Eoxy.

etha.

acro.

0

5.3

48

45

48

2.43

3

2

2.2

0.06

5.1

46

46

48

1.65

3

0.8

0.8

0.4

0.25

11.8

82

35

35

0.92

30

0.7

1.6

13.9

13.0

0.48

10.0

77

40

40

0.86

20

0.8

1.7

8.1

8.2

0.83

6.3

46

48

36

0.67

15

0.8

1.8

6.0

5.8

1.00

0.5

5

75

24

0.52

1

C3Hs/O2/I-Ie " 60/20/20 (%vol. CNTP), F = 1.5 1.h-1, m = 1.0 g

0.6

acet.

acry.

379 3.1.2 Effect of a thermal treatment at 700~ The catalytic properties of t M o O 3 - NiMoO4 ] catalysts after a thermal treatment at 700~ under an oxygen atmosphere (RA'= otMoO 3/(cxMoO 3 + 13NiMoO4 )) are reported in Table 5. Compared to the results reported in Table 4, this activation procedure causes a decrease in the propane conversion especially when RA'> 0.06, which could be related to a decrease of the surface area. These changes are particularly significant for RA'= 0.25 and 0.83 because the acid formation is completly suppressed. ESCA and XRD characterizations carried out after the catalytic experiments reveal a surface enrichment in molybdenum with a specific (010) orientation of crystal faces of MoO 3. The cristalline structure of the solids however is not very much modified so that the most significant factor of the partial covering of NiMoO4 with MoO 3 could be the thermal treatment. Table 5 Propane oxidation over (MoO3-NiMoO4) activated at 700~ under oxygen RA'

BET (m2.g-1)

Conversion (%)

Selectivity (%)

(1)

(2)

C3H8

0 2

C3I-I6

COx

Eoxy.

13NiMoO4

44

23

5.9

32

67

24

6

0.06

42

19

6.3

33

67

20

12

0.25

30

6

1.0

3

90

6

2

0.83

10

3

1.2

4

89

9

1

~MoO3 (3)

4

4

0.5

5

76

24

e

RA '= ~MoO 3 /(~MoO 3 + [3NiMoO4), C3Hg/O2/He : 60/20/20 (%vol. CNTP), T=375~ F = 1.5 l.h-1, m = 1.0 g (1) and (2) before and after the catalytic test (3) activation at 550~ (under helium)

3.2 Bismuth-nickel-molybdenum catalysts The [cxMoO3-cxNiMoO4 ] catalysts with a ratio RA=0.83 or 0.25 were chosen as references. The effect of bismuth content (precipitation) and of the water were studied. 3.2.1 Effect of bismuth content over a Ni0.17MoOx catalyst The results of Table 6 showed that the addition of a very small amount of bismuth increases significantly the selectivity towards acrolein and acrylic acid with no change in the propane conversion. The propene formation and the acetic acid production decreased at the same time, which is quite an important result of the effect of bismuth on the reaction scheme.

380 Table 6 Effect of bismuth on the propane conversion over [r Catalysts

BET 2

m .g

-1

catalysts, RA=0.83

Conv (%)

C3H8 02

Selectivity (%) C3I-I6 COx

O2/C

Zoxy etha. acro. acet. acry.

Ni0.17MoOx(*)

10

10.8

82

42

40

0.64

17

0.6

2.1

5.7

8.4

Bio.o2/Nio.lsMOOx

7

8.8

73

22

48

0.86

28

0.8

8.8

1.6

16.3

Bio.odNio.olsMOOx

8

9.6

82

21

47

0.87

29

1.3

8.4

2.8

16.6

Bio.17/Ni0.16MoOx

6

8.4

69

29

44

0.88

24

1.9

9.0

1.8

10.7

C3Hg/O2/He" 60/20/20 (%vol. CNTP), T = 425~

F = 1.5 l.h-1, m = 1.0 g, (*) T = 400~

3.2.2 Effect of bismuth and water on a [Ni0.75MoOx] catalysts Like in the preceding example adding bismuth to the ['Nio.75MoOx] catalyst increases the selectivity towards acrolein and acrylic acid and decreases the formation of acetic acid and of propene (Table 7). Table 7 Effect of bismuth and water on the propane conversion over a [czMoO3+czNiMoO4] catalyst, RA=0.25 Catalyst

1/VVH

(~

(g.h.11) C3H8 02

/ (h)

Conv (%)

Selectivity (%) C3I-I6 COx

O2/C

X;oxy etha. acro. acet. acry.

Nio.TsMoOx (29 mZ.g-1) 375/3.3

0.66

11.8

82

35

35

0.92

30

0.7

1.6

13.9

13.0

Bio.oz/Nio.7sMoO~(18 m2.g-1) 375/3.5

0.66

8.3

62

26

36

1.05

37

1.0

5.9

6.9

23.1

425/6.1

0.66

13.9

100

23

38

1.22

35

1.5

5.5

6.8

20.5

425/7.3

0.33

14.0

97

24

34

0.98

39

1.3

7.6

5.0

24.5

98

22

28

0.92

47

1.3

6.6

8.2

30.6

Substitution of HzO for He 425/10.8

0.33

14.7

C.~H~/OJX 960/20/20 (%vol. CNTP) with X = He or n 2 0 , m catalyst = 1.0 g

381 At 425~ instead of at 375~ there is an enhancement of the propane conversion but no significant change of the product distribution, except an increase of the CO2/CO ratio. In the selective oxidation of hydrocarbons to oxygenated compounds like acids, water is introduced with the reagents. This additive can create various phenomena in the oxidation : mainly kinetic effects and catalyst modification. If there is no definite evidence concerning the second point, it is well known that the addition of water favors the desorption of oxygenated compounds. In the experiment presented in Table 7 this hypothesis is well in evidence the selectivity towards oxygenated compounds particularly towards acrylic acid increasing from 25% to 31% with no change in the propane conversion. We can also notice that the total oxidation of propene to carbon oxides is inhibited by water. Under the specific conditions used in our experiment, a selectivity of about 50% towards oxygenated compounds, containing mainly acrolein and acrylic acid, was obtained for a propane conversion of about 15% and for complete conversion of oxygen.

4. CONCLUSION In agreement with previous works of Mazzocchia et coll., we have shown that the properties of nickel molybdate catalysts are greatly dependent of the molybdenum coordination and on the M o O 3 content. Propane oxydehydrogenation occurs mainly in the presence of a 13NiMoO4phase containing tetrahedral centers. When the reaction is carried out with [NiMoO4-MoO3] systems, a synergetic effect both for the activity and the selectivity is obtained when the molar ratio : otMoO3/(otMoO3+otNiMoO4) is around 0.25. According to Schrader et coll., a reciprocal coverage of the two phases (o~MoO3, otNiMoO4) can lead to changes of the surface structure. In the same way, the activity decreases after a thermal treatment at 700~ attributed to the covering of ~NiMoO4 with MOO3. A two-way reaction scheme involved for BiMo/TiO2 catalysts is also valid for bulk catalysts ( see scheme ) 02 2 ~

C3Hs

CH3CH(OH)CH3

02 > CH3COCH3

> CH3COOH + COx

3 COx J /

> C3H6

CH2CHCHO

CH2CHCOOH

Acetic acid formed via isopropanol and acetone could involve Mo 5§ species. With the second way, acrolein and acrylic acid are obtained with the participation of M o 6+ species. When the temperature increases the first way of the scheme is favored owing to a surface restructuration of the oxide: Mo+VIo3 which can contain pentacoordinated molybdenum species. The addition of bismuth to these phases decreases slightly the activity, the formation of acetic acid being supressed. When water is added to the reagent stream, the activity does not change, but the desorption of oxygenated compounds is favored and the selectivity towards acids enhanced.

382 REFERENCES

1. Y.C. Kim, W. Ueda and Y. Moro-Oka, Chem. Lett., (1989) 531. 2. Y.C. Kim, W. Ueda and Y. Moro-Oka, J. Chem. Soc. Chem. Commun., (1989) 652. 3. C. Mazzocchia, C. Aboumrad, C. Diagne, E. Tempesti, J.M. Herrmann and G. Thomas, Catal. Lett., 10 (1991) 181. 4. C. Mazzocchia, F. Di Renzo, P. Centola and R. Del Rosso, Proceeding 4th International Conference on the Chemistry and uses of Molybdenum, (1982) 406. 5. U. Ozkan and G.L. Schrader, J. Catal., 95 (1985) 120. 6. G.L. Schrader, U. Sivrioglu and M.A. Basista, Proceeding 4~ International Conference on the Chemistry and uses of Molybdenum, (1982) 415. 7. Y.S. Yoon, W. Ueda and Y. Moro-oka, Topics in Cataysis 3, (1996) 265. 8. L. Magaud, Thesis, Poitiers, France, (1994) 9. J. Barrault, L. Magaud, M. Ganne and M. Tournoux, Stud. Surf. Sci. Catal., New developments in Selective Oxidation II; Ed.: P. Ruiz and B. Delmon, Elsevier Science Publishers B.V., Amsterdam, 82 (1994) 305.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

383

Oxidative d e h y d r o g e n a t i o n o f propane on C e N i x O v (0 _< x _< 1) m i x e d oxides h y d r o g e n acceptors L. Jalowiecki-Duhamel a, A. Ponchel ~, and Y. Barbaux b aLaboratoire de Catalyse H6t6rog6ne et Homog6ne, U.R.A.C.N.R.S. D04020, Brit. C3, Universit6 des Sciences et Technologies de Lille, 59655 Villeneuve d'Ascq Cedex, France bUniversit6 d'Artois, SP 18, rue J. Souvraz, 62307 Lens Cedex, France

The oxidative dehydrogenation of propane has been studied on CeNixOy (0 < x < 1) mixed oxides previously reduced under hydrogen. The products obtained are propene and 202. No CO is obtained, whatever the temperature. An optimum propene yield of 6.9 % is obtained at 648 K on previously, at 473 K under H2, in-situ reduced CeNi0.5. Whereas without any pretreatment, a propene yield of 5.35 % is obtained. Therefore, a beneficial effect of the pretreatment under H2 is shown. The reducibility under HE of some CeNixOv mixed oxides has been previously studied and, in particular, it has been reported that these solids are able to store high quantities of hydrogen H* related to the existence of anionic vacancies. These H* species have been proposed to be one half H" species located in the anionic vacancies and the second half IT species, forming with the O 2" species of the solid, OH" groups. As dehydrogenation requires the abstraction of hydrogen from the molecule which could be performed by a lacunar phase, and by analogy to the dissociation of HE, a mechanism of the dehydrogenation step of the alkane is proposed, involving a heterolytic abstraction of a H" species by an anionic vacancy and of a IT species by an 02- species of the solid forming an OH" group.

1. INTRODUCTION Selective oxidation is one of the promising routes of utilizing alkanes which are relatively abundant in natural gas or in liquefied petroleum gas. Therefore, this research field is of great interest and of growing importance from both the industrial and the fundamental point of view. Recent studies emphasize on the importance of surface Broensted and Lewis acidity in selective oxidation of light alkanes on metal oxides [ 1,2], but the detailed mechanism and the characteristics of the active centers is still unclear. Nevertheless, it appears that it is generally admitted that i) the breakage of the C-H bond from the alkane is the rate determining step of the selective oxidation of propane [3] and ii) the reaction mechanism is of the Mars van Krevelen type (redox) [4].

384 A large variety of oxide catalysts have been claimed as beeing effective in the oxidative dehydrogenation (ODH) of propane [5-18]. Mainly vanadium based catalysts such as VPO, VMgO solids have been developed and some of them have been extensively studied in order to identify the active vanadate phases [7, 9-14]. Recently, a vanadia on precipitated silica catalyst has been found to exhibit high yield in the ODH of propane [ 17]. However, utilization of rare earth catalysts in oxidation reactions seems also to be attractive [ 18], because solids composed of oxides of Ce, Sm, Nd or Y and CeF3 are able to preserve high selectivity at high conversion [16]. It is o~en agreed that the oxidation of hydrocarbons over oxide catalysts involves surface oxygen/oxygen vacancy participation [ 19, 20], and the oxygen mobility of metal oxide catalysts has something to do with catalytic activity. The fluorite type oxides, such as ceria, zirconia and thoria, have face-centered-cubic crystal structure in which each tetravalent metal ion is surrounded by eight equivalent nearest 02- ions forming the vertices of a cube. Oxygen vacancies are created when a fluorite oxide is doped by divalent or trivalent impurity ions. Thus the fluorite oxides have been extensively studied as oxygen-ion-conducting materials due to their high oxygen vacancy concentration and mobility properties. A long time ago, a redox mechanism involving lattice oxygen/oxygen vacancy participation was proposed for carbon monoxide oxidation on cerium oxide [21]. In our laboratory, the mechanism of reduction of CeMxOv (M = Ni or Cu) mixed oxides has been studied [22]. It has been found that the insertion of Ni 2§ or Cu 2§ in ceria leads to a decrease of the reduction temperature of the host oxide and a mechanism of reduction based on the formation of anionic vacancies in ceria, facilitated by the incorporation of transition-metal cations, the heterolytic dissociation of H2 and redox reactions between Ce 4+ and the transition element, has been proposed. These rare earth based mixed oxides have been reported to be able to accept and store large quantities of hydrogen [22]. In order to elucidate further the ODH mechanism, and since dehydrogenation requires the abstraction of hydrogen from the alkane, we have studied the transformation of propane to propene on these CeNixOy compounds.

2. EXPERIMENTAL The mixed oxides CeNixOv were prepared by coprecipitation of hydroxides from mixtures of cerium and nickel nitrate using triethylamine (TEA) as precipitating agent, drying at 323 K and calcination in air at 773 K [22]. The metal loading has been verified by microanalysis. The solids will be called CeNix. The catalytic oxidation of C3H8 was performed under atmospheric pressure in a fixedbed stainless-steel tubular reactor (lenght 300 mm, internal diameter 15 mm) by co-feeding the nitrogen-diluted reaction gases (C3Hs/O2/N2 = 5/15/80). The total flow rate was 100 ml.min"~ (down flow), the reaction temperature was in the range 473-673 K, and the catalyst mass was of 0.10 g. Propene and CO2 were the only products detected and no CO was observed. The experimental details have been published previously [23]. When the oxidative dehydrogenation of C3Hs has been performed over a solid previously reduced <
385 temperatures, the catalyst has been purged under He for more than 2hrs after the reduction step. The dynamic method of titration of the hydrogen species (noted H* as the aim of the study is not their exact charge) stored by a solid has also been published previously and applied to various catalytic systems [24]. The pretreatment and catalytic experiments were carried out in-situ at atmospheric pressure in an all glass, grease free flow apparatus. The solid (65 mg) was treated first under a purified hydrogen flow at various temperatures, then after elimination of molecular hydrogen, at 423 K, under isoprene+helium flow (2 l/h) the hydrogenation activity was followed as a function of time. This hydrogenation activity involves the participation of reactive hydrogen species H* from the solid which hydrogenate isoprene and are consumed by a diffusion process. As a function of time on stream, the isoprene conversion decreases and by integrating the curve obtained, the extractable H* content that the solid is able to store has been determined. The thermogravimetric experiments were performed under purified hydrogen flow on a Sartorius balance.

3. RESULTS.

3.1. Oxidative dehydrogenation of propane. The conversions of propane to propene have been studied on the CeNix mixed oxides as a function of temperature. Appropriate blank runs showed that, under our experimental conditions, the contribution of the gas phase reaction is negligible. On Figures 1 and 2, the evolutions of the conversion and selectivity as a function of temperature obtained on CeNi0.2 and CeNi0.5 at the stationnary state are presented as examples. For the sake of comparison the catalytic activity of CeO2 was also evaluated. Propene and CO2 were the only products detected. No CO was observed whatever the temperature. On CeO2, at 373 K a propane conversion of 3% is observed with a selectivity to propene of 1.6%. As a function of temperature the conversion and selectivity increase and at 673 K, a propane conversion of 10% is obtained with a propene selectivity of 6%. On the CeNix compounds, for reaction temperatures higher than 523 K, the conversion increases while the propene selectivity decreases. On CeNi0.2 (Figure 1), propene selectivity becomes almost stable at about 10% for temperatures higher than 575 K, the yield increases up to 4.5 % for 673 K because the conversion increases. For CeNi0.5 (Figure 2), the selectivity decreases from 50 to 30 % in the temperature range 525-575 K, and an optimum yield of 5.35% is obtained at 648 K. As shown in Figure 3, when increasing the Ni content in CeNix up to x = 1, no better results are observed. One must only remark that for CeNix with x _>0.7 and temperatures higher than 675 K, CI-I4 is also observed among the products obtained. Clearly, much better performances have already been obtained on cerium based catalysts [ 16, 18] which was not the aim of the present study.

386 100

~"o<

100

8o

iI

/ :-'-,--liD,

5

-

u)

' 40

i

o

6

I

9

u)

, 40

2~

z o

o/

,

/

-3~ 9

~"

- 2~

i ip

20 0 450

,,,f.

-

I'

i

2"

dpdl

500 550 600 650 TEMPERATURE (K)

0 700

0

0 """ 450

1

500 550 600 650 TEMPERATURE (K)

O

700

Figure 2 9 Propane conversion (O), and propene yield (o) and selectivity (I), as a function of temperature on CeNi0.5.

Figure 1 9 Propane conversion (O), and propene yield (0) and selectivity (I) as a function of temperature on CeNi0.2.

t v

.

J uJ m

_

.-~.-~. t

i

;

s

9

............

_

--=

soo

.......... . ...................

~ . . .~176176 ..... :- ..........

sso 6oo TEMPERATURE

:

.

sso {K)

Figure 3 ' Propene yield as a function of temperature on CeNix with x = 0.2 (0), x = 0.5 (o), x = 0.7 ( 0 ) , x - 1 (D), and on CeO2 (o).

3.2. Ability

of the solid to accept

hydrogen

The reducibility of the CeNi0.5 mixed oxide has already been studied under H2 [22] and it has been reported that this solid is able to store high quantities of hydrogen (noted H*) in its reduced state. Taking into account that dehydrogenation requires the abstraction of hydrogen

387 species from the hydrocarbon, the ability of the solid to accept hydrogen has been studied. Therefore, the hydrogen content that the solid is able to store has been investigated by a dynamic method. Under isoprene+helium flow, at 423 K, isoprene hydrogenation activity is measured on CeNix previously treated under H2 at a given temperature. As an example on Figure 4 are presented the results obtained on reduced CeNi0.2 at 573 K under H2. The relative hydrogenation activity (HYD=t) corresponds to the ratio between the activity at time t and the activity at time zero under isoprene+helium flow. As a function of time on stream, the hydrogenation activity decreases and becomes nil. The integration of the curve obtained permits to determine the concentration of the hydrogen species H* of the solid which have reacted. Figure 5 shows the evolutions of the H* species that the solids CeNi0.2 and CeNi0.5 are able to store as a function of the treatment temperature under H2. The ability of the solid to store hydrogen depends on the treatment temperature under Hz which corresponds to various reduced states of the catalyst.

25 0.8

2O

~ O . . .

A

"O..."I3.

o'l .=...

9(16

-O

15

E

I.=

a 0.4 >"I-

/

"-o

~b ,.. 10

Q2 ~--i

0

100

*

200

Ai

300

500

600

700

800

~(rrirt) Figure 4 : Relative hydrogenation activity at 423 K under isoprene + helium flow versus time of CeNio.2 reduced at 573 K under H2.

Figure 5 9Hydrogen H* content as a function of treatment temperature under H2 of (o) CeNi0.2 and (11) CeNi0.5.

Besides, for each treatment temperature the relative hydrogenation activity can be reported as a function of the relative hydrogen H* content of the solid, and no proportionality is obtained (Figure 6). In fact, the H* consumption kinetic by the hydrocarbon is a complex phenomenon, in particular, the diffusion of the H* species from the <~bulk >>to the ~<surface >) of the solid has to be taken into account [24, 25]. For treatment temperatures higher than about 447 K, anionic vacancies are created in the CeNix catalyst by the loss of H20 (OH groups), as it is shown by thermogravimetry (Figure 3 1 , 7). After a treatment under H2 at 473 K the solid CeNi0.~ contains 17.5 10- mol.g of H

388 species, and this hydrogen storage has been correlated to the creation of anionic vacancies in the solid [22].

|

0.8 A

0.6

E E

CI

>- 0.4 -r

-8

0.2 .'

i

i

0.5

1

I.r rel.

Figure 6 : Relative hydrogenation activity at 423 K under isoprene + helium flow versus the hydrogen H* species concentration of CeNi0.5 reduced at 423 K under H2.

-12

I

273

I

473

"

TEMPERATURE (1~

673

Figure 7 9 Thermal treatment under H2 of CeNi0.5 (a) and CeNio.2 (b) followed by thermogravimetry.

A great analogy exists between the results presented in this study and those obtained in the laboratory on copper based oxides [24, 25] and molybdenum based sulfides [26, 27] which have been found to be hydrogen reservoirs. As a matter of fact different studies published on these catalysts deal with the relations existing between the active site unsaturation degree and the dienes hydrogenation activity and selectivity, as well as the existence of hydrogen H* reservoirs. It has been shown that the first hydrogen species introduced in the diene during the hydrogenation reaction is a hydride species coming from the solid [28]. Indeed, it has been proposed that under helium+alkadiene, the titrated H* species are for one half H" species located in anionic vacancies and the second halfH + species (coming from OH groups) [24-27]. These species are inserted in the solid during the activation under H2 : O 2" M n+ D + H2 -9 OH-M n§ H" (with D 9anionic vacancy). 3.3. Oxidative dehydrogenation of propane on hydrogen reservoirs. The oxidative dehydrogenation of C3H8 has been performed on previously, at different temperatures under H2, in-situ reduced catalysts. On Figures 8 and 9, propene yield and selectivity are reported as a function of propane conversion, respectively on pretreated under

389

H2 at 473 K CeNi0.2 and on pretreated under H2 at 433 K and 473 K CeNi0.5. The two treatment temperatures (433 and 473 K) lead respectively to the creation of a slight and a large hydrogen reservoir as shown previously in Figure 5. The ODH of propane obtained on the H2 pretreated CeNi0.5 at 433 K is equivalent to that obtained on the untreated solid. Besides, at 648 K, an optimum yield of about 6.9 % can be obtained on the reduced CeNi0.5 at 473 K, while at the same temperature, a maximum yield of 5.4 % is obtained on the untreated solid (Figure 8). As shown on Figure 9, a similar effect is observed on CeNi0.2. Clearly, the presence of the large hydrogen reservoir (i.e. treatment at 473 K under H2) lead to a beneficial effect on the propene yield.

100 1

6

100

8O

5

80

A

......~-.~ ......2 ~ i i 2 ~ : : 0

9

0

I

20

9

A

4

=~

3

~4o

9

5

4~ A

.

9 "'o.~

2

2O

"..~................~....

20

6

9

I

40

g=~::::::=..e ,

CONVERSION (%)

I

60

1

1 0

Figure 8 : propene yield (I, o) and selectivity (D, o) as a function of propane conversion on CeNi0.2 not treated (I, [3) and previously treated under H2 at 473 K (., o).

0

I

0

I

"

I

20 40 CONVERSION

9

I% )

I

60

Figure 9 : propene yield (m, o, 0 ) and selectivity ([2, o, 0) as a function of propane conversion on CeNi0.5 not treated (I, [2) and previously treated under H2 at 433 K ( 0 , 0) and at 473 K (., o).

The existence of the large hydrogen reservoir can be correlated to a partially reduced catalyst, as it has been published previously [22]. So, the results obtained here confirm that for the ODH of propane, the catalyst works in a partially reduced state and that a redox mechanism is involved.

4. DISCUSSION. The catalysts CeNix possess this character of being hydrogenation and oxidation catalysts in agreement with their redox properties. Knowing that in previous studies it has been shown by work function measurements that propane reacts with O 2 species located at the surface of various oxide catalysts during the activation of the alkane [23], and taking into

390 account the results obtained in the present study, a hydrocarbon activation model can be proposed. By analogy to the heterolytic dissociation of hydrogen, a heterolytic dissociation of the alkane involving the abstraction of a H" species from the hydrocarbon can be envisaged on a low coordination site involving an anionic vacancy. This active site, < OH- M "+ H" + propene (with D : anionic vacancy). It is well known that the C-H bond activation can result from different mechanisms, and even if the abstraction of a hydride species from the alkane has already been proposed in the literature on solid super acid catalysts [37], the transfert of I-I+ species has been much more otten proposed. One must remark that, in the case of a heterolytic rupture, the two species (H', I-I+) can exist, but due to its high reactivity the hydride species is much more difficult to detect. In presence of 02 the hydride species will react violently forming finally water, this permits to consume and transform 02 into selective oxygen species and regenerate the active site. 2H- + I~O 2 ~ H20 +2e" + 2 D 8 9 + 2 e " + D - ~ O2 2 H" + 02 - ~ 0 2- q- H20 + D As it has already been proposed for the hydrogen treated solid [22], the hydroxyl groups can recombine together, forming also water : 2 OH"

--~ H20 + 02- + r-]

Of course, there is no direct relationship between the H* storage and the catalytic performance. And this is true whatever the reaction (oxidation or hydrogenation), mainly because (i) the H* species storage involves <>and <<surface >>phenomena whereas the activity is related to catalytic sites located at the <<surface >>of the solid, and (ii) activity and specially selectivity for a given reaction depend on specific kind of sites. It is highly probable that the sites interacting with the alkane are much more specific (structurally for example) than

391 those interacting with H2. Moreover, the aim of the present study is not to discuss the nature and charge of the cations involved in the catalytic site. One can only surmise from the results obtained that the catalyst works in a partially reduced state. Additional characterizations are in progress, devoted to highlighting the nature and structure of the active and selective site for the ODH reaction.

5. C O N C L U S I O N A beneficial effect of pretreating CeNixOv solids under H2 is shown for the oxidative dehydrogenation of propane. Dehydrogenation requires the abstraction of hydrogen from the hydrocarbon. We have shown that, in their partially reduced state, these solids are able to accept high quantities of hydrogen H* (H', I-I+). Moreover, the hydrogen H* content that the solid is able to store depends on the reduction degree of the solid (i.e. the treatment temperature under H2), and in particular, on the creation of anionic vacancies in the solid. Therefore, by analogy to the dissociation of H2, a mechanism of the alkane dehydrogenation is proposed, involving a heterolytic abstraction of a H" species by an anionic vacancy and of a IT species by an 0 2- species of the solid forming an OH" group.

REFERENCES 1

2 3 4. 5 6. 7. 8.

9. 10. 11 12 13 14. 15 16. 17 18. 19. 20.

G. Busca, G. Centi, F. Trifiro and V. Lorenzelli, J. Phys. Chem., 90 (1986) 1337. A. Bielanski and J. Haber, in <
392 21. M. Breysse, M. Guenin, B. Claudel, H. Latreille and J. Veron, J. Catal., 28 (1972) 54. 22. G. Wrobel, C. Lamonier, A. Bennani, A. D'Huysser and A. Abouka'is, J. Chem. Soc. Farad. Trans., 92 (1996) 2001. 23. N. Boisdron, A. Monnier, L. Jalowiecki-Duhamel and Y. Barbaux, J. Chem. Soc. Farad. Trans., 91 (1995) 2899. 24. A. Sene, L. Jalowiecki-Duhamel, G. Wrobel and J. P. Bonnelle, J. Catal., 144 (1993) 544. And references therein 25. L. Jalowiecki, G. Wrobel, M. Daage and J. P. Bonnelle, J. Catal., 107 (1987) 375. 26. L. Jalowiecki, A. Aboulaz, S. Kasztelan, J. Grimblot and J. P. Bonnelle, J. Catal., 120 (1989) 108. 27. L. Jalowiecki, J. Grimblot and J. P. Bonnelle, J. Catal., 126 (1990) 101. 28. M. Daage and J. P. Bonnelle, Appl. Catal., 16 (1985) 335. 29. A. Pantazadis, A. Auroux, J.-M. Hermann and C. Mirodatos, Catal. Today, 32 (1996) 81. 30. G. Busca, G. Centi and F. Trifiro, J. Amer. Chem. Soc., 107 (1985) 7757. 31. G. Busca, G. Centi, F. Trifiro and V. Lorenzelli, J. Phys. Chem., 90 (1986) 1337. 32. G. Centi and F. Trifiro, Catal. Today, 3 (1988) 151. 33. J.C. Vedrine, J. M. Millet and J. -C. Volta, Catal. Today, 32 (1996) 115. 34. V.B. Kazansky, V. Yu. Borovkov, and L. M. Kustov, Proc. 5th International Congress on Catalysis, Berlin, 3 (1984) 3. 35. V.B. Kazansky, V. Yu. Borovkov, and A. Zaitsev, Proc. 9th International Congress on Catalysis, (Eds.M.J. Phillips and M. Ternan), Calgary, 3 (1988) 1426. 36. L. Jalowiecki-Duhamel, A. Monnier, Y. Barbaux and G. Hecquet, Catal. Today 32 (1996) 237. 37. a) H. Hatori, O.Takahashi, M. Takagi and K. Tanabe, J. Catal., 68 (1981) 132. b) O.Takahashi and H. Hatori, J. Catal., 68 (1981) 144.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

393

The Role of Adsorption in the Oxidation of ~,[3-unsaturated Aldehydes on Mo-VOxide Based Catalysts B. Stein, C. Weimer, J. Gaube* Institut far Chemische Technologie der TH Darmstadt Petersenstr. 20, D-64287 Darmstadt, Germany The kinetics of the heterogeneously catalysed vapor-phase oxidation of ~,[3-unsaturated aldehydes to the co, [3-unsaturated acid has been investigated for a Mo-V-Cu-oxid catalyst. The reaction rates of aldehyde and oxygen consumption as function of the aldehyd, oxygen, acid and water partial pressure are described by a kinetic model based on a modified Mars-van Krevelen mechanism. Also the rate of the c~,[3-unsaturated acid oxidation has been measured and described for varied partial pressures of acid, oxygen and water. It turned out that the rate of consumption of ~, [3-unsaturated aldehyde is not affected by the concentration of c~, [3-unsaturated acid and vice versa. With increasing ~, [3-unsaturated aldehyde concentration the rate of aldehyde and oxygen consumption passes a maximum. The decreasing reaction rate at elevated aldehyde concentration is interpreted by adsorptive blocking of reduced sites of the catalyst by aldehyde so that the rate of reoxidation is reduced. The increase of the rate of aldehyde oxidation by co-feeding of water vapor could be interpreted by an enhanced rate of catalyst reoxidation which has been found also in the absence of aldehyde and the related acid. The selectivity towards ~, [3-unsaturated acid is mainly determined by the consecutive oxidation of ~, [3-unsaturated acid while the parallel reaction of c~, [3-unsaturated aldehyde to by-products is almost negligible. 1. INTRODUCTION Acrylic acid is produced in a two-stage process, the oxidation of propene on catalysts of the type Mo-Bi-Fe-oxides to acrolein [1,2,3] and the oxidation of acrolein on catalysts usually composed of Mo-V-Cu-oxides [4]. Because of the moderate number of reaction products and the high selectivity towards acrylic acid the oxidation of acrolein is an appropriate model reaction to study partial oxidations. Therefore, several investigations have been carried out for elucidation of the function of catalyst components. Particularly, the interest was focussed on the dependence of activity and selectivity on the ratio of Mo 6+ / V 4+ [5,6,7]. Kinetic and infrared studies have lead to the development of a mechanism including declarations about specific interactions of adsorbed species with catalyst components such as Mo and V [8]. Infrared studies revealed that adsorbed acrylate anion is a relatively stable intermediate [5,9]. The reaction of this anion to adsorbed acrylic acid by hydroxyl groups of the catalyst surface and the desorption of acrylic acid are regarded as rate determining steps of the acrolein oxidation [ 10]. The well kown increase of the reaction rate by addition of water to the reaction gas [7,11,12] is interpreted by the enhancement

* To whom correspondence should be adressed

394 of the acrylate anion protonation caused by the increased number of hydroxyl groups [131. The aim of the presented work is a more detailed kinetic study particularly focussed on the role of the adsorption of the reactants.

2. EXPERIMENTAL

2.1. Preparation and characterization of the catalyst A typical catalyst for the acrolein oxidation composed mainly of Mo-, V-, Cu-oxides was prepared according to the patent specification EP 17000. In order to avoid the influence of mass transfer processes on the rate of conversion pellets of egg shell type with a thin active layer of about 200 ~tm thickness were used for the kinetic measurements. The BET surface of the calcined active material in the state of oxidation is about 18 m2/g which increases considerably as the degree of reduction is raised. The active mass shows macropores with a maximum at 0.14 ~tm. REM micrographs show crystallites with diameters in the range of 0.1 - 2.5 pm. 2.2. Experimental unit A flow diagram of the experimental unit is shown in figure 1a. Dosing of acrolein (ACR) and acrylic acid (AA) was performed by use of microprecision pumps. In order to prevent these monomers from polymerisation and oligomerization the pumps were kept at low temperature. Dosing of water was carried out by saturation of a nitrogen flow. Acrolein and acrylic acid are injected through a nozzle into a hot gas stream (nitrogen, water vapor) in order to instantaneously reach a temperature of 200 ~ The gas stream composed of nitrogen, water, oxygen, acrolein and acrylic acid was conducted to the differential recycle reactor, realized by a jet loop reactor [ 14], which behaves like a CSTR. Propane, used as internal standard for the GC analysis, was precisely dosed to the gas at the outlet of the reactor. A small part of the gas leaving the reactor passed a UNIVAP precision gas sampling system to the on-line GC. The main flow passed a cooling trap and finally the continuous analysis of oxygen, carbon monoxid and carbon dioxide. The reaction rates are defined and calculated as follows: Acrolein consumption t:tACR rAC R

-

9o nAC R

= m CAT

The rate of oxygen consumption is calculated from the oxygen balance, including the rate of water formation which is obtained by the hydrogen balance.

395

to2

1

,0

[ 1/2 (

-

,0

rIAC R -

riAC R ) + ( nAA

-

ti4 A ) -

ti A -

1/2 riCO

-

riCO 2 ] +

m CAT 1

[ (riAci~

~

-

o

~

nAc R)

+ (,ia~

-

o

+ ri A ]

n~)

m CAT

The formation of the by-products acetic acid (A), CO and COz is normalized to C3-equivalent. ti

-Ft

~

m CAr

propane N2

c

9 i

3

=

a c e t i c acid, C O , C O 2

=

number of C-atoms of by-product

acrolein acrylic acid 02 H20 co co2

evaporation & mixture

02

I I

N2

acrolein saturation

1

i

vacuum

jet loop reactor

,,,.--.

acrylic acid

/1 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

data registration PC

{ Go-F,o I condensat~

......................... I 02/00/002 [ ~ continuous analysis

la Figure 1

adsorption cell lb

Experimental units: 1a. Setup for kinetic measurements lb. Setup for adsorption and instationary kinetic measurements

Since the reaction kinetics is of the Langmuir-Hinshelwood type additional informations about the adsorption of reactants is helpful in the development of a kinetic model. Therefore experiments have been carried out by use of an experimental setup for adsorption measurements. Since adsorption and reaction occur simultaneously the change of gas composition in the adsorption cell is registrated by on-line GC-MS. Then the amount of adsorbed substances is calculated as the difference of added substance and substance in the gas phase of the adsorption cell given by the total pressure in the cell and the gas composition. A schematic drawing of the

396 experimental unit is shown in figure 1b. This apparatus has been also used for instationary kinetic measurements.

3. RESULTS AND DISCUSSION The dependence of the rate of acrolein consumption on oxygen pressure, figure 2, shows the typical behaviour of an oxidation which follows the Mars- van Krevelen mechanism. In the range of low oxygen pressure the reaction rate is determined by the reoxidation of the catalyst. At high oxygen pressure the rate of conversion is nearly independent of oxygen pressure, indicating the oxidation of acrolein respectively the reduction of the catalyst as rate determining step. Therefore, a kinetic model has been developed on the basis of this well accepted Mars-van Krevelen mechanism for the reaction network of acrolein oxidation.

rlS /

acrylic acid

r2

acrolein

=- by-products (CO, CO 2, acetic acid)

The catalyst is reduced by the reactants acrolein and acrylic acid and the reduced catalyst is reoxidized by oxygen from the gas phase. For the steady state of reoxidation and oxygen transfer the following set of kinetic equations results. -rl = kl PACR ~1

acrolein oxidation

-r2 = k2 PACR ~1 nl

-m

-ro21 = k021 PO 2 PACR (1

-

~1)

~2

-r3 = / % P ~

acrylic acid oxidation

n2

-r022

= k022

p o 2 (1 - (~2)

The reaction rates -rAcR, r AAand -ro2 are obtained from the experiments, v i is the st6chiometric coefficient of oxygen consumption via reaction i (v 1- 89 -rAc R = r 1 + r 2 rAA = r I - r 3 _

r02

=

_

( to21

+

r02 2)

=

v I

r 1

+

v 2

r2

+

v 3

r3

397

The parameters (I)1 and ~2 characterize the oxidation state of the catalyst. Because of the different reaction rates r 1 + r 2 and r 3 also different degrees of oxidation are assumed, which does not necessarily mean that specific sites for the oxidation of acrolein and acrylic acid exist. 12

4.0 kPa..._..._...-~--~

~(HO) = 10kPa

10. T= 280 ~ p = 200 kPa 2.6 kPa____.__,___

8 -~6

1.4 kPa

9

2O,,ff

0

1

5

,

1'0

,

1

15

'

1

2o

'

~

P(O21 [kPa]

,

30

'

/

35

'

Figure 2: Oxidation of acrolein-rAc R = f (Po2) " parameterpAcR 9solid lines: calculated The best respresentation of the experimental data shown in the following diagrams was obtained with the asumption of first order with respect to both acrolein (r 1 and r2) and acrylic acid (r3). In accordance with results of instationary reoxidation experiments employing the same catalyst the reaction order with respect to oxygen was set to 1.3 for the reoxidation term r(O2) ~ and first order for r(O2)2. 10

,c(O~=~ ~

8,

6. ~

n

n.....-n- U ~ u

2O

o

.

.

~

~

~

.

p(~Q = 10kF~

10-

ii ~m~(~

T= 3CI?~ p=2COkPa

j o

1.3 I~:~i

8-

,___.

6-

.

Figure 3 Oxidation ofacrolein 9 -racR = f (PAA) " parameter PACR

o / O ~

acrolein consumpt!on

4-

u~

~, ~ p(AA) [kPa]

~ ~

p(AA)= 62 kPa T= 300~ p =300k P a ~

0{,

~

parallel & consecutive oxidation

o

~

,

~ p(AC~ [kPa]

.

Figure 4 Oxidation of acrolein and acrylic acid; -r,~

= f (p~)

rby-pr. = f

(PAcR)

9

(C3-equivalents)

398

1214 t

12-

~S-o~2230k~

37~/2~

'

/:(O2)=m kPa

19kPa

T=2D~ p=330kPa

~10"1

~

10kPa

i ~4- ~

~ ~ ~

6I~a

~

O~

0

.

2

.

.

4

.

p(/~

6

.

8

01

10

[kPa]

Figure 5 Oxidation ofacrolein 9 -r = f (PAcR) " parameter P(O2) " solid lines: calculated

o

.

2~

5

.

.

~

.

& .

.

~

p(Aa~ [kPa]

40

Figure 6 Oxidation of acrolein ; -r = f ( P A c R ) " parameter p(HzO ) ; solid lines: calculated

The reaction rate of acrolein consumption is independent of the concentration of acrylic acid, figure 3. At constant concentration of acrylic acid the rate of by-product formation remains almost constant even though the concentration of acrolein is drastically raised, figure 4. This leads to the conclusion that the parallel formation of by-products is negligible and the consecutive oxidation of acrylic acid is independent of the partial pressure of acrolein. With increasing acrolein concentration the rate of acrolein and oxygen consumption passes a maximum, figure 5. The decreasing reaction rate at elevated acrolein concentration suggests adsorptive blocking of reduced sites of the catalyst by acrolein or an oxidized intermediate.

Po(O2)=9kPa T=280~ I ~

10-

t--

8-

~ ,

60

u

4

IV[)-V-~id

addition of water

wate.r~~t ~~ 1

in absence'of

2 0

'

I

4

'

I

3

'

I ~ I

2 decjeed r e c t u m

[%]

I

1

,I~

Figure 7 Effect of water on reoxidation of catalyst

~"

0

399 The reaction rate of acrolein and oxygen consumption increases considerably as the concentration of water vapor is raised, figure 6. This well kown effect has been interpreted by an enhanced desorption of acrylic acid causing an increase of oxygen access to the catalyst surface and finally reoxidation of the reduced catalyst [15]. However instationary reoxidation experiments in absence of acrolein have revealed that addition of water vapor also increases the rate of reoxidation of the partly reduced catalyst, figure 7. The same rate as that calculated for the acrolein oxidation under steady state condition is obtained for a degree of reduction in the range of 2 - 4 %. Since, in the presence of acrolein a reduced rate of reoxidation is assumed for the steady state of acrolein oxidation a higher degree of reduction is expected. The degree of reduction is related to the total oxygen content of the catalyst. The oxidation of acrylic acid also follows the reaction mechanism of Mars-van Krevelen, figure 8. In contrast to the acrolein oxidation no hindrance of reoxidation is observed when the partial pressure of acrylic acid is raised. Addition of water vapor causes only a small increase of the rate of acrylic acid oxidation also in contrast to the considerable effect in acrolein oxidation, figure 9.

1,0

p (I-120)= 10 kPa 0,8 T=300~ 0,6

0

~.......3~

0,2 0,0

T= 300~

_ /,,"/" " .

,/~ 3.5 kPa

.

.

.

p= 200kPa

Varialionp(H20)

37 kPa

g o.4 ,

,o(o~) = 2o kPa

VarialJon p(

p=2OOkPa

.

p(AA) [kPa] Figure 8: Oxidation of acrylic acid ; -r = f (PAA); parameter p(O2) ; solid lines: calculated

'

10

9

9 2 kPa

e/~

0

I

0

10 kPa

9 20 kPa

/

T

2

~

4

6

8

p(AA) [kPa]

Figure 9: Oxidation of acrylic acid ; -r = f (pAn); parameter P(H20) ; solid lines: calculated

Finally the selectivity towards acrylic acid of acrolein oxidation is given for constant concentration of acrylic acid, figure 10. The selectivity is defined as:

rAA

r1

-

r3

- rAC R

r1

+

r2

SAA -

400

100 I Variak~nT / 95-

260~

AJ

, ~

~

~ ,

90- 280 CO

p(u~q= 1 o ~ p(o) =m~a ~(AA)= 3 ~

85

p=23014~a 00"oC

80

0

,

1

,

2

,

:3

zl

~i

6

p (/~__,R)[kl~] Figure 10 Selectivity towards acrylic acid 9SAA = f

(PAcR) "

parameter T

Adsorption measurements have shown negligible adsorption of CO and COz, and moderate adsorption of water vapor. In order to separate the adsorption of acrolein and the following oxidation instationary experiments were carried out at 230 ~ The concentration of acrolein in the gasphase of the adsorption cell shows a fast fall attributed to adsorption followed by acrolein consumption and an increase of the concentration of reaction products, figure 11. The reaction product water is nearly completely adsorbed. The broken line indicates the adsorbed amount of substance calculated as C3-equivalents. The diagram shows a strong adsorption of acrylic acid or acrylate anion. Related to the surface which is evaluated from BET (N/) data the adsorption layer of C3-equivalents is very dense. 2,0

T = 230~ Po = 7 kPa

l

4-" 1,5

IVb-V-~id, R0= 2% R = 2,4%

/

~ 1,0~qukrolent \ acrolein \ o -.7.0,5.

o

0,0~ t~ ' ' ' o 0

"~o

~

r

20

o------o - - - - - o o acblicadd CO CO vvat"~ "--<

8~~~~40

~

60

o

,

80

t [min] Figure 11 Instationary oxidation of acrolein in absence of oxygen ; c i = f (t) Rx = degree of reduction (x- o: t = 0 min ; x - e: t = 80 min)

401 4. CONCLUSIONS The kinetics of acrolein oxidation is of first order with respect to acrolein, which can be easily understood by the competition of strongly adsorbed acrylic acid and less strongly adsorbed acrolein. The reduced rate of acrolein oxidation at elevated acrolein concentration is interpreted by adsorption of acrolein on reduced sites and consequently described by a negative reaction order with respect to acrolein in the reoxidation t e r m r(O2) 1. In this context the independence of the acrolein oxidation rate from acrylic acid concentration is surprising. This result may be understood by the reaction mechanism proposed by T.V. Andmshkevich in which formed acrylate anions are shifted to sites on vanadium cations and after protonation are desorbed as acrylic acid. Parallel to the protonation the reduced sites are reoxidized. The experiments of acrolein oxidation in absence of molecular oxygen confirm a dense adsorption layer of oxidized C3-compounds, acrylic anion + acrylic acid, figure 11. It is remarkable that despite of the increasing adsorption of the oxidized C3-compounds the rate of acrolein consumption is hardly reduced. IR analysis of surface intermediates have lead to the assumption that deprotonation of chemisorbed acrolein and oxidation of the intermediate to acrylic anion are fast reactions. Therefore, the protonation of the acrylic anion and the following desorption of acrylic acid are regarded as rate determining steps leading to accumulation of the acrylic anion on the surface of the catalyst [12]. Both the rate of acrolein oxidation and the rate of catalyst reoxidation depend on the concentration of acrolein so that also steps towards the formation of acrylic anion must be taken into account as rate determining. In the whole the acrolein oxidation appears as the result of a network of coordinated reaction steps. The influence of water on the rate of acrolein respectively oxygen consumption is still included in the reaction rate constant of acrolein oxidation k 1 and in the constant of catalyst reoxidation k(Oa)l. These constants are listed in the table for different water vapor contents of the reaction gas. p(H20) P,o,al = 200 kPa

kl*

k(02),*

260 ~

280 ~

300 ~

260 ~

280 ~

300 ~

1%

0.33

0.78

0.90

0.31

0.19

0.55

5%

1.00

1.00

1.00

1.00

1.00

1.00

10%

1.20

1.01

1.01

1.40

1.49

2.17

* For comparsion,the constants are related to the values at a water vapor content of 5 %.

While at a temperature of 280 ~ and above the constant of acrolein oxidation k 1 is hearly independent of the water content, the reaction rate constant of the reoxidation k(O2)~ strongly increases with increasing water content. In absence of acrolein the reoxidation rate of the reduced catalyst also increases markedly as the water vapor concentration is raised, figure 7. This effect may partly explain the

402 considerable increase of the rate of acrolein oxidation in continuous operation. However, also enhanced protonation of the acrylic anion, enhanced desorption of acrylic acid and the liberation of reduced sites from adsorbed acrolein may be taken into account to explain the rate increasing effect of water. On the other hand the interpreted results support in the main the mechanism proposed by T.V. Andrushkevich and on the other hand reveal complicated relations of competitive and noncompetitive adsorptions and reactions.

REFERENCES 1.

J.L. Callahan, R.K. Grasselli, E.C. Milberger and H.A. Strecker, I&EC Prod. Res. and Dev., 9 (1970) 134.

2.

R.K. Grasselli and J.D. Burrington, Adv. Catal., Acad. Press, 30 (1981) 133.

3.

R.K. Grasselli, J. Chem. Ed., 63 (1986) 216.

4.

N. Nojiri, Y. Sakai, Y. Watanabe, Catal. Rev.- Sci. Eng., 37 (1995) 145.

5.

T. Andrushkevich, L. Playsova, T. Kuznetsova, V. Bondareva, T. Groshkova, !. Oienkova, N. Lebedeva, React. Kinet. Catal. Lett., 12 (1979)463.

6.

T. Kuznetsova, G. Boreskov, T. Andrushkevich, Y. Grigorkina, N. Maksimov, I. Olenkova, L. Plyasova, T. Gorshkova, Reakt. Kinet. Catal. Lett., 19 (1982)405.

7.

T. Kuznetsova, T. Andrushkevich, T. Groshkova, React. Kinet. Catal. Lett., 30 (1986) 149.

8.

G. Popova, A. Davydov, I. Zakharov, T. Andrushkevich, Kinet. Katal., 23 (1982) 692.

9.

A. Davydov, G. Popova, T. Andrushkevich, React. Kinet. Catal. Lett., 25 (1984) 1175.

10. E. Erenburg, T. Andrushkevich, V. Bibin, React.. Kinet. Catal. Lett., 41 (1990) 33. 11. J. Tichy, J. Kf~shka, J. Machek, Coll. Czech. Chem. Comm., 48 (1983) 698. 12. R. Recknagel, L. Riekert, Chem. Technik, 46 (1994) 324. 13. E. Erenburg, T. Andrushkevich, G. Popova, A. Davydov, V. Bondareva, React. Kinet. Catal. Lett., 12 (1979) 5. 14. G. Luft, H.-A. Herbertz, Chem.-Ing.-Techn., 41 (1969)667. 15. T.V. Andrushkevich, Catal. Rev.- Sci. Eng., 35 (1993) 213.

Acknowledgement Financial support by the Bundesministerium fiar Bildung und Forschung (BMBF) is gratefully acknowledged.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

403

A n e w catalyst for propane ammoxidation: the S n N / S b m i x e d oxide S. Albonetti a, G. Blanchard b, P. Burattin b, S. Masetti a and F. Trifir6 a aDepartment of Industrial Chemistry and Materials, University of Bologna, V.le Risorgimento 4, 40136 Bologna, Italy bRhone Poulenc Chimie, Rue De La Haie Coq 52, 93308 Aubervilliers, France.

The effect of the preparation and calcination method of a Sn/V/Sb mixed oxide on its catalytic performance are analyzed with respect to ammoxidation of propane. Different samples have been prepared with coprecipitation technique (by dissolving the starting materials with ethanol, iso-butanol or water), or by solid state reaction between oxide and tin hydroxide. The sample prepared by ethanol shows the best catalytic performance: this solvent makes the partial alkoxides quite stable in solution and brings about a better coprecipitation with a more intimate mixture of tin, antimony and vanadium. The thermal transformation of the precursor of this sample has been followed during the calcination, both in air and in nitrogen with several characterization techniques. The thermal treatment in air at 700~ leads to the best catalytic performance, that is, good activity and high selectivity. This calcination procedure leads to a homogenous mixed oxide, with numerous well-crystallized microfields of tin oxide, incorporating antimony and vanadium ions in limited concentrations and an excess of poorly crystallized antimony oxide. 1. INTRODUCTION Acrylonitrile (ACN) is produced nowadays by the ammoxidation of propene on catalysts made of promoted Fe-Bi-Mo-O (SOHIO) or promoted Fe-Sb-O (Nitto); nevertheless in recent years some companies have decided to invest in the ammoxidation of propane research [ 1]. One of the more interesting catalytic system for the direct ammoxidation of the paraffm is Sb/V/O [2-7]. The synthesis of these catalysts is usually made by solid state reaction between V20 5 and Sb20 4 at 700~ or by reaction between NH4VO 3 and Sb20 3. Mossbauer analysis shows that antimony is mainly in its pentavalent state, so a large amount of vanadium is reduced, with the formation of VIIISbVO4 and the possible presence of V IV in substitutional or interstitial solid solution coordinations. Centi et al. [3] and Nilsson et al. [4] have shown that an excess of antimony brings about an activity decrease, but also a large increase in the selectivities and in the yields in acrylonitrile and propene. The best catalyst for the synthesis of acrylonitrile from propane has a large excess of antimony (Sb/V=5.0) [6]. This excess seems to quicken the transformation of the intermediate propene to acrylonitrile; on the contrary, an excess of vanadium brings about an activity increase, but also a low selectivity in acrylonitrile due to the increased production of propene and carbon oxides. Sn/Sb/O system [7] has been widely studied in recent years as a catalyst active in the allylic oxidation and ammoxidation: the best preparation method for these compounds involves coprecipitation from a solution of Sn(IV) and Sb(V) chlorides [8, 9]; calcination at temperatures higher than 700-900~ (depending on the bulk antimony/tin ratio) provokes

404 segregation on the surface of ~-Sb204 particles [10]; an antimony surface enrichment, exspecially for a concentration of antimony lower than 20-30%, [11] and the formation of Sb(III) [9] are evidenced during the thermal treatment; the catalytic activity strongly depends on the Sn/Sb ratio and on the calcination temperature [10, 12, 13]. In the literature there is also a disagreement about the hypothesis on the nature of the active sites for allylic oxidation on this system: isolated Sb 3+ species surrounded entirely by Sn4+ ions in a specific environment [9, 10]; an oriented film of Sb204 supported on a solid solution of Sb 5+ in Sn(IV) oxide [8, 13]; a solid solution of Sb 5+ ions in a rutile type matrix [14]; two "gem" SbS+=o groups [15]. In order to make a new catalyst for the ammoxidation of acrylonitrile, we decided to mix these two systems: Sb/V, in order to activate the paraffin, and Sn/Sb, having already been utilized to ammoxidate the relative olefin. In this paper, the effects of the method of preparation and calcination of a Sn-V-Sb mixed oxide on its catalytic performance are analysed. In particular we changed the medium in which the starting material were dissolved, utilizing ethanol, iso-butanol and water, in order to understand the interaction between the solvent and the metallic ions and to enhance the performance of the catalyst. Moreover we made a mixture of oxides and hydroxide in n-hexane to compare the coprecipitation with a simple mixture. The catalyst with the best performance was then calcined at different temperatures and in different atmospheres in order to optimize the thermal treatment.

2. EXPERIMENTAL 2.1. Catalyst preparation Every sample was prepared with a relative molar ratio between Sn:V:Sb of 1:0.2:1 (this ratio was optimized in a previous work of this research group). Coprecipitation of vanadium, antimony and tin oxohydrates was achieved as follow: initially a solution of anhydrous SnC14 in an organic medium (absolute ethyl alcohol for sample 1 and iso-butanol for sample 3) or in an acidic aqueous medium (around HC1 3M for sample 2) was prepared; then VO(acac) 2 and SbC15 were dissolved, in this order, in the solution, in order to obtain the desired Sn/V/Sb ratio. The sequence of the dissolution of the salts is very important in order to obtain an homogenous solution. The VO(acac) 2 complex hydrolyzes in a stepwise fashion: hydrolysis of the intermediate VO(acac) + is much slower than the bis-chelate complex at room temperature. When anhydrous SnC14 is dissolved in the organic medium, strongly acid conditions are developed, that are required for the rapid cleavage of both the vanadyl diketonate ligands; this allowed a complete hydrolysis of the stable VO(acac) 2. In addition, the solvolysis effect produced during the dilution of SnC14 in the organic medium further enhanced the hydrolysis. This solution was added dropwise to an aqueous solution of CH3COONH 4, having an initial pH of around 7.0. During the precipitation of the oxohydrates the pH, that decreases due to the release of HC1, was mantained constant by the addition of ammonia solution. In fact, the pH must be not higher than 7.5 (in order to avoid the V 4+ oxidation and to allow the precipitation of antimony oxohydrated) and not lower than 5 in order to avoid the vanadium and tin oxohydrates redissolution. When the precipitation was carried out at pH outside the mentioned range, there is no simultaneous precipitation of the catalyst components. The resulting precipitate was filtered, washed and dried overnight at 120~ then it was calcined at 350 ~ for 1 hour and at 700~ for 3 hours. Sample 4 was prepared by a solid state reaction between V205, Sb20 3 and a tin oxohydrated, prepared by precipitating an ethanolic solution of SnC14 in a pH controlled aqueous solution (similar to the preparation of the sample 1). The precipitate of Sn(OH) 4 was filtered, washed and dried at 120~ overnight (drying at 100-140~ leads to the partial

405 dehydratation of ortho-stannic acid Sn(OH)4 to meta-stannic acid SnO(OH)4). It was then mixed with V20 5 and Sb20 3 by suspending the powders in vigorously stirred n-hexane. After evaporation of the solvent under reduced pressure, the mixture was dried at 120~ overnight and finally calcined at 350 ~ for 1 hour and at 700~ for 3 hours. Sample 1, which is the catalyst with the best performance, was re-prepared and, after the drying treatment, was then calcined up to different temperature in air or in nitrogen in order to study the influence of the thermal treatment.

2.2. Characterization The catalysts were tested in a conventional laboratory apparatus with a tubular fixed-bed (length 10"; diameter 1/4"), working at atmospheric pressure and at 360-480~ The analysis of propane, propene, acrylonitrile, acetonitdle and uncondensable gases were made by gaschromatographic techniques (with Poropack QS and Carbosieve packed columns); ammonia and hydrogen cyanide was detected by absorptions and titrations. The composition of the feed was propane 25%, oxygen 20%, ammonia 10% and the remainder helium. The catalyst (2 ml) was loaded as grains (30-40 mesh) and the contact time was around 2 s. A thermocouple, placed in the middle of the catalyst bed, was used to verify that the axial temperature profile was within 3~ Surface areas were determined using the B.E.T. method with nitrogen absorption at -196~ on a Carlo Erba instrument. Fourier transformed infrared (FT-IR) spectra in transmission were recorded using a PerkinElmer 7200 Fourier transform spectrometer and KBr disk technique. X-ray diffraction patterns (powder technique) were obtained using Ni-filtered CuKtx radiation (~,=1.542/~) with a Philips computer controlled instrument (PW 1.050/81). XPS analysis was detected with a Perldn Elmer Model Philips 5500. 3. RESULTS AND DISCUSSION

3.1. Effects of the preparation method 3.1.1. Surface area Surface areas of the four sample are reported in Table 1. The catalysts have a similar value of surface area; the one prepared from oxides and hydroxides (4) shows the lowest value. Table 1. Sample prepared with different method and surface area Sample Medium Temperature (~ Atmosphere 1 EtOH 700 Air 2 H20 700 Air 3 i-BuOH 700 Air 4 n-hexane 700 Air

Asup (m2/g) 46 54 62 32

3.1.2. FT-IR spectroscopy FT-IR spectra of the four samples are reported in Figure 1. The catalysts prepared by coprecipitation (1,2,3) have the same FTIR spectra. It is possible to see the stretching absorbance of the V=O bond at around 988 cm-l: this shifting to lower frequencies with respect to the absorbance of the crystalline V20 5 (1022 cm-1) may be attributed to two effects, firstly to the interactions between the vanadyl species and the ruffle type SnO2 matrix and

406

11 l /

/

1000 800 600 cm-1400 Fig. 1. FTIR of the samples 1-2-3 and 4

10 ' 20 ' 30 ' ,~0 ' 50 ' 60 2070 Fig. 2. XRD of the samples 1, 2, 3 and 4

secondly to the electronic effect due to the presence of neighboring reduced vanadium sites. The absorbance at 629 cm -1 can be attributed to the stretching of the Sn-O bond: this spectra is different from the tin oxide one (660 and 620 cm -1) due to the interaction of tin with antimony and vanadium. The spectra of sample 4, prepared by mixture of oxide and hydroxide in hexane, is very different from the other ones: it shows the absorbance bands of ~-Sb20 4 at 745, 648, 605 and 529 cm -1 and a weak one at 1020 cm -1, typical of crystalline V20 5.

3.1.3. X-ray diffraction The diffractograms of these four samples are reported in Figure 2. The samples prepared by coprecipitation methodology are very similar and show clearly the diffraction lines of the rutile type SnO 2 phase; the reflections are very broad, due to the very small crystallites. No other phases are present, however, it can be observed that, due to the broadness of the diffraction lines of SnO 2, the possible presence of the ruille type VSbO 4 phase would be difficult to evidence. When the preparation is carried out by solid state reaction (sample 4), the XRD patterns drastically change, as it was observed in the FT-IR characterization. The main phase present is o~-Sb20 4, moreover it is possible to see the weak peaks of SnO 2.

3.1.4. Catalytic tests In Figures 3 and 4 are reported respectively conversion vs. temperature and acrylonitrile selectivity vs. conversion for these four samples. It is possible to see, that even if the catalysts reach similar conversions, they have different activities, due to their different working temperature range. The acrylonitrile selectivity is very different: sample 1 (EtOH) is the best

407 70

25

E (9

20 -

4:n-hexane~/ . ~ / 3:i-BuO~,J~

/

~15 09 > c" O

o10

,

/

~50

.. isoyield "-.. 10%

30-

T

03

-

40 -

(3O

O

1:EtOH",

60-

/

<

50

360

'

i

380

i

i

1

400 420 440 Temperature, ~

i

460

Fig. 3. Conversion vs. temperature for the samples 1, 2, 3 and 4

480

10

0

-...

4:n-hexane

20-

'

5

'

'

'

10 15 20 C3H8 conversion, %

isoyield - - _ 5%

'

25

30

Fig. 4. ACN selectivity vs. conversion for the samples 1, 2, 3 and 4

one, reaching value of 60%, then the order decreases as follows, sample 3 (i-BuOH), sample 2 (water) and sample 4 (n-hexane). Hence, while the catalytic data of the samples prepared by coprecipitation technique are different, the structural characterization has not revealed m e a n i n ~ differences. The use of ethanol as solvent for the dissolution of the ions brings about the partial formation of tin-, antimony- and vanadyl-alkoxide, which slows down the precipitation of the different compounds, making their respective velocities uniform. In this way a better interaction among the hydroxides is achieved and the better dispersion of the active sites brings about the necessary multifunctionality to convert ~ saturated hydrocarbon into a functionalized molecule with a good selectivity. The alkoxides are more difficult to obtain with a sterically hindered alcohol as iso-butanol and similar equilibria doesn't exist with water. Hence, by using water or iso-butanol as solvent, the precipitation velocities of the different compounds can be different, leading to a worse dispersion of the active sites; that means low acrylonitrile selectivity. The method of preparation by solid state reaction (sample 4) leads to a mixture of oxides and not to a mixed oxide: in fact the characterization has revealed a segregation of the system into the single oxides. Consequently the mixture of the active sites is not good and the acrylonitrile selectivity falls considerably. 3.2. Effects of the thermal treatment

Sample 1 (Ethanol) was re-prepared and calcined at 500, 700 and 800~ nitrogen, as reported in Table 2.

in air and in

3.2.1. Surface area

Table 2. Sample prepared with differente thermal treatment and surface area Sample Medium Temperature (~ Atmosphere 1A500 EtOH 500 Air 1A700 EtOH 700 Air IA800 EtOH 800 Air 1N500 EtOH 500 Nitrogen 1N700 EtOH 700 Nitrogen 1N800 EtOH 800 Nitrogen

Asup (m2/g) 114 46 32 74 26 18

408 The surface areas of the sample calcined at different temperature, both in air and in nitrogen are summarized in Table 2. The surface area is reduced with an increase of the calcination temperature, due to the destruction of the microporosity; moreover the sample calcined in air has a surface area higher with respect to the one calcined at the same temperature in nitrogen. 3.2.2. FT-IR spectroseopy The FTIR spectra of samples calcined at high temperature, both in air and in nitrogen are shown in Figure 5 (1A700, 1A800, 1N700 and 1N800). The FT-IR spectra of the sample calcined in air are quite similar, it is possible to see the typical absorbance bands of the tin oxide (around 629 cm -1) and of vanadyl species (around 988 cm -1) described in the previous section. After treatment at higher temperatures, with an increase in crystallinity and the progressive evolution of volatile compounds, the band, typical of Sn-O bond in the Sn/Sb system, becomes more well defined. Moreover, some initial modifications of spectra due to Sb-O-Sb bond of antimony oxides are observed in the sample calcined at 800~ The band typical of vanadyl species is much less intense in the spectra of samples calcined in nitrogen. Moreover, in nitrogen atmosphere, the typical bonds of antimony oxides are observed, even after treatment at 700 ~ indicating a higher amount of segregation of different phase during calcination in inert gas. Hence, the evolution of phases with temperature, as evidenced using FT-IR, strongly depends on the atmosphere of calcination.

L,, o ra~

.< ~

~

~

1N700

,~~ 1000

800

600 cm-1400

Fig. 5. FT-IR of the samples 1A700, 1A800, 1N700 and 1N800

10

20

~ 30

40

50

60 20 70

Fig. 6. XRD of the samples 1A700, 1A800 and 1N700

409

3.2.3. X-ray diffraction The XRD of the samples calcined in air at 700~ and 800~ and in nitrogen at 700~ are reported in Figure 6 (respectively 1A700, 1A800 and 1N700) In the XRD of the sample calcined at 700~ it is possible to see only the reflections of the rutile type SnO2; the diffractogram of the sample calcined in air at 500~ is similar, even if less crystalline. Moreover, it can be observed that, due to the broadness of the diffraction lines, the possible presence of the rutile VSbO4 phase would be difficult to observe. On the contrary, the XRD of the sample calcined in air at 800~ is very different: both SnO2 and Sb20 4 phases (13Sb20 4 and ~-Sb204) and an increase in the crystallinity of the samples, as evidenced by the narrowing of the diffraction lines and lowering of the signal background, are obtained at 800~ The formation of ~-Sb20 4 under these conditions was not expected, since direct oxidation of antimony oxide produces ~-Sb204, which is converted in the 13-Sb204 at 1130~ Berry [9] proposed that the 13-Sb204 phase obtained during the preparation of the VSbO4 phase was more accurately described as a solid solution of vanadium in 13-Sb204 and that the vanadium facilitates and stabilises the low temperature formation of this phase. Thus, also in ours case, the formation of ~-Sb20 4 at low temperature can be favoured by the migration of some vanadium ions to form a solid solution within this phase. In spite of the high content of Sb in the catalyst, the X-ray diffraction patterns up to 700~ show only the lines due to rutile-type SnO2. No lines due to antimony oxides are observed, but the diffraction lines of rutile phase in all the SnN/Sb samples are shifted towards lower d values with respect to those of pure SnO2. This observation can be attributed to the substitution by Sb and/or V ions in the tin(IV) oxide lattice. Table 3 shows the rutile tetragonal cell volume for samples calcined at different temperatures in air. In these samples the cell volume increases, with increasing the temperature of calcination, reaching values closer to (but still lower than) that of pure SnO2. This behavior indicates a substitution of Sn4+ by ions with atomic radius lower than 0.69/~ (V3+=0.64/~; V4+=0.58~; VS+=0.54A; Sb3+=0.76~; Sb5+=0.60/~), this substitution should be higher at lower temperature and it seems to decrease with increasing the calcination temperature. Moreover, since it is known that an antimony concentration as low as 5-6% may indicate the upper limit of antimony incorporation in the tin(IV) oxide lattice [9], it is probable that the excess of antimony can be present as an amorphous phase together with the ruffle doped structure. Table 3. Cell parameter and volume of rutile SnO9 of the sample calcined in air and in nitrogen. Sample a (/~) c (~) v (/~3) 1A700 4.70 3.14 69.36 1A800 4.70 3.16 69.83 1N700 4.70 3.18 70.24 1N800 4.70 3.17 69.97 Pure SnO 2 4.74 3.19 71.55

In Figure 6 is also shown the X-ray diffraction profde of sample calcined in nitrogen at 700~ (1N700). The profde of a sample calcined at 600~ shows the SnO2 phase only (like sample calcined in air at 700~ on the contrary, in nitrogen atmosphere, the formation of crystalline antimony oxides is shifted to a lower temperature with respect to in air (700~ instead of 800~ and the main phase evidenced, besides SnO2, is only a-Sb204 and not 13Sb204.

410 Table 3 shows the effect of the calcination temperature in nitrogen on the crystallographic parameters. Also in this case, the cell volume of the Sn/V/Sb system is lower than pure SnO2, but the values are different from samples calcined in air, evidencing a different kind of substitution into the rutile structure. Since the cell volume obtained by calcination in air is lower, we can hypothesize that, during calcination in nitrogen, a lower amount of antimony and vanadium ions enter into the rutile structure to form a solid solution; thus, the presence of oxygen during calcination seems to favour the formation of a solid solution.

3.2.4. X-ray photoelectron spectroscopy The material, calcined at different temperatures and under different atmospheres was studied by XPS analysis in order to determine the surface composition and the oxidation state of tin, vanadium and antimony in each of the samples. Table 4 reports the atomic percentage of tin, vanadium and antimony in the various catalysts: the bulk atomic ratio analysis is made by atomic absorption and the surface atomic ratio by XPS spectroscopy. Table 4. Bulk and surface relative atomic ratio of the sample calcined in air Sample Atomic ratio Sn:V:Sb Bulk Surface 1A500 5.0:1.0:4.9 5.0:1.0:4.8 1A700 5.0:1.0:4.9 5.0:1.2:5.9 1A800 5.0:1.0:4.9 5.0:1.2:3.7 1N500 5.0:1.0:4.8 5.0:0.8:4.7 1N700 5.0:1.0:4.8 5.0:0.8:4.5

The results show that surface vanadium content in catalysts calcined in air is quite close to bulk composition. Moreover, the Vxp S intensity remains almost constant increasing the temperature of calcination. On the contrary, the Sb surface composition is similar to bulk composition only at low temperature and we have observed an enhancement of the signal up to 700~ and a decrease after calcination at 800~ in air. This last phenomena can be due to: i) migration of antimony into the ruille structure, ii) aggregation of antimony particles. The latter is more acceptable because XRD analysis of the sample calcined at 800~ in air shows the segregation of Sb20 4 phase. In a nitrogen environment a much less pronounced effect of temperature is observed, but the vanadium surface composition of these samples is lower than bulk, indicating a probable aggregation of vanadium particles. XPS analysis was also utilized to estimate the oxidation state of the different elements present in the system. Table 5 shows the oxidation degree for vanadium and antimony ions; the oxidation state of tin is always Sn4+. Table 5. Estimation of antimony and vanadium oxidation state in the SnN/Sb system. Atomic percentuage (%) Sample Antimony Sb 3d 5/2 Vanadium 2p Sb3+ SbS+ V4+ V5+ 1A500 4 96 33 67 1A700 5 95 39 61 1A800 5 95 44 56 1N500 2 98 51 49 1N700 3 97 51 49

411 The XPS analysis revealed the presence of Sb 3+, Sb 5+, V 4+ and V 5+ species in oxygen environments. In all the samples, the surface antimony is mainly present as Sb 5+ and the maximum content of Sb 3+ is 5%. The surface vanadium is present both as V 5+ and V4+ and the amount of V 4+ present on air calcined samples increases with increasing the temperature of calcination. This behavior can be due to the formation of same stable structure containing V4+ also if the content of surface vanadium measured by XPS has not showed any decrease due to segregation of new phase. On the contrary, the Vxp S signals of samples calcined at different temperature in nitrogen remain unchanged and show a higher degree of reduction of these samples with respect to the samples calcined in air.

3.2.5. Catalytie tests In Figures 7 and 8 are reported respectively conversion vs. temperature and acrylonitrile selectivity vs. conversion for the six samples calcined at different temperature, both in air and in nitrogen. By comparing these data with the structural characterization reported before, the following conclusion can be pointed out. The activity of the samples calcined at 500~ is quite high, but the selectivity in acrylonitfile is very low, owing to combustion. The high surface area of the sample calcined at low temperature negatively influences the catalytic performance. The sample calcined in air at 800~ is less selective in acrylonitfile with respect to one calcined at 700~ this decrease is caused by the structural evolution of the system with segregation of the antimony oxide. The sample calcined in nitrogen at 700~ is less active with respect to one calcined in air at the same temperature and the selectivity in acrylonitrile is very low, owing to combustion. The absence of vanadyl species, evidenced by FTIR spectra, seems to influence the catalytic performance of the mixed oxide in a dramatically negative way. By increasing the temperature of the calcination in nitrogen atmosphere (from 700 to 800~ the sample becomes more active and selective in acrylonitrile, but still less with respect to the catalyst calcined in air at 700~ The presence of different phases and of a nonhomogenously distributed mixed oxide leads to a less selective sample for propane ammoxidation.

25

70 1A800

o~ 20 -

o~

E

",

60

1 "

.O m

~ 1A800

"-

~15> cO

isoyield

~, 4o

1

~

o]0-

T

<

CO

O

50

380

20 i

t

i

i

i

400

420

440

460

480

500

Temperature, "C

Fig.7. Conversion vs. temperature for the samples calcined in air and in nitrogen

10 0

1 N7

5

-. 10% "-..

0 1N500 Z . " " - -

isoyield

. . . . . . . U_'

- A,y00 I

10 15 20 C3H8 conversion, %

i

25

30

Fig. 8. ACN selectivity vs. conversion for the samples calcined in air and in nitrogen

412 4. CONCLUSION The general conclusion that can be drawn from this work is that our original method of preparation using a coprecipitation procedure from the ethanolic phase [16] leads to a more intimate mixture of tin, antimony and vanadium, inducing a spreading of the active component that promotes active site isolation [17] and stabilizes the active phase. During the thermal treatment the catalytic system evolves from a homogeneous mixed oxide, where numerous well-crystallized microfields of SnO 2 incorporating antimony and vanadium ions in a limited concentration are dispersed in an excess of poorly crystallized SbO x, to a mixture of antimony oxide and ruffle type solid solution. In fact, when the calcination temperature is increased up to 800~ (in air) or 700~ (in nitrogen) the antimony ions migrate forming a segregated antimony oxide. The catalytic performance for the ammoxidation of propane appears to be related to an enhanced antimony content in the surface (evidenced by XPS analysis) in poorly crystalline materials, i.e. samples in which the antimony in excess is highly dispersed and amorphous. High temperatures of calcination cause the migration and phase segregation of antimony to the surface and a corresponding decrease in the selectivity to acrylonitrile formation. REFERENCE

1. 2. 3. 4.

5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17.

G. Centi, R.K. Grasselli and F. Trifirb, Catal. Today, 13 (1992) 661. G. Centi, D. Pesheva and F. Trifirb, Appl. Catal., 33 (1987) 343. G. Centi, F. Trifirb and R.K. Grasselli, La Chimica & L'Industria (Milan) 72 (1990) 617. R. Nilsson, T. Lindblad, A. Andersson, C. Song and S. Hansen in: New Developments in Selective Oxidation II, eds. V. Cortes Corberan and S. Vic Bellon (Elsevier, Amsterdam) 82 (1994) 293. G. Centi, R.K. Grasselli, E. Patan6 and F. Trifirb in: New Developments in Selective Oxidation, eds. G. Centi and F. Trifirb (Elsevier, Amsterdam) 55 (1990) 515. A.T. Guttmann, R.K. Grasselli, J.F. Bradzil and D.D. Suresh, U.S. Patent 4,746,641 (1988), assigned to The Standard Oil Co.. G. Centi and F. Trifirb, Cat. Rev. - Sci. Eng. 28 (1986) 165. J.C. Volta, P. Bussiere, G. Coudurier, J.M. Herrmann and J.C. Vedrine, Proc. IX IberoAm. Symp. Catal. (Lisbon, 1984) p. 888. F.J. Berry, Adv. Catal. 30 (1981) 97. H.J. Hemiman, D.R. Pyke and R. Reid, J. Catal. 58 (1979) 68. Y.M. Cross and D.R. Pyke, J. Catal., 58 (1979) 61. F. Trifirb and I. Pasquon, La Chimica & L'Industria (Milan) 52 (1970) 228. Y. Boudeville, F. Figueras, M. Forissier, J.L. Portefaix and J.C. Vedrine, J. Catal., 58 (1979) 52. G.M. Godin, C.C. McCain and E.A. Porter, Proc. 4th Int. Congr. Catal., vol. 1, Akademiai Kiado, Budapest 1971, p. 271. F. Sala and F. Trifirb, J. Catal. 41 (1976) 1. S. Albonetti, G. Blanchard, P. Burattin, F. Cavani and F. Tdflrb, France Patent 94-079821994, assigned to Rhone Poulenc Chimie. J.L. CaUahan and R.K. Grasselli, AIChE J. 9 (1963) 755.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

413

F o r m a t i o n o f a c t i v e p h a s e s in the S b - V - , A 1 - S b - V - , a n d A 1 - S b - V - W - o x i d e systems for propane ammoxidation Jerker Nilsson a, Angel R. Landa-C~inovas b, Staffan Hansen c and Arne Andersson a aDepartment of Chemical Engineering II, Lund University, Chemical Center, P.O. Box 124, S-221 00 Lund, Sweden bCentro de Microscopia Electronica, Universidad Complutense de Madrid, Ciudad Universitaria, E-28040 Madrid, Spain CNational Center for HREM, Inorganic Chemistry 2, Lund University, Chemical Center, P.O. Box 124, S-221 00 Lund, Sweden The Sb-V-O, A1-Sb-V-O and A1-Sb-V-W-O systems were investigated for the ammoxidation of propane to acrylonitrile. Compared with pure =SbVO4, the =SbVO4 that was obtained by cocalcination of pure =SbVO4 with o~-Sb204 at 800 ~ and subsequent sieving was found to be much more selective to acrylonitrile. XRD data showed no variation of the futile unit cell for synthesis ratios Sb:V > 1, indicating that the active phase is =SbVO4 with a surface enriched with Sb. Characterisation of syntheses in the A1-Sb-V-O system allowed the identification of a trirutile-like phase with the composition All_xSbVxO4 (0 < x < 0.5). The synthesis of this phase, which is active and selective to acrylonitrile, requires excess of aluminium. Over a fresh preparation with A1, Sb, V and W the activity and the selectivity to acrylonitrile increased considerably with time-on-stream. This behaviour shows that the active rutile phase is formed in situ, and EDX analyses gave the average composition A10.1Sb0.8V0.7W0.404. The highest yield to acrylonitrile that was observed for the three systems was 37 % and was obtained over the A1Sb-V-W-oxide. 1. I N T R O D U C T I O N In patents [ 1] it has been indicated that for propane ammoxidation =SbVO4 is an active phase of rutile type in catalysts with Sb, V and W on an alumina-rich support. It has been found that when the amount of antimony in the synthesis is in excess of the stoichiometric ratio, i.e. with Sb:V > 1, a catalyst is produced which is more selective than the pure =SbVO4 phase for acrylonitrile formation. An Sb:V ratio in the range 2-10 was reported as being optimal [1-3]. Over Sb-V-O catalysts with excess Sb and without A1 and W the yield achievable to acrylonitrile is 10-12 %, to be compared with 3-4 % for a corresponding sample with an Sb:V ratio equal to one. It was shown by adding alumina to SbsVOx that the yield to acrylonitrile was further increased up to about 25 %, i.e. by a factor more than two. Another key element in the catalysts of rutile type is tungsten. A composition Sb5VWOx supported on A1203 with 20 wt.% of SiO2 was reported to give a yield to acrylonitrile of 40 % at almost 70 % propane conversion [ 1]. Considering the yield to acrylonitrile, the data show a remarkable step-wise improvement when progressively adding excess Sb, A1 and W to --SbVO4. In the present paper the formation and structure of the active phases in the three systems Sb-V-O, A1-Sb-V-O and A1-Sb-V-W-O will be described and compared considering various characterisation results. It will be shown that there are significant differences between the systems concerning the composition of the active rutile phase, the stage at which the active structure is formed, and catalytic performance.

414 Moreover, it will be demonstrated that data can be rationalised in terms of the site isolation principle, which has been found applicable in other catalyst systems with antimony [4].

2. EXPERIMENTAL 2.1. Preparation of samples =SbVO4 (2.0 m2/g) was prepared by heating an equimolar mixture of V205 (Riedel-deHahn, p.a.) and Sb203 (Merck, p.a.) in air at 800 ~ for 18 hours. The same method was used to prepare a sample with excess antimony oxide (Sb:V = 2), having a specific surface area of 1.4 mZ/g. Another Sb-V-O sample with Sb:V = 2:1 was prepared from a slurry of Sb203 in water solution with NH4VO3 (Merck, p.a.), which was heated under reflux before drying and final calcination at 610 ~ [3]. The specific surface area was 3.6 m2/g. For investigation of phase formation in the A1-Sb-V-O system, weighed amounts of AI(OH)3 (Riedel-de-Ha~n, p.a.), Sb203 and V205 were mixed and ground together and finally heated at 680 ~ for 4 days with one intermittent grinding. A sample, which was used as catalyst, and with the atomic ratio AI:Sb:V = 21:5:1 was prepared from the same reactants following a slurry method which has been described in detail elsewhere [5]. The final calcination was performed at 610 oC and the resulting specific surface area of the sample was 157 m2/g. A catalyst with the nominal composition AI:Sb:V:W = 21:5:1:1 was prepared according to the slurry method [1,5] starting from AI(OH)3, Sb203, V205 and (NH4)6W12(OH)2038 (Strem Chemicals, p.a.). The specific surface area of the sample was 121 m2/g after the final calcination at 610 ~

2.2. Activity measurements The activity measurements were performed at 480 ~ using a plug-flow reactor made from glass. Dilution of the catalyst with quartz grains was necessary to have isothermal conditions. The composition of the inlet flow to the reactor was propane 14.3 vol.%, ammonia 14.3 vol.%, oxygen 28.6 vol.%, water vapour 7.1 vol.%, and nitrogen 35.7 vol.%, corresponding to the stoichiometric ratio between propane, ammonia and oxygen for acrylonitrile formation. Propane and the products CO, CO2, propylene, acrylonitrile and acetonitrile were analysed on a GC equipped with a Porapak Q column, a sample valve, a methanisation column for conversion of the carbon oxides to methane, and an FID detector. Analyses of the conversion of ammonia and the formation of HCN were performed using titrimetric methods [6]. Analysis of conversions and selectivities were performed with time-on-stream after heating the catalyst in air to the reaction temperature and subsequently switching to the reactant composition. The influence of propane conversion on the product distribution was studied varying the amount of catalyst at constant flow rate.

2.3. Characterisations Specific surface areas were measured with a Micromeritics Flowsorb 2300 instrument, using adsorption of N2 at liquid N2 temperature. The samples were degassed at 350 ~ For X-ray powder diffraction (XRD), the samples were crushed and mounted on adhesive tape. Films were recorded using a Guinier-H~igg focusing camera with CuKo~I radiation and with Si as internal standard. Energy dispersive X-ray microanalysis (EDX) was carried out in a transmission electron microscope JEM-2000FX fitted with a Link AN10000 analysis system. The phases were first identified by electron diffraction and thin edges were then analysed using a beam approximately 500/k in diameter and an acceleration voltage of 200 kV. High resolution transmission electron microscopy was performed inoa JEM-4000EX instrument operated at 400 kV and possessing a structural resolution of 1.6 A. Samples were lightly ground in methanol and the dispersion was then transferred to copper grids covered with a holey carbon film. In the microscope, thin crystals positioned over the holes in the carbon film were examined by diffraction and imaging techniques.

415 3. R E S U L T S

AND DISCUSSION

3.1. T h e S b - V - O system Stoichiometric SbVO4 does not form when heating an equimolar mixture of V205 and Sb203 under various conditions, but two [7] or three [8] compositional series with SbVO4 as an end member were used to characterise the rutile-type phases that were produced. M6ssbauer data indicated Sb 5+ and consequently mixed V3+/V 4+ in this rutile [7]. By heating equimolar mixtures of V205 and Sb203 at 800 ~ in flowing gas with various O2/N2 ratios, we have recently demonstrated that the data can be interpreted in terms of a single cation deficient series, namely Sb0.9V0.9+xD0.2-xO4 (0 < x < 0.2), where D denotes cation vacancies [9]. In Fig. 1 are the electron diffraction pattems of selected samples from the series. The pattern of the oxidised end member Sb(V)o.9V(IV)o.sV(III)o.IDo.204 (Fig. lc) shows beside the basic rutile lattice, satellite diffraction spots which are due to ordering of the cation vacancies. An approximate supercell 2~/2a,2,4T2b,4c was observed for a partial oxygen pressure equal to or higher than that for air. For the reduced end member Sb(V)o.9V(IV)o.2V(III)o.9Do.oO4 the diffraction pattern in Fig. 1d shows a different rutile superstructure (~/2a, a/2b,2c), resulting from cation ordering [10]. At intermediate pressures of oxygen (0.1 < P(O2) < 0.21 atm) only diffraction spots from basic rutile are observed (Fig. l a). When the synthesis was carried out with excess of vanadia, we observed [9] formation of a solid solution series Sb0.9VI.IO4 - VO2, i.e. Sb0.9-yV1 l+yO4 with 0 < y < 0.7. The V-rich end composition, which is approaching V(IV)O2, becomes Sb(V)0.2V(IV)I.6V(III)0.204. The diffraction pattern of the V-rich structure showed a basic rutile lattice with strong diffuse intensity between Bragg spots. Investigation of the Sb-V-oxide by XRD showed only reflections from the basic rutile cell.

Figure 1. Electron diffraction patterns of rutile structures recorded with the electron beam parallel to the [100] direction. Superlattice reflections are indicated by arrows. (a): intermediate --SbVO4; (b): (A1,V)SbO4; (c): oxidised =SbVO4; and (d): reduced =SbVO4.

416 Figure 2 shows catalytic data for =SbVO4 and a sample with the atomic ratio Sb:V = 2:1, which both were prepared by solid state synthesis. XRD showed the latter sample to consist of =SbVO4 and Sb204. The plots for =SbVO4 of the propane conversion and the selectivities to propylene, acrylonitrile and C 1-2 degradation products against time-on-stream show no variation with time. At =5 % of propane conversion the selectivities to propylene and acrylonitrile are 54 % and 9 %, respectively. The sample prepared with Sb:V = 2:1 shows at the same conversion level a modest decrease in propane conversion with time-on-stream. Simultaneously the selectivity to acrylonitrile increases from 22 % up to 26 %, while the selectivities to propylene and degradation products correspondingly decrease. The increase in the selectivity to acrylonitrile can possibly be associated with our previously reported finding that migration of antimony from Sb204 to the surface of =SbVO4 occurs during the catalytic reaction [3]. We have, moreover, reported that samples with excess of antimony are slightly less reduced than the pure =SbVO4 after use in the ammoxidation [3,11]. However, these findings cannot fully explain the differences between the data for =SbVO4 and the Sb:V = 2:1 sample in Fig. 2, since in fact the data show that the sample with excess antimony is already initially much more selective than the pure =SbVO4 to acrylonitrile formation. In contrast to previous results [ 12], the XRD data in Table 1 for the samples show no difference in the dimension of the unit cell of the rutile, which for both agrees with the composition Sb0.9Vo.904. 70

70

"~

60

v

60

"= o

50

~= o

50

m

40

o~

40

"o

tO

~ t-

oO

9

9

~-

e.-

30

30 " k'~L--~__ ~ A A A .

o m

'-

20

ID tO

10

..

..

--

==

o

=.

20

10 o----o

0 0

i 1

i 2

i 3

i 4

i 5

T i m e - o n - s t r e a m (hours)

i 6

0

7

0

I

i

i

I

I

I

I

1

2

3

4

5

6

7

Time-on-stream

(hours)

8

Figure 2. Propane conversion (O) and the selectivity to propylene (@), acrylonitrile (11) and degradation products (A) vs. time-on-stream over (a): Sbo.9Vo.904; and (b): solid state preparation with the atomic ratio Sb:V = 2:1. Table 1 Unit cell parameters for the =SbVO4 in solid state preparations with two nominal Sb:V ratios c/a Sb:V atomic ratio a (/k) c (/k) 1:1" 1:1, fresh 1:1, used 2:1, fresh 2:1, used

4.625(4) 4.6228(5) 4.6247(4) 4.6228(8) 4.6251(12)

3.040(2) 3.0385(4) 3.0397(4) 3.0387(6) 3.0369(9)

0.657 0.657 0.657 0.657 0.657

*The average unit cell as determined for one sample using four different methods; from ref. [ 13].

417 Thus, the apparent effect of excess antimony in Sb-V-O preparations to enhance the selectivity to acrylonitrile cannot be explained by a change of the bulk composition of the =SbVO4 phase. To gain further insight into this matter a mechanical mixture (Sb:V - 4:1) with pure --SbVO4 and o~-Sb204 was heated for 8 hours at 800 ~ The fractions of particle size were 0.150 - 0.425 mm and 0.100 - 0.150 mm for =SbVO4 and ~-Sb204, respectively, allowing separation of the two phases by sieving after the co-calcination. Activity data are presented in Figure 3 for the mechanical mixture after the heat treatment and for untreated =SbVO4, not calcined with o~-Sb204. We have previously reported that a simple mechanical mixture of the two phases, which was not submitted to the co-calcination step, does not show improved catalytic performance as compared with the =SbVO4 phase [ 11]. The present data in Fig. 3 for the co-calcined mixture, on the contrary, show considerable improvement. The selectivity to acrylonitrile at --24 % propane conversion is only =2 % over =SbVO4, while it is substantially higher (42 %) for the calcined mixture. Consequently, calcination of -SbVO4 with (~-Sb204 has produced a catalyst which gives less degradation and is more selective to C3 products. Data obtained for each phase from the mixture, after separation by sieving, are as well included in Fig. 3. These data show that the co-calcined =SbVO4 phase alone is responsible for the high selectivity to acrylonitrile, though it is less active than the untreated =SbVO4. Although it is known that both o~-Sb204 and ~-Sb204 have low activity and are not selective to acrylonitrile formation [5], the co-calcined Sb204 phase has become active and selective. XRD data indicate that the reason can be the o~-Sb204 during the co-calcination has transformed partly into Vdoped Sb204 [14], which has been reported to be selective to acrylonitrile formation [5]. Comparison of the data for the co-calcined mixture with those measured for each of its constituents shows that the propane conversions over the latter samples are almost addable, and the data also indicate that the propylene formed over the V-doped Sb204 is converted to acrylonitrile over the =SbVO4. When the co-calcined phases after separation are used in consecutive beds the corresponding data in Fig. 3 show almost no difference as compared with the phases are being used in form of a mixture. This fact indeed shows that the activated =SbVO4 is a major phase, determining the catalytic performance in Sb-V-O catalysts with excess antimony oxide. ~. 100 v

~>9

80

-~

60

"0 c

~

['9 [~ 9 I-]

sel. propylene sel. acrylonitrile propane conversion N H 3 conversion

40

c

o

~ c 0

o

20

0

I

a

i

b

[

c

d

i

e

f

Figure 3. Conversion of propane and ammonia, and the selectivity to propylene and acrylonitrile in propane ammoxidation over (a): =SbVO4; (b): mixture obtained by co-calcination of particles of =SbVO4 with c~-Sb204; (c): =SbVO4 obtained by co-calcination with ff-Sb204 and subsequent separation by sieving; (d): V-doped Sb204 obtained by co-calcination of =SbVO4 with otSb204 and subsequent separation by sieving; (e): co-calcined samples in two serial beds with =SbVO4 in the first and V-doped Sb204 in the second bed; and (f): co-calcined samples in two serial beds with V-doped Sb204 in the first and -SbVO4 in the second bed. The co-calcination was performed at 800 ~ for 8 hours. Amounts: =SbVO4 1.0 g, and Sb204 1.9 g.

418 The data described conclusively demonstrate that the active phase in the antimony-rich side of the Sb-V-O system is formed during the calcination step in the catalyst synthesis. The active phase is -SbVO4, possibly with a surface which is enriched with antimony. In the present work the enrichment can be due to extraction of vanadium by the Sb204 phase, which is in contact with the =SbVO4 surface. However, at lower calcination temperatures where no V-doped Sb204 can form, the enrichment with antimony possibly is due to migration of antimony species from o;-Sb204 to the surface of =SbVO4. For such inference we have previously presented support [3,15]. The present results can clearly be rationalised in terms of isolation of the vanadium centres at the surface is needed to a suitable degree to have a surface which is selective to acrylonitrile, agreeing with the idea on site isolation that originally was presented by Callahan and Grasselli [16]. For a Sb-V-O catalyst prepared by the slurry method [3] and with the atomic ratio Sb:V = 2:1, the selectivity to propylene as a function of the propane conversion is plotted in Fig. 4a, and the corresponding plots of the selectivity and the yield to acrylonitrile are in Fig. 4b. The data are representative for Sb-V-O preparations. The highest yield to acrylonitrile obtained per single pass is 11% and is achieved at about 40 % of propane conversion. The corresponding selectivity passes through a maximum of 37 % at lower conversion (25 %). --

sel. p r o p y l e n e S b - V - O

-"

- -m-- sel. p r o p y l e n e AI-Sb-V-O - o- - sel. p r o p y l e n e A I - S b - V - W - O 60 , , , , ,

I II

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I i,

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o

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o

0

20 40 60 80 C o n v e r s i o n of p r o p a n e (%)

. . n - - y i e l d acn. A I - S b - V - O - o - -yield acn. A I - S b - V - W - O i

6o ~ 40 "8 g

~ 30

0

b t

~ 40 ".~,

20

i

._. 5O

-1z

0

..

--a---yield acn. S b - V - O

80

a

.;

4o

sel. acn. Sb-V-O

- - m - - s e l . acn. AI-Sb-V-O - o - - s e l . acn. A I - S b - V - W - O

~9

,_~

_

"~

20

0

o)

10 0

.o

~

""

"O

..

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i

-/-I~

,' 0

~ 20

40

60

80

100

Conversion of propane (%)

Figure 4. (a): The selectivity to propylene vs. the propane conversion in propane ammoxidation over Sb-V-O with Sb:V = 2:1, A1-Sb-V-O with AI:Sb:V = 21:5:1, and A1-Sb-V-W-O with AI:Sb:V:W = 21:5:1:1. The conversion of propane at =70 % ammonia conversion is indicated for Sb-V-O (A), A1-Sb-V-O (!-3), and A1-Sb-V-W-O (O). (b): The selectivity and the yield to acrylonitrile over the same samples as in (a). 3.2. The AI-Sb-u system Syntheses in the A1-Sb-V-O system were performed starting with mixtures of AI(OH)3, Sb203 and V205 in various ratios, which were calcined in air at 680 ~ [5]. Characterisation with XRD, electron diffraction and EDX revealed the formation of crystalline 8-A1203, o~Sb204, V205 and A1VO4, as well as two rutile related phases Sbo.9Vo.904 and (A1,Sb,V)204. No crystalline A1SbO4 was observed to form at 680 ~ since its formation requires higher synthesis temperature [ 17]. The 8-A1203phase, which was undetectable with XRD, consisted of 5-10 nm sized crystals and was identified using electron diffraction and EDX. The XRD analyses showed that oc-Sb204, A1VO4 and V205 are formed close to their ideal stoichiometry, though due to some decomposition of A1VO4, V205 appears also in the Al-rich side of the system. In Fig. 5 are the fields of formation for Sbo.9V0.904 and (A1,Sb,V)204 depicted for

419 different starting compositions. Formation of Sb0.9Vo.904 occurs over a wide range of starting compositions, while (A1,Sb,V)204 appears as product exclusively in the Al-rich part of the system. o

o

§

/.: " v v ok*

v v-vv o\O

§

/ 20,.s

O%Sb

.\0

.\o

0\0

ok*

Figure 5. Fields of formation in air at 680 ~ of Sb0.9V0.904 (left triangle) and (A1,Sb,V)204 (fight triangle) as determined by X-ray diffraction. The circles are marked at the starting composition (at.%) of the synthesis. Three samples synthesised under various conditions and containing (A1,Sb,V)204 were selected for elemental analysis. Crystals of (A1,Sb,V)204 in each sample were first identified by electron diffraction and were then analysed using EDX. The data are plotted in Fig. 6. Sb204

SbV

V204

bO4

AI203

Figure 6. EDX analysis of ruffle-related (A1,Sb,V)204 crystals in three representative samples, which were prepared by calcination of mixtures with AI(OH)3, Sb203 and V205 in various ratios. O: starting composition AI:Sb:V = 9:9:2, calcined at 900 ~ for 6 days; O: starting composition AI:Sb:V = 6:3:1, calcined at 680 ~ for 4 days; and +: starting composition AI:Sb:V = 1:2:1, calcined at 1000 ~ for 4 days. In spite of a slight overestimation of the antimony content, the plot shows that a solid solution series between A1SbO4 and SbVO4 exists (A13+ ~ V 3+) and can be formulated as All_xSbVxO4 with 0 < x < 0.5. The series evidently does not correspond to substitution of both A13+ and Sb 5+ with V 4+, i.e. a solid solution between A1SbO4 and V204. The gap in the

420 composition between =SbVO4, or more precisely Sbo.9Vo.904, and A105SbV0504 is reasonable, since the former phase is essentially a V 4+ compound, while the "latter is a V 3+ compound. Comparison of the field of formation for (A1,Sb,V)204 in Fig. 5 with its range of composition in Fig. 6 shows that at 680 ~ (A1,Sb,V)204 is not formed in syntheses which start from compositions corresponding to the phase composition. It appears that the formation of Sb0.9V0.904 is kinetically favoured and that a large excess of alumina (> 60 at.%) is required to obtain (A1,Sb,V)204 as the sole rutile structure. (A1,Sb,V)204 likewise Sbo.9Vo.904 has a rutile-related structure and for both phases only the basic rutile-type reflections are observed with powder X-ray diffraction. Selected area electron diffraction, on the other hand, admits the easy distinction between (A1,Sb,V)204 and Sb0.9V0.904 (see Fig. 1). The diffraction pattern in Fig. l b of (A1,Sb,V)204, or (A1,V)SbO4, shows a 3-fold supercell of trirutile-type [ 18]. We have reported that (A1,Sb,V)204 is the active phase for propane ammoxidation in Al-rich A1-Sb-V-O catalysts [5]. Figure 7 shows for a slurry preparation with AI:Sb:V = 21:5:1 the dependence of the catalytic performance with time-on-stream. The propane conversion and the selectivities show almost constant behaviour, indicating that the active structure is formed in the synthesis of the catalyst. Characterisation with FTIR, FT-Raman, XPS and X-ray diffraction before and after use in ammoxidation showed no difference [5]. The selectivities to propylene and acrylonitrile together with the yield to acrylonitrile are shown in Fig. 4 as a function of the conversion of propane. At about 30 % propane conversion, the selectivity to acrylonitrile passes through a maximum of 45 %. The yield to acrylonitrile reaches 20 % at around 55 % propane conversion and 75 % ammonia conversion. Further increase is limited by the complete consumption of the oxygen in the stream. 60

i

o-----o

!

o

r f~

OU

1

~

I

i

I

I

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>,

._> 45

o

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-

t-

o

$

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> cO

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O 0

0

I

I

I

2

i

3

Time-on-stream (hours)

Figure 7. Propane conversion (O) and the selectivity to propylene (O) and acrylonitrile ( I ) vs. time-on-stream over a slurry preparation with the ratio AI:Sb:V = 21:5" 1. 3.3.

2O

t-

(..)

The

AI-Sb-V-W-O

0 ~

0

10

20

Time-on-stream

30

40

(hours)

Figure 8. Propane conversion (O) and the selectivity to propylene (O) and acrylonitrile ( I ) vs. time-on-stream over a slurry preparation with the ratio AI:Sb:V:W - 21"5:1" 1.

System

No detailed characterisation of this catalyst system regarding phase composition has previously been reported in the literature, though kinetic data have been given [2,6]. A freshly charged catalyst with the nominal composition AI:Sb:V:W = 21:5:1:1 showed a considerable change in catalytic behaviour with time-on-stream. Figure 8 shows that both the activity and the selectivity to acrylonitrile increase during the first 15 hours, after which the selectivity to acrylonitrile reaches 44 % at about 70 % of propane conversion. On the freshly charged catalyst, initially, mainly degradation products are formed.

421 Characterisation of the used sample with transmission electron microscopy combined with electron diffraction and EDX revealed the presence of a rutile phase with A1, Sb, V and W. This phase is present in form of polycrystalline aggregates consisting of very small crystallites, which are less than 10 nm in diameter, see Fig 9. Eight EDX analyses showed the average composition A10.1Sbo.8V0.7W0.404, which is quite different from the nominal metal composition. The polycrystalline aggregates were observed to be much less frequent in the freshly prepared sample. It appears that the activation behaviour is due to reduction in propane ammoxidation of W 6+ to form W 4+, which can accommodate in a structure of rutile-type. Considering the composition of the rutile with tungsten, it seems that it can be described as a solid solution between (V,A1)SbO4 and WO2, where two W 4+ replaces one Sb 5+ and one V3+/A13+.

Figure 9. Electron micrograph of polycrystalline (A1,Sb,V,W)204 (left), and the corresponding electron diffraction pattern with basic rutile tings (fight). Besides the rutile phase, XRD and electron diffraction of the flesh catalyst showed the presence of ff-Sb204 and WO3, while 8-A1203 was identified by electron diffraction only. The same phases were present in the used catalyst, though the tungsten oxide phase appeared less frequent and showed defects. The selectivity to propylene together with the selectivity and the yield to acrylonitfile are shown in Fig. 4 as a function of the conversion of propane. Compared with the Sb-V-O and A1Sb-V-O samples, the A1-Sb-V-W-O preparation is more selective to C3 useful products at high propane conversions (> 40 %). The selectivity to acrylonitrile reaches almost 50 % at 55 % propane conversion. The yield to acrylonitrile becomes almost 37 % at 80 % propane conversion, where the consumption of ammonia is complete. The results presented for the A1-Sb-V-W-O system show that in this system the active phase is formed in situ in propane ammoxidation. The improvement of the catalytic properties, compared with the less complex subsystems, can be considered due to the introduction of tungsten in the rutile, adjusting the structure and properties of the active vanadium ensemble. 4. C O N C L U S I O N S In the Sb-V-O system the structure that is active for propane ammoxidation is formed in the catalyst synthesis. It is comprised of --SbVO4 with a ruffle superstructure and a surface which is enriched in antimony. The formation of the active surface requires an excess of antimony in the synthesis (calcination), usually Sb:V - 2-5 [1-3,15], as compared with the equimolar ratio for forming Sbo.9Vo.904.

422 The active structure in the A1-Sb-V-O system is a bulk phase, which is directly formed in the catalyst synthesis. It is a trirutile-like phase All_xSbVxO4 (0 < x < 0.5), and the presence in the synthesis of an excess of aluminium is critical for its formation. Thus, the aluminium is not only a catalyst support in the form of ~-A1203, but it is also an element in the active phase. The Sb:V ratio in the trirutile usually falls in the range 2-5 (cf. Fig. 6). In the A1-Sb-V-W-O system a ruffle-type phase with the approximate composition A10.1Sb0.8V0.7W0.404 is active for propane ammoxidation. It is predominantly formed in situ during the ammoxidation. Compared with the rutiles in the other two systems, the A1-Sb-V-Woxide has a lower Sb:V ratio. Under the reaction conditions that were used in the present investigation the highest yields to acrylonitrile obtained were 11% for the Sb-V-oxide, 20 % for the A1-Sb-V-oxide and 37 % for the A1-Sb-V-W-oxide. Thus, considerable enhancement of the yield is achieved by adding aluminium and tungsten to the catalyst. The results can be rationalised in tenns of site isolation [4,16]. On active Sb-V-oxide the excess antimony at the surface creates appropriate isolation of the active vanadium centres. In the A1-Sb-V-oxide the aluminium effects isolation of vanadium not only at the surface, but also in the bulk of the active phase. Compared with these two systems, the introduction of tungsten in the rutile structure further tunes the active ensemble to a configuration which is more efficient for propane ammoxidation. Surrounding the vanadium sites with aluminium and tungsten produces a catalyst which not only gives less amount of C 1-2 products, but also less degradation of ammonia (see Fig. 4). REFERENCES

1. 2. 3.

4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18.

A.T. Guttmann, R.K. Grasselli and J.F. Brazdil, US Patents 4 746 641 (1988) and 4 788 317 (1988); assigned to The Standard Oil Company (Ohio). G. Centi, R.K. Grasselli, E. Patane and F. Trifir6, in G. Centi and F. Trifir6 (eds.), New Developments in Selective Oxidation, Studies in Surface Science and Catalysis, Vol. 55, Elsevier, Amsterdam, 1990, pp. 515-526. R. Nilsson, T. Lindblad, A. Andersson, C. Song and S. Hansen, in V. Cort6s Corber~in and S. Vic Bell6n (eds.), New Developments in Selective Oxidation II, Studies in Surface Science and Catalysis, Vol. 82, Elsevier, Amsterdam, 1994, pp. 293303. R.K. Grasselli and J.D. Burrington, in D.D. Eley, H. Pines and P.B. Weisz (eds.), Advances in Catalysis, Vol. 30, Academic Press, New York, 1981, pp. 133-163. J. Nilsson, A.R. Landa-C~inovas, S. Hansen and A. Andersson, J. Catal., 160 (1996) 244. R. Catani, G. Centi, F. Trifir6 and R.K. Grasselli, Ind. Eng. Chem. Res., 31 (1992) 107. T. Birchall and A.W. Sleight, Inorg. Chem., 15 (1976) 868. F.J. Berry, M.E. Brett and W.R. Patterson, J. Chem. Soc. Dalton Trans., (1983) 9. A. Landa-C~inovas, J. Nilsson, S. Hansen, K. StS.hl and A. Andersson, J. Solid State Chem., 116 (1995) 369. A.R. Landa-C~inovas, S. Hansen and K. Stgthl, Acta Crystallogr. B, in press. J. Nilsson, A. Landa-C~inovas, S. Hansen and A. Andersson, Catal. Today, 33 (1997) 97. G. Centi and P. Mazzoli, Catal. Today, 28 (1996) 351. S. Hansen, K. St~hl, R. Nilsson and A. Andersson, J. Solid State Chem., 102 (1993) 340. R.G. Teller, M.R. Antonio, J.F. Brazdil and R.K. Grasselli, J. Solid State Chem., 64 (1986) 249. R. Nilsson, T. Lindblad and A. Andersson, J. Catal., 148 (1994) 501. J.L. Callahan and R.K. Grasselli, AIChE J., 9 (1963) 755. K. Brandt, Ark. Kemi, 17 (1943) 1. S. Hansen, A. Landa-C~inovas, K. St~hl and J. Nilsson, Acta Crystallogr., A51 (1995) 514.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

423

hlfluence of A n t i m o n y Content in the Iron A n t i m o n y Oxide Catalyst and Reaction Conditions on the ( A m m ) O x i d a t i o n of Propene and P r o p a n e Eric van Steen, Gunther Kuwert, Alvin Naidoo and Marco Williams Catalysis Research Unit, Department of Chemical Engineering, University of Cape Town, Private Bag, Rondebosch 7700, South Africa The influence of antimony content and of ammonia partial pressure on the selectivity of the (amm)oxidation of propene and propane with FeSbO4 can be understood in terms of the degree of reduction of the catalyst surface at steady state conditions. The higher the degree of reduction which can be caused either by a low antimony content or by a high ammonia partial pressure the higher the selectivity for the combustion/degradation products. 1. INTRODUCTION Valuable chemicals can principally be produced from paraffins if these unreactive compounds can be functionalized. This primarily requires the activation of the rather unreactive paraffinic C-H bond or C-C bond. Therefore, high temperatures are usually applied in the functionalization of paraffins. An interesting route is the partial oxidation of these compounds. This may yield olefins by oxidative dehydrogenation, oxygenates by oxygen insertion, or even nitrile compounds if a reactive nitrogen compound is added to the feed. The partial oxidation of olefins over a variety of mixed oxide catalysts is well known [ 1] and has been studied in great detail. For example, iron antimony oxide is known to be a selective catalyst for the partial oxidation of propene yielding acrolein [2-4] and of 1-butene yielding 1,3 butadiene and 2-butenal [5,6]. The oxidation of paraffins, like propane, with this type of catalysts yields only the combustion products CO and CO2 [7]. Propane can however be selectively oxidised over this type of catalysts in the presence of ammonia to yield acrylonitrile [8]. The active and selective phase in the iron antimony oxide catalyst for the selective partial oxidation of propene is FeSbO4 [3,8]. The selectivity to acrolein, however, improves dramatically if the antimony to iron ratio in this catalyst exceeds one [3 ]. It has been observed, that the surface of iron antimony oxide with Sb/Fe = 1 is enriched with antimony [8,10]. If the Sb/Fe ratio in the catalyst formulation is increased, then a higher enrichment of the surface with more antimony was observed [8]. It has been indicated for the analogue U-Sb system [9], that the surface antimony inhibits the reduction of the catalyst during the reaction. Different oxygen species are involved in the partial (amm)oxidation reaction [1]. The presence of e.g. 0 2. will lead to a nucleophilic addition to the adsorbed hydrocarbon species resulting in the formation of the selective partial (amm)oxidation products. The presence of electrophilic oxygen species like O2-and O lead to the formation of degradation and combustion products. All these oxygen species can be formed from gas phase oxygen. 0 2. can also act as lattice oxygen. Although nucleophilic oxygen species are formed in the re-oxidation process of the catalyst, the relative concentration of these species will be less on surface with a

424 higher degree of reduction. The degree of reduction of the catalyst surface during the reaction will therefore influence the relative composition of the pool of oxygen species at the surface. The presence of ammonia in the ammoxidation is expected to change the degree of reduction of the catalyst at steady state conditions [9], and hence a change in selectivity should be observed. In this paper we want to compare the influence of antimony and ammonia on the (amm)oxidation of propene and propane, which are expected to have opposite effects. We will explain the observed changes in terms of a change in the degree of reduction of the surface, and hence in terms of changed composition of the pool of oxygen species at the surface.

2. EXPERIMENTAL The iron antimonate catalysts with an antimony to iron ratio of 1:1 and 2:1 was prepared using the method described by Allen et al. [4]. Briefly, Sb203 was added to a solution of iron nitrate in its own crystal water at 60 ~ The temperature was subsequently raised to 80 ~ and an aqueous 25 % NH3-solution was added to raise the pH to 3. The precursor was then dried for 16 hours at 125 ~ in an oven. Finally, the catalyst precursor was calcined at 800 ~ for 8 hours. Two different methods were attempted to enrich the iron antimony catalyst with Sb/Fe = 1 with antimony. With both methods the catalyst was contacted with a Sb-containing solution (2 ml/gcatalyst), which was prepared by dissolving Sb203 in HC1. Method A was the impregnation of iron antimonate with antimony. For method B the antimony was precipitated by adjusting the pH to 8 by adding an aqueous 25 % NH3 solution. Subsequently, the catalyst was dried at 120 ~ for 4 hours and calcined at 800 ~ for 8 hours. For method A the impregnation and calcination steps were repeated once. The phases present in the catalysts were characterised using XRD. The bulk composition of the catalysts were determined using ICP and XRF. The ammoxidation was carried out in a fixed bed glass reactor [6]. The catalyst (0.5 g; dp < 0.1 mm) was diluted with sand (4 g; dp = 0.2 - 0.3 mm). In all experiments the inlet partial pressures of propene/propane and of oxygen were kept constant at respectively 15 kPa and 30 kPa. The inlet partial pressures of ammonia and nitrogen were adjusted to give a total of 90 kPa. The space velocity for the propene and propane (amm)oxidation was kept constant at 1.2 mmol hydrocarbon/(gcat min). Methane was added to the effluent as internal standard. The organic products in the effluent were analysed with a GC equipped with an FID using the offline ampoule sampling technique [ 11 ]. The combustion products CO and CO2 were analysed using an on-line IR-analyzer. HCN was trapped in a silver nitrate solution. Unconverted NH3 was trapped in a 0.1 M HCI solution. The amounts of HCN and unconverted NH3 were then determined by titration.

3. RESULTS AND DISCUSSION

3.1 Catalyst preparation and characterisation The XRD of the base iron antimony oxide catalyst with Sb/Fe = 1 showed only the presence of the FeSbOn-phase. The catalyst with Sb/Fe = 2 showed also the presence of c~-

425 Sb204. The two phases in this catalyst could not be visually identified and therefore not be separated easily. The calcined iron antimony oxide with Sb/Fe = 1 had a BET-surface area of 28.9 m2/g. Based on the surface area and an assumed diameter of Sb203 of 5.5 9 10a~ m the amount of antimony added to the catalyst either with method A or method B was calculated to be between 0.2 and 10 times a monolayer coverage. The final amount of antimony on the surface with method A will have been lower than intended, since Sb in a HCl-solution partly forms SbC13, which evaporates at 283 ~ The precipitation of the antimony from a HCl-solution (method B) led to the formation of the FeSbO4 and ot-Sb204 phase, which could be easily identified and separated. The phases were separated. The phase, which showed by XRD only the existence FeSbO4, was used in the catalytic studies. The preparation of the catalysts according to method B was more reproducible than the preparation according to method A. The catalysts used in the catalytic studies, which were prepared according to either method A or method B, showed by XRD solely the existence of FeSbO4. The antimony was either finely dispersed on the surface or very small Sb204 crystals were formed, which could not be detected by XRD. An attempt was made to measure the change in the bulk Sb/Fe ratios in these catalysts, but the observed changes were within the accuracy of the analyses. The results are therefore given based on the initial amount of Sb added to the base iron antimony oxide catalyst (Sb/Fe = 1).

3.2 Initial activity/selectivity in the ammoxidation of propene

Figure 1 shows the conversion of propene in the ammoxidation and the selectivity to the partial (amm)oxidation products, acrylonitrile plus acrolein, as a function of time on stream for the base iron antimony oxide catalyst (Sb/Fe =1) and for the impregnated catalyst with 0.27 g Sb added per g FeSbO4 (method A), which corresponds to a coverage with antimony of 5 100

30 Base CatTlyst

O

E 25

Impregnated catalys!

8O

.ooel 20

~60

9

tD

> 15

~ 10 o~

ml.m u

9

9

40

0

• m

5 !mpregnated ~

0 -1

0 1 log (Time on Stream/min)

Base Catalyst (Sb/Fe = 1)

~
9

0 2

-1

0 1 log (Time on Stream/min)

2

Figure 1 Time on stream behaviour in the ammoxidation of propene (pNH3,inlet - - 10 kPa; T = 380 ~ over the base iron antimony oxide catalyst (Sb/Fe = 1) and the impregnated iron antimony oxide catalyst (method A; 0.27 g Sb/g FeSbO4) left: conversion ofpropene right: selectivity of the partial (amm)oxidation products acrolein plus acrylonitrile

426 monolayers of SbzO4. The conversion decreased with time on stream as was observed in the partial oxidation of propene [7]. This might be ascribed to an increasing degree of reduction with time on stream. The formation of all products involves surface oxygen which was depleted, since the catalyst became partially reduced during the reaction. Impregnating the base catalyst with antimony reduced the activity. This might be attributed to a lower reducibility of the Sb-covered surface, which would result in a lower activity for oxygen release. The selectivity of the partial oxidation products, acrolein plus acrylonitrile, remained essentially constant with a small decrease alter prolonged times on stream (ca. 45 C-% for the base catalyst and 80 C-% for the impregnated catalyst). This behaviour is contradictory to that observed in the partial oxidation of propene where in the first minutes of the reaction a strong decrease in the acrolein selectivity accompanied by a strong increase in the selectivity of the combustion products CO and CO2 was observed [7]. This was explained in terms of a change in the composition of the pool of oxygen species, i.e. an enhancement of the surface concentration of electrophilic oxygen species relative to the concentration of nucleophilic oxygen species. The content of acrylonitrile in the Impregnated catalyst (method A) fraction of acrylonitrile plus acrolein 100 increased with time on stream (see Figure 2). It can be assumed that acrolein and 80 o acrylonitrile are both formed via the common ~ 60 rt-allyl intermediate, which can undergo either oxygen or NH-insertion [9,12,13 ] ~9 40 o ~" 20 b/Fe = 1) _ _ _+LOc _ CH2_-CH_.CH2_~, < 0 .,44 o... CH ,~H2 + ~ ~ ~ ~ !

0

i

-1

i

I

I

I

;

I

I

I

I

I

I

I

0 1 log (Time on Stream/min)

I

2

Figure 2 Influence of reaction time on the acrylonitrile content in the fraction of acrylonitrile plus acrolein (pNH3, inlet -- 10 k P a ; T = 380 ~ over the base iron antimony oxide catalyst (Sb/Fe = 1) and the impregnated iron antimony oxide catalyst (method A; 0.27 g Sb/g FeSbO4)

CH2 ::

+

CH2=CH--CH~.NH

The acrylonitrile content in the fraction of acrylonitrile plus acrolein can therefore be interpreted as the probability of a NHinsertion instead of an oxygen insertion in the x-allyl complex. The increase with time on stream in the acrylonitrile content then means an increase in the probability of insertion of

the NH-species or a decrease in the probability of oxygen insertion in the rt-allyl complex with time on stream. This can be understood by realising, that the degree of reduction of the oxidation catalyst increases with time on stream This would lead to a decrease in the concentration of surface oxygen species. Since the probability of oxygen insertion depends amongst others on the concentration of surface oxygen species relative to the concentration of NH-species, the probability of oxygen insertion is expected to decrease with increasing degree of reduction.

427 The acrylonitrile content was initially much higher for the impregnated catalyst. This indicates, that oxygen release by the catalyst and hence oxygen insertion is much more difficult if the surface is enriched with antimony. The catalyst impregnated with antimony showed an increase in the acrylonitrile content in the fraction of acrylonitrile plus acrolein from 90 to 95 C-% with time on stream, whereas the content of acrylonitrile with the base catalyst increased from 6 to 95 C-%. The relatively small increase with the impregnated catalyst, in comparison to that with the base catalyst, indicates that the antimony covered surface of the impregnated catalyst inhibits the reduction of FeSbO4. At steady-state reaction conditions the probability of oxygen insertion in the propene ammoxidation over the iron antimony oxide base catalyst (Fe/Sb = 1) is the same as for the with antimony impregnated catalyst. This can only be the case if the ratio of the rate constant for oxygen insertion times the concentration of the surface oxygen species relative to the rate constant for the NH-insertion time the surface concentration of NH-species are approximately equal. This can only be realised if the ratio of the rate constant of oxygen insertion times the surface oxygen concentration and the rate constant for NH-insertion times the concentration of these NH-species is approximately the same for both catalysts. The surface of the base catalyst is more reduced and therefore the concentration of the surface oxygen species is lower than with the impregnated catalyst. If the rate constant for NH-insertion is assumed to be independent of the surface coverage with antimony, then increasing the surface coverage with antimony must lead to either a higher concentration of surface NH-species higher or a lower rate constant for O-insertion. 3.3 Influence of antimony precipitated on FeSbO4 surface The effect of antimony precipitation on FeSbO4 (method B) was evaluated using the ammoxidation of propene and propane. The exact amount of antimony which was precipitated on the surface could not be determined. The catalyst used in the catalytic study showed by XP~ only the presence of FeSbO4 not of ~-Sb204. The conversion of propene and propane both decreased with increasing the amount of antimony added to the base catalyst, FeSbO4 (see Figure 3). Also, iron antimony oxide catalysts prepared according to the method described by Allen et al. [4] showed a decrease in the rate at which hydrocarbons are converted with increasing antimony content [3,7]. If the surface of FeSbO4 is covered with antimony, then the ability to release oxygen becomes less [9]. The conversion of propane at a temperature of 60 ~ higher is much lower than the conversion of propene. This illustrates the difficulty of the activation of a paraffin in comparison to the activation of an olefin. The selectivity for the partial (atom)oxidation products acrolein and acrylonitrile increased with increasing amount of antimony on the iron antimony oxide surface. This can be ascribed to a change in the composition of the pool of oxygen species when changing the surface antimony content. Since antimony inhibits the surface reduction, a higher concentration of nucleophilic oxygen will be present resulting in a higher selectivity for acrylonitrile/acrolein. In the propane ammoxidation a lower selectivity for acrolein plus acrylonitrile is observed. The formation of partial (amm)oxidation products from propane requires more elemental steps than their formation from propene. All these intermediates can undergo a side reaction with electrophilic oxygen species yielding degradation products.

428

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0.1

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Antimony added, g Sb/g cat

Figure 3: Influence of amount of antimony precipitated (method B) on the iron antimony oxide catalyst (Sb/Fe = 1) on the ammoxidation of propene (PNm = 10 kPa; T - 370 ~ and of propane (pNm = 10 kPa; T = 430 ~ left: conversion of propene and propane right: selectivity of the partial (amm)oxidation products acrolein plus acrylonitrile

100 !

~

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,,

9 Propene

so

r,.) .-~

70

~

60

-

Propane

o

<

50 0.05 0.1 0.15 0.2 Antimony added, g Sb/g cat

Figure 4: Influence the amount of antimony added to FeSbO4 (method B) on the acrylonitrile content in the fraction of acrylonitrile plus acrolein in the ammoxidation of propene (pyre, inlet = 10 kPa; Y = 370 ~ and of propane (p~3, i.lot = 10 kPa; T = 430 ~

The acrylonitrile content in the fraction of acrylonitrile plus acrolein in the partial(amm)oxidation of propene (see Figure 4) was almost independent of the amount of antimony added to FeSbO4, whereas in the propane (atom)oxidation a minimum was observed. With increasing antimony content the degree of reduction of the surface decreases and hence the concentration of surface oxygen species increases. This would lead to an enhanced probability for oxygen insertion and thus to a decreased content of acrylonitrile in the fraction of acrylonitrile plus acrolein. There is apparently another factor countering the effect of increasing surface oxygen concentration. This must be then either the rate constants for oxygen or NH-insertion or the concentration of surface NHspecies.

429 3.4 Influence of ammonia partial pressure on the (amm)oxidation of propene

The influence of ammonia on the partial (amm)oxidation of propene was studied over the iron antimony oxide catalyst (Sb/Fe = 2) at 375 ~ (see Figure 5). The yield of the partial (amm)oxidation products acrylonitrile plus acrolein decreased with increasing ammonia partial pressure. The yield of the combustion products CO and CO2 first decreased and then increased with increasing ammonia partial pressure. The opposing trends for the yield of both product groups resulted in a complex behaviour of the conversion of propene as a function of the partial pressure of ammonia. The rate of formation of the partial (amm)oxidation products can be easily modelled as a surface reaction ocupying one or two active sites, and ammonia occupying one of the sites. racrylonitrile + acrolein -A B + pyH3 The rate of formation of acrolein could be modelled assuming that the formation of the rtallylic complex, i.e. the first hydrogen abstraction, is the rate determining step [ 13,14]. Oxygen occupies a different site and its coverage is negligible. If the rate determining step remains unchanged in the presence of ammonia, then the observed dependency indicates, that propene and ammonia are competing for similar sites. The complex dependency of the yield of combustion products on the ammonia partial pressure indicates that various factors are influencing the rate of formation of this product class. If ammonia is only inhibiting the adsorption of propene, then the yield of combustion products is expected to follow the same trend as the yield of the partial (amm)oxidation products. Ammonia is not only consumed for the formation of acrylonitrile, but can also reduce the surface under the formation of e.g. N2. This will change the degree of reduction of the surface, and hence the composition of the pool of oxygen species. If ammonia cannot

_ A

100 (amm)oxidation products

~ 8o -pP~a~'~~amm)oxidafion products r..)

4

~ 6o ]

.,..~ >" 2

9

roduc

~ 40 ~ 20

degradation p~ucts

" combustion products 0

10 20 30 40 Ammonia partial pressure, kPa

0

10 20 30 40 Ammonia partial pressure, kPa

Figure 5: Influence of ammonia partial pressure on the (amm)oxidation of propene over an iron antimony oxide catalyst (Sb/Fe = 2) at 375 ~ left: yield of the (amm)oxidation products acrylonitrile and acrolein and of combustion products CO and CO2 right: selectivity of the partial (amm)oxidation products for the partial (atom) oxidation products acrolein plus acrylonitrile, the combustion products CO and CO2, and the degradation products

430 occupy sites which can accomodate the electrophilic oxygen species, then the rate might increase. The selectivity for the partial (amm)oxidation products first increased and then decreased with increasing ammonia partial pressure. A very high selectivity (close to 100 C-%) was obtained at a molar ammonia to propene ratio of less than one. The selectivity for the combustion products showed the opposite trend. At high ammonia partial pressures a slight increase in the selectivity for the degradation products was observed, which are thought to be formed similarly to the combustion products. The acrylonitrile content in the fraction 100 of acrylonitrile plus acrolein as a function o 80 of the ammonia partial pressure (see Figure 6) increased strongly at low ammonia = O 60 partial pressures. This can be attributed to the increase in the concentration of NHN 40 surface species. Ammonia in the feed will N 20 lead to a higher degree of reduction of the oxidation catalyst. This will lead to a lower I I I < 0.. concentration of nucleophilic oxygen 10 20 30 40 species on the surface. Thus, the increase Ammonia partial pressure, kPa in the acrylonitrile content will be enhanced. At higher ammonia partial Figure 6 Influence the ammonia partial pressures the acrylonitrile content pressure on the acrylonitrile decreased. The occurance of this maximum content in the fraction of at a relatively low ammonia partial pressure acrylonitrile plus acrolein in the was unexpected. This means that the ammoxidation of propene over an oxygen insertion becomes more likely than iron antimony oxide catalyst the NH-insertion at higher ammonia partial (Sb/Fe = 2) (pNH3, inlet -- 10 kPa; T pressures. NH-species can be formed by = 370 ~ successive hydrogen abstraction from an adsorbed NH3-species. The hydrogen will be transformed to an oxygen species yielding OHgroups and water. If the catalyst is getting more reduced than the adsorbed propene and the adsorbed NH• (x = 2-3) are competing for the same oxygen surface species on a partially reduced surface. Due to the lower availability of this surface oxygen species, the surface concentration of NH-species decreases and the acrylonitrile content will decrease. This can be easily illustrated using following, highly simplified, reaction steps assuming that the oxidative dehydrogenation of ammonia is reversible, the existence of one single type of oxygen species, and the rate determining step in both the formation of acylonitrile and acrolein is the addition of a NH-species or an oxygen species to the allylic complex (C3Hs,ads) [9,12,13]" NH3, ads + Oad~~

NH2, ads + OHad~

NH2, ads + Oads ~

NI~d s

+

OHads

C3 H5, ads d- NI-Iad s

>

......

> acrylonitrile

Ca Hs, ads +

>

......

~ acrolein

Oads

It can then be easily derived that the ratio of the rate of formation of acrylonitrile relative to the rate of formation of acrolein will be:

431 racr,/ionitrilc ~ [Oads] [NH3,~ds] racrolein [OHads] 2 If the concentration of oxygen species decreases due to the increase in the ammonia partial pressure, then a maximum is expected for acrylonitrile content, i.e. for the probability of NHinsertion instead of oxygen insertion in the ~-allylic complex. 3.5 Influence of t e m p e r a r e on the a m m o x i d a t i o n of p r o p e n e / p r o p a n e

The temperature influence on the ammoxidation of propene and propane was studied over FeSbO4 at a fixed ammonia partial pressure of 10 kPa (see Figure 7). The apparent activation energy for the formation of acrolein and acrylontrile from propane was almost twice the apparant activation energy for the formation of these compounds in the propene conversion (90 vs. 55 kJ/mol). This might reflect a different rate determining step in the propane conversion compared to the one in the propene conversion. In the latter case the formation of the allylic complex is most likely to be the rate determining step [ 13,14]. For the activation of propane a surface propyl group is a very likely initial intermediate [8], whose formation might be energetically less favoured than the second hydrogen abstraction and the third hydrogen abstraction yielding the ~-allylic adsorbed species. -2

Ea (kJ/mol) = -3 - 55 ~

~s

-4

E -5

-

100 Propene

~

9o

~60

-

e

+

' 1.3

1.4

n nnn~nn \9

40

"~ "~ "-6 ~ ~ 20 < 0

"~ -6 = -7

opane

80

1.5 1.6 1000/T, 1/K

1.7

Prop ' 300

ene~ ' ~ '

'

'

'

400 Temperature, ~

'

' 500

Figure 7 Influence of temperature on the formation of acrylonitrile plus acrolein in the ammoxidation of propene and propane (pNm, inlot = 10 kPa) over an iron antimony oxide catalyst (Sb/Fe = 1) left: Arrhenius' plot right: Selectivity of acrylonitrile plus acrolein The selectivity for the partial (amm)oxidation products decreased with increasing temperature. This indicates a higher activation energy for the formation of combustion/ degradation products than for the selective formation of the partial (amm)oxidation products. With propane a higher selectivity for acrylonitrile and acrolein was observed. In the propane (amm)oxidation the content of acrylonitrile increased with increasing temperature, which indicates that the insertion with oxygen becomes less probable with increasing temperature. This might be attributed to either the difference in the activation energy for the NH- and Oinsertion and/or a higher degree of reduction of the partial oxidation catalyst. In the propene

432 ammoxidation the acrylonitrile content showed a maximum, which can also be explained by an increase in the degree of reduction of the catalyst surface (vide supra). The decrease in the selectivity of the partial (amm)oxidation products acrolein plus acrylonitrile might therefore not only be attributed to the differences in activation energy, but can also be caused by the change in the degree of reduction. 4. CONCLUSIONS Iron antimony oxide catalysts enriched with antimony have been prepared. The impregnation of FeSbO4 is not very reproducible. The catalyst prepared according to method A showed in the propene partial oxidation results similar to iron antimony oxide catalysts with Sb/Fe = 2. In the preparation of the catalysts according to method B most of the antimony precipitated as a separate phase. The selectivity changes in the (amm)oxidation of propene and propane can be (partly) understood in terms of a change of the degree of reduction of the surface, which would cause a change in the distribution of nucleophilic and electrophilic oxygen species in the pool of surface oxygen species. Antimony inhibits the surface reduction and the catalysts with an antimony enriched surface therefore exhibit a higher selectivity with a lower activity. Ammonia inhibits the formation of partial (amm)oxidation products by competitve adsorption. It can also modify the degree of reduction of the surface and thereby shitting the selectivity of the reaction. To obtain a maximal selectivity to the partial (amm)oxidation products a molar ammonia/propene ratio of less than one should be applied. The ammoxidation of propene has a higher apparent activation energy than the ammoxidation of propene, but shows a higher selectivity. REFERENCES A. Bielanski and J. Haber, Oxygen in Catalysis, Marcel Dekker Inc., New York, 1991 [1] G.K. Boreskov, S.A. Venyaminov, V.A. Dzisko, D.V. Tarasova, V.M. Dindoin, N.N. [21 Sanobova, I.P. Olenkova, and L.M. Keteil, Kinet. Katal. 10 (1969) 1530 I. Aso, S. Furukawa, N. Yamazoe, and T. Seiyama, J. Catal. 64 (1980), 29 [31 M. Allen, R. Bettely, M. Bowker, and G. Hutchings, Catal. Today 9 (1991) 97 [4] G.I. Straguzzi, K.B. Bischoff, T.A. Koch, and G.C.A. Schuit, J. Catal. 104 (1987) 47 [5] E. van Steen, M. Schnobel, and C.T. O'Connor, in Heterogeneous Hydrocarbon [6] Oxidation (B.K. Warren, S.T. Oyama, eds.), ACS Symposium Series 638, ACS Washington DC, 1996, p. 276 M. Schnobel, PhD Thesis, University of Cape Town, 1997 [7] M. Bowker, C.R. Bicknell, and P. Kerwin, Appl. Catal. A: General 136 (1996) 205 [8] R.K. Grasselli and D.D. Suresh, J. Catal. 25 (1972), 273 [9] [10] S.R.G. Carraz/m, L. Cadus, Ph. Dieu, P. Ruiz, and B. Delmon, Catal. Today. 32 (1996), 311 [11] H. Schulz, W. Bohringer, C. Kohl, N. Rahman, and A. Well, DGMKForschungsbericht 320, DGMK Hamburg, 1984 [12] J.D. Burrington, C.T. Kartisek, and R.K. Grasselli, J. Catal. 81 (1983), 489 [13] J.D. Burrington, C.T. Kartisek, and R.K. Grasselli, J. Catal. 87 (1984), 363 [14] G.W. Keulks, and M.-Y. Lo, J. Phys. Chem. 90 (1986), 4768

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

433

Catalytic selective o x i d a t i o n of C2-C4 alkanes over reduced heteropolymolybdates Wen Li and Wataru Ueda Department of Environmental Chemistry and Engineering, Tokyo Institute of Technology, Nagatsuta-cho 4259, Midori-ku, Yokohama 226 JAPAN Catalytic oxidation of ethane, propane, and isobutane to corresponding unsaturated acids and acetic acid with molecular oxygen were carried out over reduced H3PMo12040 catalysts above 300~ A highly reduced state of H3PMo12040 was formed by the heat-treatment of the pyridinium salt under N2 flow up to 420~ This catalyst, denoted by H3PMo12040(Py), gave the highest conversion of alkanes and selectivities to the objective products and its reduced state was highly stable under the conditions of catalytic oxidation. The highly stable reduced state was composed of defect Keggin unit with oxygen vacancies and a secondary structure formed by the defect Keggin unit and remaining pyridine in the structure. Another reduced catalyst generated by the heat-treatment of (NH4)3PMo12040 was active for the propane and isobutane conversion but poorly selective to acrylic acid and methacrylic acid, because this catalyst was readily oxidized during the oxidation. Non-reduced H3PMo12040 catalysts showed poor activity for all of the reactions. A reaction mechanism is proposed over the reduced H3PMol2040(Py) catalyst. 1. INTRODUCTION Because of the global abundance of liquefied petroleum gas (LPG), interest in the potential use of ethane, propane, and butanes as sources of the corresponding alkenes or their derivatives is increasing [1]. In the last decade much progress has been made, particularly in the selective partial oxidation of light alkanes with molecular oxygen in gas phase [ 1, 2]. For economic reasons, molecular oxygen is usually used as the primary oxidant[3]. To promote both the conversion of reactants and the selectivity to partial oxidation products, many kinds of metal compounds are used to create catalytically active sites in different oxidation reaction processes [4]. The most well-known oxidation of lower alkanes is the selective oxidation of n-butane to maleic anhydride, which has been successfully demonstrated using crystalline V-P-O complex oxide catalysts [5] and the process has been commercialized. The selective conversions of methane to methanol, formaldehyde, and higher hydrocarbons (by oxidative coupling of methane [OCM]) are also widely investigated [6-8]. The oxidative dehydrogenation of ethane has also received attention [9,10]. Heteropolycompounds, having unique structures and the resulting acidic and redox properties, exhibit high oxidation abilities in the selective oxidations of alkanes [ 1, 2, 11, 12]. Especially, it should be noted an interesting fact that the heteropolycompounds containing molybdenum showed high catalytic activities in the partial oxidation of hydrocarbons to corresponding acids [10, 13-23]. The direct oxidation of light alkanes to corresponding acids has tremendous economic value. Although it seems very difficult, some related research results have been reported with the catalysts of heteropolycompounds containing molybdenum for the reactions of isobutane to methacrylic acid [3, 13, 19, 22, 23] and propane to acrylic acid [13, 17, 21]. There appears to be no published literature pertaining to the production of

434 acetic acid from ethane oxidation directly over heteropolycompound catalysts. Recently we found that highly reduced H3PMo12040 which was formed by the heattreatment of pyridinium salt can catalyze the propane oxidation to acrylic acid and acetic acid selectively [24, 25]. After activation in N2 flow at 420~ for 2hr, the catalyst of H3PMo 12040 (Py) shows reduced state of molybdenum and a new stable structure in which pyridine remains as the linkage of the secondary structure. The activated H3PMo12040 (Py) also gives catalytic activities in the partial oxidations of ethane and isobutane to acetic acid and methacrylic acid respectively. In this paper, we will report the oxidation results of C2-C4 alkanes and discuss the roles of reduced state and activation of molecular oxygen over this catalyst.

2. EXPERIMENTAL 2.1 Catalyst preparation 12-molybdophosphoric acid (I--I3PMol2040.xH20) was obtained from Nippon Inorganic Color & Chemical CO., LTD.. H3PMol2040.xH20 was dissolved in distilled water, then the solution was filtered, evaporated at 40-50~ with stirring until there was little crystals on the surface of the solution. The mixture was kept at 5~ for overnight to recrystallize, and the crystal was filtered and dried at 40~ for 8hr. By this way, H3PMo 120,0.10H20 was obtained. Pyridine-treated heteropolycompound, denoted by (Py), was prepared by precipitation method. Recrystallized H3PMo12040 was dissolved in distilled water with stirring at 40-50~ and desired amounts of pyridine was added slowly, then the solution containing precipitates was evaporated to dryness at 40~ The obtained solid was further dried in N2 flow at 120~ for 8hr. For the synthesis of (NH4)3PMo12040, desired amounts of (NH4)3MoTO24 and H3PO4 were dissolved in distilled water at about 40-50~ then HNO3 was added slowly with stirring at 50~ The addition of HNO3 resulted in a yellow precipitate, which was filtered, washed with distilled water, and dried at 40~ overnight. The obtained solid contains the desired composition and shows cubic structure as reported in literature [20]. Before catalytic reaction, H3PMol2040(Py) and (NH4)3PMo12040 were activated in N2 flow for 2h at 420~ and 450~ respectively.

2.2 Catalytic testing Alkane oxidations were carried out at an atmospheric pressure in a conventional flow system with a fixed bed Pyrex tubular reactor( qb 12mm). Catalyst(3g) was diluted by 2g sands to prevent the catalyst from overheating during reaction. The catalysts were heat-treated in a nitrogen flow at a selected temperature, then set to desired reaction temperatures. The feed compositions were controlled with mass flow controller (KOFLOC 3510). The total flow rate was 30 ml.min l. The feedstock and products were analyzed by on-line gas chromatography operating with two sequential columns, molecular sieve 13X lm at room temperature for separation of 02, N2 and CO, and Porapak Q 4m at 60"C-140~ for hydrocarbons and CO2. There was a ice cooling trap to collect other products at the exit of the reactor, and the collected products were analyzed quantitatively by another gas chromatography (TC-WAX 60m capillary).

2.3 Characterization TPD The experiments were carried out using a standard apparatus. Helium was used as a carrier gas. The thermal conductivity detector was employed to detect the changes of desorption. A dry ice trap was used to remove the water in the carrier gas. The weight of sample was 10mg. The sample was pretreated in He flow at room temperature for lh before the experiment. The heating rate was 10~ l. XPS A Shimadzu ESCA-750 electron spectrometer with an aluminum anode (1486.6eV) was used to obtained X-ray photoelectron spectra. All binding energies were referenced to gold (Au 4f7/2 line; 83.8eV) which was deposited on samples in vacuum. The activated sample in the flow system was outgassed in the preparation chamber of the spectrometer

435

under 10~ Torr for 15 minutes. 3. R E S U L T S AND DISCUSSION 3.1 Ethane Oxidation We examined the ethane oxidation over the mentioned catalysts under different reaction conditions at 280~176 About 1% conversion of ethane and 10% selectivity to acetic acid were obtained over the activated H3PMol2040(Py) at 3400C. Ethane oxidation did not occur over non-reduced H3PMo12040 and activated (NH4)3PMo12040 catalysts. There seems to be no literature about the direct oxidation of ethane to acetic acid over heteropolycompounds catalysts. Nevertheless, there is a limited amount of literature[ 10, 26-28] about direct oxidation of ethane to acetic acid over oxide catalysts at low temperature (200350~ It seems that vanadium and molybdenum are necessary to those catalysts, and the addition of water is useful to increase the production of acetic acid. Roy et al. [ 10] has proved that vanadium and molybdenum phosphates supported on TiO2-anatase were effective in the direct oxidation of ethane to acetic acid. Considering previous research results, it is suggested that other promoters, such as transition-metal oxides, are necessary to enhance the catalytic activity of the activated H3PMol2040(Py) in the direct oxidation of ethane to acetic acid. 3.2 Propane Oxidation The catalytic activities of H3PMo 12040, (NH4)3PMo 12040 and I-I3PMo 12040(Py) in the propane oxidation are shown in Figure 1. The non-reduced, acidic form of molybdophosphate catalyst, H3PMo12040, revealed an activity and yielded propene mainly. The activated H3PMol2040(Py) catalyst showed distinctly enhanced activities for the propane oxidation, and the catalyst gave a significantly different product distribution, where acrylic acid and acetic acid were main organic oxygenated products. Such drastic chaiage of the product distribution was observed even at a low conversion of propane under a short contact time. Obviously, this results from an intrinsic change of catalyst properties by the treatment with pyridine. The product distribution in activated H3PMo12040(Py) catalyst is different from that in the reported metal oxides catalyst or heteropolycompounds catalysts promoted by metal atoms, such as Bi, Fe, Ni, or Cs substitution for proton and V s+ substitution for Mo 6. in

Catalyst (precursor) [S.A.: m2/g ]

0

Conversion of Propane (%) 2 4 6 8 |

lO [] Acetic acid 9 Acrylic acid [] Propene

w

(NH4)3PMo12040 [ 15.21 H3PMol2040(Py) [12.8] H3PMo12040(Py)

(0.Sg) H3PMo12040 [3.8]

73% NgJ/F. .

.

[

!

0

10

.

.

. i

~ !

,t~ "

20 30 40 Selectivity(%)

Figure 1. Catalytic performance for the propane oxidation at 340~ Catalyst weight: 3g; total flow rate: 50ml/min; reactants" C3H8:O2:H20:N2=20.4:9.9:19.9:49.8 (mol%).

436 the direct oxidation of propane to acrylic acid [13, 17, 21, 29]. The former shows relatively higher ability to form organic oxygenates than the latter. The mentioned metal promoters seemed to be favorable to form propene, and oxidize it to CO and CO2 directly. Figure 1 also shows the catalytic performance of the activated (NH4)3PMo12040 catalyst which has a surface area similar to the H3PMo12040(Py)catalyst. Non-activated (NH4)3PMo 12040was completely inactive but became active after the heat-treatment under the N2 flow at 450~ The catalyst, however, gave a broad product distribution. The special ability of the activated H3PMo 12040(Py) in the partial oxidation of propane allows us to investigate its catalytic performance in detail. The conversion of propane and selectivities to acrylic acid and acetic acid as a function of time on stream at 3400C are shown in Figure 2. The main products were acrylic acid and acetic acid, and the remainder is carbon oxides. At the beginning of the reaction (t<6h), the conversion and the selectivity to acrylic acid increased with the time on stream, then became to be stable. The selectivity to acetic acid was almost unchanged during the reaction. Since the H3PMo 12040(Py) catalyst was activated in N2 flow at 4200C for 2h and then was allowed to the reaction without being pre-oxidized, the changes of the activity are slightly different from the former case[25] in which the initial activity changes was not so long because the activated catalyst was pre-oxidized in 20%O2/N 2 at 380~ Although the activity of the catalyst without being pre-oxidized was unstable at the beginning of reaction, it became stable and gave the higher conversion and selectivity than that with being pre-oxidized. The result suggests that the catalyst can keep highly reduced state under the suitable reaction conditions to supply more active sites. 60

~_ 40

OI

9

C

/'~e---O.--. 0

0 _

0

0

--0-"

20

9

40 -

/

"3

;,- 20

_

~0

0 ~ ~ ~

/o.,,.o ~

r,..) 0

0

-.~.....

I

I

5 10 Time on stream (h)

I

15

Figure 2. Catalytic performance of reduced H3PMo12040(Py) in the propane oxidation at 340~ Catalyst weight: 3g; total flow rate: 30ml/min; reactants: C3H8:O2:H20: N2=3.3:3.3:33.3:60.1 (mol%); D: C3H8; O: acrylic acid; O: acetic acid.

0

0

I

I

I

5 10 15 20 Partial pressure of 0 2 (kPa)

Figure 3. Effects of 0 2 partial pressure on the propane oxidation over reduced H3PMol2040(Py) at 340~ Symbols and reaction conditions except oxygen partial pressure are the same as those in Figure 2.

The effects of oxygen partial pressure on the conversion of propane and the selectivities to acrylic acid and acetic acid under the typical reaction conditions are shown in Figure 3. With the increase of oxygen partial pressure, the conversion of propane slightly increased. On the other hand, the selectivity to acrylic acid increased greatly, and the selectivity to acetic acid decreased slightly. For the production of acetic acid, two pathways can be assumed: one is consecutive pathway from produced acrylic acid and the other is parallel production through a common intermediate for the acrylic acid formation. Since the result showed the constant formation of acetic acid irrespective of the oxygen pressure, it can be concluded that the latter

437 pathway is predominant. We tentatively consider that in the parallel production pathway a propene-like common intermediate readily changes into acetic acid by acid-promoted oxidative hydration process through acetone. For the production of acrylic acid, on the other hand, catalysts need to have an ability for allylic oxidation which can take place on coordinately unsaturated molybdenum site and with lattice oxygen. In this sense a reduced state is suitable for providing such sites, but if a highly reduced state can be kept by slow reoxidation with molecular oxygen the surface can not supply lattice oxygen sufficiently to achieve high selectivity for allylic oxidation products. Therefore, a higher oxygen pressure is necessary to have more lattice oxygen on surface. Since the overall reaction rate of propane was almost independent of the oxygen pressure, the selectivity to acrylic acid formation tends to increase with the increase of the oxygen pressure as observed in Figure 3. We observed that at the high oxygen partial pressure condition the activity began to decrease after 30h reaction accompanied by the formation of MoO3 phase in the catalyst. Nevertheless, it was found that the structure of activated H3PMol2040(Py) was unchanged under 20%O2/N2 flow at the temperature (T<380~ These results suggest that the reduced structure of H3PMol2040(Py) is stable and does not decompose by reoxidation with only molecular oxygen, but the structure of the catalyst is not stable under the catalytic reaction conditions specially with high oxygen partial pressure after a long reaction time. This phenomenon can be simply explained as follows: Once redox type reaction such as the allylic oxidation takes place over oxide surface, structural deformation is inevitably induced in oxide catalysts by the removal of lattice oxygen during the reaction. In the case of heteropolycompounds each discrete structure like a Keggin unit has to sustain the structural deformation, which is obviously hard. It might not be easy neither in the highly reduced H3PMol2040(Py) catalyst used in this work even though the reduced state of the catalyst is stable under oxygen, because its high resistance against reoxidation does not mean highly resistant against reduction-oxidation cycle of the catalyst. For overcoming this problem, addition 60 of certain metal element seems to be profitable but in many cases results in poor catalytic performance for alkane oxidation because a highly reduced state is often sacrificed. "-7.40Figure 4 shows the effects of reaction temperature on the conversion of propane and the selectivities to acrylic acid and acetic acid under the typical reaction conditions in the partial 20 oxidation of propane. It is clear that higher reaction temperature is preferable for the formation of 3 ~/ ~o_.o.acrylic acid in the propane oxidation over the reduced H3PMol2040(Py) catalyst. This occurs 0 I i i t also because higher reaction temperatures are 320 330 340 350 360 suitable for allylic oxidation. Under presently Temperature (~ optimized reaction conditions 50% selectivity to .~,.0~,--

acrylic acid and 24% to acetic acid were achieved at 12% conversion of propane. However, when the catalyst was used in the propane oxidation for a prolonged time more than 30 h at high temperature (T>350~ weak additional peaks identified to M o O 3 again appeared as in the reaction under the high oxygen pressure.

--

,-.,

Figure 4. Effects of reaction temperature on propane oxidation over reduced H3PMolZO40(Py). Reactantes: C3H8: O2:H20:N2=3.3:6.7:33.3:56.7; the other reaction conditions and symbols are the same as those in Figure 2.

3.3 Isobutane Oxidation

The h e t e r o p o l y m o l y b d a t e catalysts, H3PMo12040, (NH4)3PMo12040 and H3PMol2040(Py) were tested for the partial oxidation of isobutane to methacrylic acid. The

438 Conversion of isobutane (%) 10 20

Catalyst (precursor) [S.A.: m2/g ]

!

[71 Acetic acid m Methacrolein II Methacrylicacid

(NH4)3PMo12040 [ 15.2] H3PMol2040(Py) [12.8] H3PMo 12040(Py) (0.5g) H3PMo 12010 [3.8] 0

20 40 Selectivity(%)

60

Figure 5. Catalytic performance for the isobutane oxidation at 3000C. Catalyst weight: 3g; total flow rate: 30ml/min; reactants: i-C4HlO:O2:H20:N2=3.3:13.3:33.3:50.1 (mol%). activities of them at 300~ are shown in Figure 5. The main organic oxygenate products are methacrylic acid, methacrolein, acetic acid and a little amount of acrylic acid. The activity of the H3PMo12040 catalyst is quite similar to that reported in the literature[30]. The (NH4)3PMo 12040 catalyst showed a tendency to form methacrolein. The H3PMol2040(Py) catalyst gave the highest conversion (about 22%) and highest selectivity to methacrylic acid (about 52%). With time on stream, the conversion of isobutane increased slightly, and the selectivity to methacrylic acid was relatively stable. The effects of contact time on the isobutane oxidation over the H3PMol2040(Py) 60 30 catalyst are again shown in Figure 6. With the increase of the contact time, the conversion of isobutane increases greatly, and the selectivity of methacrolein decreases 34020 3 markedly to disappear. The change of the selectivity to methacrylic acid exhibits a mountain shape, and the selectivities to acetic acid and acrylic acid still increased slightly O O"~ 101~' after the top of the mountain. It is, therefore, clear that methacrolein is the intermediate to form methacrylic acid. Figure 7 shows the effects of oxygen 0 .. .. 0 0.000 0.025 0.050 0.075 0.100 partial pressure on the isobutane oxidation over the H3PMo12040(Py) catalyst. With the W/F (g.min.m1-1) increase of the oxygen partial pressure, the conversion of isobutane increased greatly, and Figure 6. Effects of contact time on isobutane the selectivities to methacrylic acid and oxidation over reduced H3PMo 12040(Py) at methacrolein decreased especially up to 300~ Catalyst weight: 3g; total flow rate: 14kPa. The selectivities to acrylic acid and 30ml/min; reactants: i-C4H 10:O2:H20:N2= acetic acid increased continuously with the 3.3:13.3:33.3:50.1 (mol%); D: i-C4H10; increase of the partial oxygen pressure. ~: acrylic acid; O: acetic acid; Obviously, these changes are different from 0: methacrylic acid; O: methacrolein. those observed in the partial oxidation of

"l

S

o

O

439 60

n/

'4o-

30

60

-2o

~40 (D

"" "~,

Nr 20"

),./o

.,_~ ca~ r

-10

raq 73 :2 20 9

L) (?

0

I ''~

,,

i

5 10 15 20 25 Partial pressure of 0 2 (kPa)

0

Figure 7. Effects of 0 2 partial pressure on isobutane oxidation over reduced H3PMolZO40(Py) at 3000C. Symbols and reaction conditions except oxygen partial pressure are the same as those in Figure 6.

280

300

320

340

Temperature (~ Figure 8. Effects of reaction temperature on isobutane oxidation over reduced H3PMo12040(Py). Reactants: i-C4H10: O2:H20: N2=3.3:6.7:33.3:56.7 (mol%); the other reaction conditions and symbols are the same as those in Figure 6.

propane (Figure 3). Two reasons for the difference can be considered. First the reactivity of isobutane is much higher than that of propane because of lower C-H bond strength of isobutane. Then it is quite reasonable that appreciable oxygen dependency on the isobutane oxidation is observed if the formation rate of surface active oxygen from mol~cular oxygen controls isobutane activation. Second, the selectivity changes are able to be ascribed to that the C-C bonds of isobutane and methacrylic acid are broken on acidic catalyst surface easier than those of in propane and acrylic acid. At a low ratio of isobutane to oxygen, the effects of reaction temperature on the catalytic oxidation over H3PMo12040(Py) are shown in Figure 8. The conversion of isobutane tremendously increased from 8% at 280~ to 45% at 340~ Accordingly, the selectivity to methacrylic acid decreased greatly from 58% to 19%, and the selectivities to acetic acid and acrylic acid increased greatly. The changes of the selectivity to methacrolein exhibited a mountain shape just like the result in the effect of contact time, so that this result also supports the consecutive formation of methacrylic acid through methacrolein. In summary of the isobutane oxidation, the H3PMol2040(Py) catalyst not only shows high activity of the direct oxidation of isobutane to methacrylic acid, but also keeps a highly reduced state even at low ratios of isobutane to oxygen, while in many other cases on direct oxidation of isobutane to methacrylic acid the high ratio of isobutane to oxygen was used probably in order to prevent catalysts from collapse, and products from over-oxidation[2].

3.4 Catalysts Properties and Reaction Mechanism In order to clarify why the H3PMol2040(Py) catalyst is so highly active for the formation of organic oxygenates in the partial oxidation of alkanes, we briefly investigated the reduced state of the catalysts. In many cases [31], it had been demonstrated that the heteropoly compound catalysts containing molybdenum exhibited reduced state which was formed due to the reaction between reactants or proton and lattice oxygen. In the study of silica supported 12-molybdophosphoric acid, Moffat[32] had proposed that the defect Keggin unit, formed by dehydration, was stabilized by SiO2, as the active sites in partial oxidation of methane. In addition, the desorption or reaction of the linkage molecules in the bulk of heteropoly compounds can also induce reduced states which result in a variety of catalytic and structural properties

440

C

b

I

100

I

I

200 300

I

I

I

400

500

600

Temperature (oc) Figure 9. TPD of Heteropolymolybdates. (a) H3PMo12040; (b) H3PMol2040(Py); (c) (NH4)3PMo12040.

225

230 235 240 Binding Energy (eV)

Figure 10. Mo3d XPS spectra of H3PMol2040(Py). (a) before activated; (b) activated.

such as the production of active sites and active phases. We, therefore, conducted TPD measurements of H3PMo 12040, (NH4)3PMo12040 and H3PMol2040(Py) samples to compare reduction step during heating under an inert gas. The results of TPD are shown in Figure 9. The assignments of peaks had been verified by GC-MASS-TPD technique. The prepared H3PMo12040.10H20 gives a similar result as that given by Moffat [32]. As shown in Figure 9-a, the water desorbing at about 125"C was assigned to those being hydrogen-bonded to the protons within the secondary structure, and the another wide peak at about 355"C was attributed to the dehydration from protons and lattice oxygen. At this stage the catalyst was reduced to some extent. However, the peaks at 442"C, attributing to the desorption of water also, appeared perhaps due to a collapse of Keggin-type structure [33], which means that the reduced state generated by the heat-treatment of acidic form of H3PMo 12040 compound is not stable. In fact, it was found in our work that the H3PMo12040 catalyst was always in a highly oxidized state after the alkanes oxidation irrespective of the reaction conditions. In the TPD of H3PMo12040(Py) (Figure 9-b), there were three peaks due to the desorption of pyridine at 240~ 400~ and 527~ Therefore, after the pretreatment at 4200C, one kind of pyridine which desorbs at about 5270C remained in the structure. Another peak at 467~ is attributed to the desorption of CO2 which is formed by the reaction of pyridinium ion in the lattice with the lattice oxygen. At the same time, a newly reduced structure was formed having a orthorhombic structure [25]. The reduced state generated in the H3PMo12040(Py) by this way was found to be highly stable even under the above mentioned oxidation conditions. XPS was used to ensure the reduction. Figure 10 shows the Mo3d spectra of H3PMol2040(Py). H3PMo12040(Py) (Figure 10-a) exhibits the Mo vI 3d at 234.2 and 231.1eV only. After the catalyst is treated in N2 flow at 420~ for 2hr, it reveals that the most part of molybdenum in surface Keggin units changes to the reduced state MoV(Figurel0-b). The binding energies of doublet of Mo v 3d are shown at 233.2eV and 230.3eV. Undoubtedly, the reduced molybdenum exists on the surface. Two main desorption peaks were observed at 265~ and 550~ in the TPD of the (NH4)3PMo 12040 sample as shown in Figure 9-c. Therefore during the pretreatment of the

441 (NH4)3PMo 12040 sample in N2 flow at 450~ for 2h, NH3 (including N 2) and water desorbed accompanied by reduction. The structure formed by this way was a cubic structure, but the cubic structure and the reduced state was unstable under the oxidation conditions, readily oxidized into normal H3PMo12040. Obviously, the pyridine treatment has an effect on the catalysts to be able to keep highly reduced states during the oxidation. It appears that pyridine induces the particular secondary structures for keeping the highly reduced state of Keggin unit against oxidation, since it was observed that adsorptive pyridine in the H3PMol2040(Py) catalyst was unchanged after the catalytic reaction (0.8 pyridine per Keggin unit by chemical analysis). The reduced states induced by the reaction between surface protons and lattice oxygen seem not stable against oxidation. We, therefor, think that a oxygen deficient structure of Keggin unit generated by the reaction with pyridine is completely different from that by the reaction between surface protons and lattice oxygen. As a consequence, the particular reduced structure of heteropolymolybdophosphoric acid is inevitably important for the formation of acrylic acid and methacrylic acid in the propane and isobutane oxidation with molecular oxygen. 4. CONCLUSIONS Our study suggests that both a protonic, acidic property and a reduced state of molybdenum in H3PMo 12040 play an important role in the selective oxidation of alkanes with molecular oxygen. A special, oxygen deficient structure which provides Lewis acid sites seems to be important, and the formation of active oxygen on the reduced surface and the activation of alkanes with the active oxygen thus formed are the key steps in the catalytic alkane oxidation with molecular oxygen. On the basis of these results and suggestions, we propose a possible reaction mechanism for alkane oxidation in Figure 11, where the cooperation of proton and electron in the reduced H3PMo12040 catalyst having acidity in the activation of molecular oxygen and the subsequent activation of propane are shown along with the consecutive allylic oxidation. Molecular oxygen first reacts with two protons and reduced Mo to form one molecule of water. The remaining oxygen atom is coordinated on Mo site and play the active oxygen species for propane activation. Then propane changes into adsorbed propene-like species, followed by the rapid consecutive allylic oxidation to acrylic acid. It should be noted ~o[-@

A n ndac,ivation

2H+ + 2e(Mo) weakacid reduced ~

(Oxygenactivation)

I|

Mo- OH

Acrylicacid

~

HN)

r ~176 ]

[.~-,

OH HO

,,~MjO _

Lhydration)

rio ]

Loxidation) HO - Mo= O

10Io" H20

HO -IVlo/

Figure 11. Proposed mechanism of propane oxidation.

442 that the highly reduced state is suitable for allylic oxidation because the reduced state provides not only coordinately unsaturated Mo for the adsorption (stabilization) of propene-like species but also basic lattice oxygen which can only be formed on the reduced Mo sites. A part of acrylic acid is oxidized deeply to acetic acid, CO and CO2 by the activated oxygen species on the surface' Acetic acid and the other C-C bond broken products are formed by various proton-catalyzed reactions. We, therefore, need to control the protonic property of the catalyst in order to achieve much higher selectivities to unsaturated oxygenate products. This work was supported by the Grant-in-Aids for Scientific Research on Priority Areas from the Ministry of Education, Science, Sports and Culture of Japan. REFERENCES

1. 2. 3.

Y. Moro-oka and W. Ueda, Catalysis, Royal Society of Chemistry, Vol 11 (1994) 223. Cavani F., and Trifiro, F., Catalysis, Royal Society of Chemistry, Vol. 11 (1994) 247. R.A. Scheldon, "The Activation of Dioxygen and Homogeneous Catalitic Oxidation", Ed. by D. H. R. Barton et al., Plenum Press, New York, (1933) 9. 4. H.H. Kung, Ind. Eng. Chem. Prod. Res. Dev., 25 (1986) 171. 5. B.K. Hodnett, Catal. Rev.-Sci. Eng., 27 (1985) 373. 6. G.J. Hutchings, Chem. Soc. Rev., 18 (1989) 251. 7. Y. Amenomiya, Catal. Rev.-Sci. Eng., 32 (1990) 163. 8. G.N. Kastanes, G.A. Tsigdinos and ,l. Schwank, Proc. A.I.Ch.E., Mtg, 172 (1983) C4. 9. E.M. Thorsteinson, T.P. Wilson, F.G. Young, and P.H. Kasai, ,l. Catal., 52 (1978) 116. 10. M. Roy, M. Gubelmann-Bonneau, H. Ponceblanc, and J. C., Volta, Catal. Lett., 42 (1996) 93. 11. M. T. Pope, "Heteropoly and Isopoly Oxometalates", Springer-Vedag, New York (1983). 12. R. J. J. Jansen, et al, Recl. Tray. Chim. Pays-Bas, 113 (1994) 115. 13. K. Harold, European Patent, 0010 902 (1979). 14. J. B. Black, J. D. Scott, E. M. Serwick and ,l. B. Goodenough, J. Catal., 106 (1987) 16. 15. E. M. Serwick, J. B. Black, and J. B. Goodenough, J. Catal., 106 (1987) 23. 16. G. B. McGarvey, and J. B. Moffat, J. Catal., 132 (1991) 100. 17. J. P. Battek, et al., US Patent 5,198,580 (1993). 18. A. Aboukais, D. Ghoussoub, E. Blouet-Crusson, M. Rigole and M. Guelton, Appl. Catal. A: General, 111 (1994) 109. 19. N. Mizuno, M. Tasaki and M. Iwamoto, Appl. Catal. A: General, 118 (1994) L1. 20. S. Albonetti, F. Cavani, F. Tifiro, M. Gazzano, M. Koutyrev, F. C. Aissi and A. Aboukais, J. Catal., 146 (1994) 491. 21. N. Mizuno, M. Tasaki and M. Iwamoto, Appl. Catal. A: General, 128 (1995) L165. 22. F. Cavani, E. Etienne, M. Favaro, A. Galli and F. Trifiro, Catal. Lett., 32 (1995) 215. 23 N. Mizuno, M. Tasaki and M. Iwamoto, ,l. Catal., 163 (1996) 87. 24. W. Ueda and Y. Suzuki, Chem. Lett., (1995) 541. 25. W. Ueda, Y. Suzuki, W. Li and S. Imaoka, 1 lth Inem. Congr. Catal., Baltimore (1996) 1065. 26. E. M. Thorsteinson, T. P. Wilson, F. G. Young and P. H. Kasai, J. Catal., 52 (1978) 116. 27. M. Merzouki, B. Taouk, L. Monceaux, E. Bordes and P. Courtine, Stud. Sur. Sci. Catal., Vol. 72 (1992) 165. 28. L. Tessier, E. Bordes and M. Gubelmann-Bonneau, Catal. Today, 24 (1993) 335. 29. M. Ai, J. Catal., 101,389 (1986). 30. S. Yamamatsu and T. Yamaguchi, .lap. Patent 02-042,032 (1990). 31. Z. Sojka, Catal. Rev. - Sci. Eng., 37 (1995) 461. 32. Payen, E., Kasztelan, S., Moffat, J. B., J. Chem. Soc. Faraday Trans., 88 (1992) 2263. 33. K. Bruckman and ,l. Haber, in: Adv. Catalyst Design, Eds. C. N. R. Rao, and M. Graziani, World Scientific, Singapore, (1994) 111.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

443

The role of metal oxides as promoters of V205/y-AI203 catalysts in the oxidative dehydrogenation of propane J.M. L0pez Nieto a, R. Coenraads a'l, A. Dejoz b and M.I. Vazquez b a) Instituto Tecnologia Quimica, UPV-CSIC, Avda. Los Naranjos s/n, 46071-Valencia (Spain), Fax: 6-3877809; b) Departamento de Ingenieria Quimica, Universidad de Valencia, Burjassot (Spain).

Summary The physicochemical properties of potassium-, bismuth-, phosphorous- and molybdenum-doped (Me/V atomic ratios of 0 to l) V2Os/7-AI203 catalysts and their catalytic behavior in the oxidative dehydrogenation of propane have been compared. The incorporation of metal oxides modifies the catalytic behavior of alumina-supported vanadia catalysts by changing both their redox and their acid-base properties. In this way, the addition of potassium leads to the best increase in the selectivity to propylene. This performance can be related to the modification of the acid character of the surface of the catalysts. The possible role of both redox and acid-base properties of catalysts on the selectivity to propylene during the oxidation of propane is also discussed. 1. INTRODUCTION

Supported vanadium oxides have been proposed as selective catalysts in partial oxidation reactions [1] and more specifically in the oxidative dehydrogenation (ODH) of short chain alkanes [2, 3]. However, it has been observed that the catalytic behavior of these catalysts during the oxidation of alkanes depends on the vanadium loading and the acid-base character of metal oxide support. In this way, aluminasupported vanadia catalysts with low V-loading are highly active and selective during the ODH of ethane [4-7] and propane [8] but they show a low selectivity in the ODH of n-butane [4, 5, 9, 10]. The catalytic behavior of vanadium-based catalysts can be generally modified with the addition of a second element which acts as promoter. In this paper we present a comparative study on the catalytic properties in the ODH of propane of undoped and Me-doped AI203-supported vanadium catalysts, in which metal oxides with redox and/or acid-base properties (K, Bi, Mo, P) have been used as potential promoters. 1On leave from the Eindhoven University of Technology

444 Table 1. Characteristics of undoped and doped alumina-supported vanadia catalysts. Samplea VIAl Bi(0.2) Bi(1.0) K(0.1 ) K(0.7) P(0.5) P(1.0) Mo(0.2) Mo(0.9)

TPR results T m (oc) 465 n.d. 360, 520 490 385, 500 470 485 n.d. 463, 772

Relative acidityb Lewis Br~nsted 1.0 1.0 0.5 0.4 0.4 0 0.1 0 0 0 0.6 0 n.d. n.d. 0.6 0.7 0.5 2.4

V5+-species (DR)C (VO3) n n.d. VO4 + VO6 (VO3)n ; VO4 (VO3)n ; VO4 n.d.. (VO3)n ; VO4 n.d. (VO3)n ; VO4

a) In parenthesis the Me/V atomic ratio, b) The number of acid sites was determined from the intensity of the corresponding band in the IR sprectrum of pyridine adsorbed, c)VO4 and VO6 are isolated tetrahedral and octahedral V 5§ species, respectively, while (VO3)n indicates polymeric species.

2. EXPERIMENTAL 2.1. Catalyst preparation Alumina-supported vanadia catalyst (V/AL) was prepared by "wet" impregnation method of a Girdler T126 7-AI203 support (SBET = 188 m2 g-l), with an ammonium metavanadate solution at a pH of 7 [5]. The concentration of the vanadia solution was selected in order to achieve a final vanadium loading of 3.5 wt % of vanadium atoms. The sample was calcined at 600~ for 6 h. Mo-, P-, Bi- and K-doped catalysts were prepared by impregnation of the alumina-supported vanadia catalyst with an aqueous solution of a salt of the corresponding metal. The amount of metal oxide was varied in order to obtain MeN atomic ratios between 0 and 1. All the catalysts tested were previously calcined at 600~ during 6 h. The catalysts will be named with the metal promoter and, in parenthesis, the MeN atomic ratio. Table 1 shows the physico-chemical characteristics of some of the studied catalysts. 2.2. Catalyst Characterization Characterization of catalysts was carried out by several techniques, i.e. BET, Diffuse reflectance in the UV-vis region (DR), 51V MAS-NMR, TPR-H2 and FTIR of adsorption/desorption of pyridine [3-5, 10]. 2.3. Catalytic test The catalytic experiments were carried out in a fixed bed quartz tubular reactor at atmospheric pressure in the 500-550~ temperature interval. The catalyst charge was 0.2-1.0 g mixed with 8 g of Norton Silicon Carbide. The feed consisted of propane,

445 oxygen and helium mixture in a molar ratio of 4/8/88. The analyses of reactant and products were carried out by gas chromatography, using two column types: i) Porapak Q (3.0 m x 1/8 in); ii) molecular sieve 5A (1.5 x 1/8 in). Blank runs showed that under the experimental conditions used in this work the homogeneous reactions could be neglected [3]. 3. RESULTS AND DISCUSSION

XRD patterns of calcined samples indicated the absence of crystalline materials, except in the case of Bi-doped catalysts in which crystalline BiVO4 was observed. The TPR profiles of doped and undoped alumina supported vanadium oxide catalysts are shown in Figure 1, while the temperatures of the maximum hydrogen consumption, Trn, are shown in Table 1. One peak with a maximum at 465~ is observed in the undoped sample (Fig. la)in agreement with previous results [5]. H2-uptake (a.u.)

Figure 1. TPR profiles of aluminasupported catalysts: a) V/AI; b) K(0.1); c) Z(0.7); d) P(0.5); e) P(1.0); f) Mo(0.9); g) Bi(1.0) and h) 30wt% of Bi203/A1203.

d

I

9 200

" 400

600

800

T e m p e r a t u r e (~

One reduction peak is also observed in the TPR patterns of K- (Fig.1 b-c) and Pdoped (Fig.1 d-e) catalysts. The addition of potassium increases the temperature of maximum hydrogen consumption from 465~ (in undoped VIAl sample) to 500~ for the K(0.7) sample, indicating that the reducibility of V5+ species in K-doped catalysts is lower than in the undoped sample. However, although the temperature of maximum H2consumption also increases with the phosphorous loading, this effect is smaller than on K-containing catalysts. On the other hand, two peaks are observed in the TPR profiles of Bi- and Mocontaining samples. The TPR pattern of the Mo(0.9) sample shows two peaks with temperatures of maximum hydrogen consumption at 463~ and 772~ (Fig.If). The first peak corresponds with the reduction of V5+, which appears at the same temperature as

446 in the undoped sample, while the second peak is related to the reduction of Mo6+ species. In the case of Bi-doped catalysts the peaks appear at 360~ and 520~ respectively (Fig. l g). Alumina-supported Bi203, free of vanadium, shows two peaks with maxima 310~ and 435~ (which correspond to a reduction from Bi3+ to Bi0) (Fig. l h). Thus, the peak at 360~ probably corresponds to the reduction of Bi3+ in BiVO4 (with a reducibility lower than in Bi203), whereas the peak at 520~ is related to the reduction of both Bin+ and V5+ species. In addition, it appears that the presence of bismuth reduces the reducibility of the V5+ species in the catalyst, as the temperature of the maximum hydrogen consumption related with V5+ reduction shifts from 465~ to 520~ The average oxidation state (AOS) of the vanadium species after TPR can be derived from the amount of hydrogen consumed by the catalysts during the experiment. In this way, it can be noticed that, AOSs near to 3 have been obtained in all the catalysts studied. DR spectra in the UV-vis region of the undoped and doped samples are shown in Figure 2. In all cases, the presence of an absorption band at about 270-320 nm indicates the presence of isolated and/or polymeric VO4 tetrahedra [4,5]. Only in the case of the Bi(0.9) sample a second band can be seen at 475 nm, which could be related to the presence of V 5§ species in octahedral coordination [12].

a

a.u.

200

300

400 500 Wavelength, nm

600

Figure 2. DR spectra of alumina supported vanadia catalysts: a) V/A1; b) K(0.7); c) P(1.0); d) Mo(0.9); and e) Bi(1.0). The 51V NMR spectra under static conditions of K-doped catalyst with a KN atomic ratio of 0.7, are shown in Figure 3. The wideline spectrum shows a small peak at c.a. -270 ppm, which suggests the minoritary presence of octahedral V 5§ [13-15]. Two more components are observed at -530 ppm and -900 ppm. Under MAS conditions the main line appears at -583 ppm. This chemical shift value and the wideline spectrum with two components a t - 5 4 0 ppm and -900 ppm, probably corresponding to (~2 and (~3 respectively, suggest the presence of tetrahedral V 5§ of Q(2)

447

type like in metavanadates NH4VO3 or NaVO3, in agreement with the assignment made by Eckert and Wachs [13]. Since a similar spectrum is also observed in the undoped sample [4], it can be concluded that in the catalysts studied here, the vanadium species are mainly from tetrahedral VO4 chains. In addition, we can also conclude that in the range of the potassium contents studied, the incorporation of potassium does not modify the nature of the vanadium species.

I

0

,

I

-300

,

I

-600 5(ppm)

,

I

-900

,

I

-1200

F i g u r e 3. 51V wideline spectrum of potassium doped V2OJA12Oa with a K/V atomic ratio of 0.7.

Infrared spectrum of pyridine adsorbed on undoped alumina-supported vanadium oxide catalyst, after evacuation at 150 ~ shows an absorption band at 1450 cm-1 characteristic for pyridine retained on Lewis acid sites, which has been related to V-free alumina [4,10, 11]. The intensities of the bands at 1450 cm-1 (related to Lewis acid sites) and 1545 cm-1 (related to Brbnsted acid sites) have been used to determine the numer of Lewis and Brensted acid sites on the surface of catalysts. The results are outlined in Table 1. The intensity of the band at 1450 cm-1 observed in the spectrum of the undoped alumina-supported vanadia catalyst decreases upon the addition of bismuth, potassium, phosphorous or molybdenum. These results indicate a reduction of the number of Lewis acid sites after the incorporation of metal oxides. In the case of K-doped catalysts the low intensity of the band at 1450 cm-1 clearly demonstrates that the majority of the surface acid sites disappears with the incorporation of potassium. Thus, Lewis acid sites have completely disappeared in the K(0.7) catalyst. In the case of the molybdenum doped samples the increase of the intensity of the band at 1545 cm-1 with the molybdenum loading is striking, and can be attributed to a major presence of Brensted acid sites on the molybdenum doped samples compared with the undoped catalyst. This could be explained by the interaction between Mo6+ species and the alumina support sites free of vanadium.

448 3.2. Oxidation of propane

The variation of the selectivity to the main reaction products, i.e. propylene, CO and CO2, with the propane conversion during the oxidation of propane on the VIAL sample is shown in Figure 4. The selectivity to propylene does not depend on the reaction temperature but strongly depends on the conversion of propane.

........B.--

60 %%

>~ 40 ....

>

,,41.-I

O

(U

-~ 20 O0

0

' .9 - , ,,,,

0

.

I

10

I

I

20

I

I

30

i

I

4O

I

_ . _

I

50

Conversion, % F i g u r e 4. Variation of the selectivities to propylene (D); CO J ) and CO2 (V) in the ODH of propane on a VzOJT-A1203 catalyst (3.5 wt% V-atoms) at 500-500~ Experimental conditions in text.

Since a similar trend is observed on all the catalysts studied it can be concluded that propylene is a primary unstable products while CO and CO2 can be considered as primary and secondary products. However, oxygenated products other than COx were not observed. Figure 5 shows the variation of the catalytic activity for the oxidation of propane at 500~ with the reducibility of V-based catalysts, determined as the inverse of the temperature of maximum H2-consumption (1/Tm). A parallelism between the reducibility of catalyst and the catalytic activity for propane conversion similar to those observed on supported vanadium oxide catalysts [1, 10] can be proposed. On the other hand, the selectivity to propylene varies with the amount and type of the metal oxide added. Figure 6 shows the variation of the selectivity to propylene with the MeN atomic ratio, at a propane conversion of 10%. In all cases, the doped catalysts with a low MeN atomic ratio show a higher selectivity to propylene than the undoped sample, although K-containing catalysts, specially the sample with a KN atomic ratio of 0.7, are the most selective ones.

449

60

9 Bi.1 [3

K.0.7

O

K.0.1

I

P.1

o

A

Bi.0.2

o

V

Mo.0.2

-~ 20

+

V.AL

A

Mo.0.9

~7

P.0.5

40

>

<

o

0

1,28

1,24

1,32

1,36

1,40

Reducibility, 1/Tm Figure

5.

Variation of the catalytic activity (in gcat h/mol-C4) with the reducibility of catalysts (1/Tin).

80 o ...m o ~

9

.o~

60

9

9

i

/"

,,

j

9

"

'

> o

9

..o

o

HI""

~,

o-~

. o ~

s" ~''

['1.,.

" , ,..,, .., ,,.,

9

..,..,_,

t_.J ......

9 .,..

................

"',,

I

,

I

0,2

,

I

0,4

r-i

~'~

,,

"V ..................................

0,0

9

" ............

. ..,...

40

9

,

I

0,6

I

,

0,8

9

,

I

1,0

Me/V atomic ratio F i g u r e 6. Variation of the selectivity to propylene with the M e N atomic ratio of Me-doped catalysts at a propane conversion of 10%. Me= K J ) ; Bi (O); P (CI); and Mo(V).

450

It has been proposed that the acid-base nature of catalyst and support can modify the adsorption/desorption of reactants and products [2,3,5]. In this way, and according to both the catalytic and the physicochemical properties, it can be concluded that the incorporation of potassium initially increases the selectivity to propylene as a consequence of lowering the number of acid sites. Basic sites, in addition to acid sites, can be also proposed in Bi-doped catalysts. In addition, it has been observed that V-Bi-O catalysts are selective in the ODH of propane [16]. However, in our case, Bi-doped catalysts are not selective. This apparent contradiction can be explained when considering the presence of Bi3§ probably vanadium free, with a high reducibility which can favor the deep oxidation reactions. On the other hand, acid sites have been observed on Mo- and P-doped catalysts, specially on those with high MeN atomic ratios, and they show a selectivity to olefins lower than the undoped samples. Thus the variation of the selectivity to propylene from propane on metal oxide-doped catalysts could be explained on the basis of the acid-base character of catalysts. The addition of potassium to AI203- [5] or TiO2-supported [17-19] vanadia catalysts increases the selectivity to olefins during the ODH of n-butane and propane, respectively. This variation of the selectivity to oxydehydrogenation products has been related to a decrease in the number of Lewis acid sites [5] or to a decrease in the heat of the propylene adsorption [17-19]. In order to determine the influence of the addition of potassium on the catalytic performance of K-doped catalysts, Figure 7 shows the variation of the yield of propylene with the contact time (W/F) at 550~

16 12 4,,P ~

o-.e -o-

8

.er"

4 0

0

20

40

60

Contact time, W/F F i g u r e 7. Variation of the yield of propylene with the contact time (W/F) in the oxidation of propane at 550~ on the undoped V/AL catalyst (~1); and Kdoped catalysts with a K/V atomic ratio of 0.1 (0); 0.7 (V) and 0.9 (O).

451

At low contact time, the higher the K-content on the catalysts the lower the yield of propylene (Fig. 7). Since the formation rate of propylene is related to the Y~/(W/F) ratio, it can be concluded that the incorporation of K decreases the formation rate of propylene. However, the best yields of propylene are obtained on K-doped catalysts at high contact times (Fig. 7). This indicates that a lower rate in the combustion of propylene (consecutive reaction) is achieved on K-doped sample. In addition, since the maximum yield of propylene increases with the potassium content it can be concluded that on Kdoped samples, the rate of consecutive reactions decreases in a longer extension than the formation of propylene, favoring the obtention of higher selectivites to ODH products (Fig. 6). Thus, the incorporation of potassium, remplacing Lewis acid sites on V205/AI203 catalysts, would eliminate nonselective sites in the deep oxidation of propylene. Only in the case of samples with KN atomic ratios higher than 0.7, potassium oxide or potassium vanadate could be formed which could favor a low formation rate of propylene but also a high deep oxidation of propylene. In fact, a low selectivity to propylene has been observed on sample K(0.9) (Fig. 6), indicating that carbon oxides are directly formed from propane on samples with high K-contents. In the case of the addition of Mo and P, the selectivity to propylene initially increases with the metal incorporation, probably as a consequence of the decrease of the number of Lewis acid sites of the support. However, the selectivity to propylene decreases with the Me-loading. Since the appearance of Br5nsted acid sites have been observed in both P- and Mo-doped samples, it could be tentatively concluded that the presence of acid sites does not favor the formation of olefins. It has been suggested that the incorporation of alkali metals on TiO2-vanadia catalysts decreases both the V=O stretching frequencies and their polarizing power, while the incorporation of acid anions produces an opposite trend [20]. In addition, the presence of alkali ions decreases the heat of the propylene adsorption [17,18, 21]. Thus the different catalytic behavior of doped alumina supported vanadia catalysts, could be explained on the bases of the influence of the acid-base character of catalysts on the adsorption/desorption of propane and propene. In any case, the redox properties must be also considered. In this way, it will be interesting to study if, realy, a lower reducibility of the active sites could favor a lower rate of the consecutive reactions, as it has been observed in the case of K-doped catalysts. 4. CONCLUSIONS

In conclusion, this paper shows the effect of the addition of different metal oxides (K, Bi, P and Mo) on the catalytic behavior of an alumina-supported vanadia catalysts in the ODH of propane. In all cases, the addition of small amounts of metal oxide (MeN atomic ratio of 0.1) increases the selectivity to propylene, probably as a consequence of the elimination of non selective sites (Lewis acid sites) on the surface of the support. However, only in the case of K-doped catalysts the selectivity and the yield of propylene increases with the metal content. The varition of the acid-base character of catalysts and its influence on the adsorption/desorption of reactants and products could be responsible of the different performances observed. In this way,

452

potassium can be considered as a promoter of supported vanadia catalysts, not only in the case of the ODH of n-butane [5] but also in the ODH of propane. ACKNOWLEDGEMENT

Financial support by Comisi6n Interministerial de Ciencia y Tecnologia, CICYT, from Spain (Project MAT 94-0898) is acknowledged. REFERENCES

1. 2. 3. 4 5 6. 7. 8. 9.

10. 11. 12 13. 14. 15. 16. 17. 18. 19. 20. 21.

G.C. Bond and S.F. Tahir, Appl. Catal. 71 (1991) 1. T. Blasco and J.M. L6pez Nieto, Appl. Catal. A: General, (1997) in press. P. Concepci6n, A. Galli, J.M. L6pez Nieto, A. Dejoz and M.I. V&zquez, Topics in Catal. 3 (1996) 451. T. Blasco, A. Galli, J.M. L6pez Nieto and F. Trifiro, J. Catal. 168 (1997) in press. A. Galli, J.M. L6pez Nieto, J.M., Dejoz, A. and Vazquez, M.I., Catal. Lett. 34 (1995) 51. J. Le Bars, A. Auroux, S. Trautmann and M. Baerns, in Proceeding DGMKConference "Selective Oxidation in Petrochemistry", Ber.-Dtsch. Wiss. Ges. Erdoel, Erdgas Kohle, Tagungsber, 1992, p. 59. J. Le Bars, A. Auroux, M. Forissier and J.C. Vedrine, J. Catal., 162 (1996) 250. J.G. Eon, R. Olier and J.C. Volta, J. Catal., 145 (1994) 318. P.J. Andersen, and H.H. Kung, in New Frontiers in Catalysis, (L. Guczi, F Solymosi and P. Tetenyi, Editors), Studies in Surface Science and Catalysis, Vol. 75, Elsevier, Amsterdam, 1993, p. 205. T. Blasco, J.M. L6pez Nieto, A. Dejoz and M.I. Vazquez, J. Catal., 157 (1995) 271. J. Le Bars, J.C Vedrine, A. Auroux, S. Trautmant and M. Baerns, Appl. Catal. A: General, 119 (1994) 341. K.V.R.Chary and G. Kishan, J. Phys. Chem., 99 (1995) 14424. H. Eckert and I.E. Wachs, J. Phys. Chem., 93 (1989) 6796. T. Blasco and J.M. L6pez Nieto, Colloid Surface A: Phys. Eng. Aspects, 115 (1996) 187. O.B. Lapina, V.M. Mastikhin, L.G. Simonova and Yu O. Bulgakova, J. Molec. Catal., 69 (1991 ) 61. A.Corma, J.M. L6pez Nieto, N. Paredes, M. P~rez, Y. Shen, H. Cao and S.L. Suib, Stud. Surf. Sci. Catal., 72 (1992)213. R. Grabowski, B. Grzybowska, J. Sloczynski and K. Wcislo, Topics in Catal., 3 (1996) 277. R. Grabowski, B. Grzybowska, K. Samson and J. Sloczynski, Appl. Catal. A: General, 125 (1995) 129. D. Courcot, A. Ponchel, B. Grzybowska, Y. Barbaux, M. Rigole, M. Guelton and J.P. Bonnelle, Catal. Today, 33 (1997) 109. G. Ramis, G. Busca,and F. Bregani, Catal Lett., 18 (1993) 299. C. Martin, V. Rives and A.R. Gonzalez-Elipe, J. Catal.,114 (1988) 473.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

453

A l k a n e O x i d a t i o n o v e r B u l k and S i l i c a - S u p p o r t e d VO(HzPO4)z-Derived Catalysts

G.K. Bethke, D. Wang a, J.M.C. Bueno b, M.C. Kung and H.H. Kung Center for Catalysis and Surface Science, Ipatieff Laboratory, Northwestern University, Evanston, I1 60208, U.S.A. aon leave from Department of Chemical Engineering, Tianjin University, Tianjin, 300072, P.R. China UDEQ/UFSCar, Caixa Postal 676, 13560, Sao Carlos(SP), Brazil ABSTRACT XRD and LRS were used to characterize both supported and unsupported catalysts derived from VO(H2PO4) 2. Supported VO(H2PO4)2-derived catalysts were found to be more active but slightly less selective in butane oxidation to maleic anhydride than unsupported samples, but the difference in selectivity could be eliminated by adding a small amount of phosphorus to the supported samples. For butane oxidation, the activity of the catalysts was much enhanced by the presence of ~I-VOPO4. The relative rates of oxidation of propane, butane, and pentane were consistent with alkane activation occurring predominantly at the secondary carbon. 1. INTRODUCTION The selective oxidation of alkanes is a desirable prospect due to their potentially low environmental impact and the relatively low cost of raw material. Accordingly, the catalysts for such reactions are of great industrial interest. Vanadium-phosphorus oxides are one such group of compounds which have been heavily studied due to their commercial use in the selective oxidation of butane to maleic anhydride. Although the bulk phase of the active commercial catalyst is (VO)2P2OT, it has been reported that VO(PO3)2-containing catalysts are more selective, but less active [ 1]. Up to a conversion of 8%, the VO(PO3) 2 catalysts could produce maleic anhydride and furan with a combined selectivity of 100%. In our previous studies of butane oxidation over silica-supported vanadium-phosphorus oxide (VPO) catalysts, it was observed that the selectivity for maleic anhydride increased with the P/V ratio, from about 25% for catalysts with a P/V ratio of unity to about 50% for a P/V ratio of about two [2,3]. In order to more fully understand the properties of VPO catalysts with P/V ratios near 2, we have prepared and characterized silica-supported and unsupported catalysts derived from VO(H2PO4)2, the precursor to VO(PO3)2, and studied the selective oxidation of a series of alkanes ranging from C3H8 to C5H12 over the bulk catalyst.

454 2. EXPERIMENTAL V O ( H 2 P O 4 ) 2 w a s prepared by an aqueous procedure based on the method reported by Sananes et al. [ 1] Aqueous H3PO 4 w a s combined with V204 and allowed to sit overnight. The deep blue, glassy mixture was then refluxed between 180~ and 200~ for 1 to 6 h until the mixture became predominantly pale blue. The solution was then washed with diethyl ether and centrifuged to remove e x c e s s H3PO 4. The pale blue solid was dried at 100~ overnight. It was observed that in order to consistently obtain the pale blue solid, the starting V204 had to be free of V 5+. If a significant amount of V 5+ was present, a green solid would form. Catalyst A was prepared by impregnation of Cab-O-Sil L-90 silica with an aqueous solution ofVO(HRPO4) 2 and allowed to dry overnight at 40~ - 50~ Catalyst B was prepared identically except that additional H3PO4 was added to the impregnation solution to increase the P/V ratio by 0.08. The amount of VO(H2PO4)2 used in the impregnation yielded nominal loadings of 8.5 wt% and 8.4 wt% V, respectively. Both supported and unsupported catalysts were formed by activating the precursors on-stream in a C4H~0/OR/He (2/24/74) mixture flowing at 70 ml/min. The temperature was increased from room temperature to 100~ at 2~ then held at 100~ for 1 h. The temperature was then increased to the reaction temperature at 10~ Precursors and catalysts were characterized in ambient conditions by X-ray diffraction (XRD) on a Rigaku Powder Diffractometer using CuK~ radiation with a Ni filter. LiF was used as an internal standard for the activated catalysts. Laser Raman spectra (LRS) were collected using Ar ion laser excitation (514.5 nm) at a power of 25 mW at the sample. Spectra for the precursors were collected in ambient conditions, and reaction-used catalysts were characterized in-situ at 400~ in a 70 ml/min flow of C4H~0/O2/He (0.99/10.2/88.81). Phosphorus to vanadium ratios (molar) were determined by inductively coupled plasma (ICP). Diffuse reflectance spectra (DRS) were collected in ambient conditions using polytetrafluoroethylene as a reference.

3. RESULTS AND DISCUSSION Figure 1 shows the XRD pattern, of a representative precursor sample. With the exception of the small peaks at 12.6 ~ 25.4 ~ and 39'0 ~ the peak positions are in agreement with those presented in the literature for VO(H2PO4) 2indexed to a tetragonal cell with dimensions of a=8.956/~ and c=7.967 A [4,5]. The presence of the unidentified peaks suggest an impurity phase(s). The relative intensities of all the peaks varied slightly from both the literature as well as from one batch of precursor to another. This was most noticeable for the peak at 20= 14~ corresponding to the [ 100] plane of VO(H2PO4)2. This slight difference in intensities, however, was not correlated with either the reaction data or the LRS which were identical for the precursors. The LRS of the precursors (figure 2) also agreed with the literature spectra of VO(H2PO4) 2 with the main peak at 937 c m l (P-O stretch of the phosphate group) and other minor bands at 1146 cm -~, 1105 cm l, 902 cm -1, 576 cm ~ and 221 cm ~ [1,6]. Three smaller peaks at 339 cm l, 371 cm ~ and 518 cm ~ were also present but have not previously been reported. Figure 3 shows that the DRS of the precursor matched that of V O ( H 2 P O 4 ) 2 with three broad bands at 205 nm, 378 nm and 618 nm. The bands at 618 nm and 378 nm can be assigned to the promotion of the single 3dxy electron in a 2B2 state to the 2B 1 and 2A 1 excited states, respectively; the band at 205 nm can be assigned to charge transfer [4]. ICP analysis determined the P/V ratio of a typical precursor to be 2.05 & 0.08.

455

,--2,.

.,..a

(D

.,..a

10

15

20

25

30

35

40

45

50

2-Theta

Figure 1" XRD of the precursor, VO(H2PO4) 2.

.,..a ~ ,.,,~

.,..a

~o'oo 86o

66o

480

200

Wavenumber (cm 1) Figure 2" LRS of the precursor, VO(H2PO4)2.

456

800

-

600 -~

,.o

400 200 0

200

4;0

6;0

800

Wavelength (nm) Figure 3: DRS of the precursor, VO(H2PO4) 2. The activity and product selectivity for the oxidation of C3H8, C4H10 and C5H12 over the unsupported catalyst are shown in Table 1. The product distribution varied according tothe reactant alkane. For butane and pentane oxidation, maleic anhydride selectivities were equal to or greater than 50% and changed little with conversion. In addition, for pentane oxidation, small amounts of phthalic anhydride were also formed. Propane oxidation was quite nonselective with 87% of the products being carbon oxides. The activity was roughly proportional to the number of secondary carbons in the alkane and increased from C3H8 to C4H8 to CsHn. These results agree with those presented by Patel et al. in their study of the oxidative dehydrogenation of ethane, propane and butane over V-Mg-O and Mg2V207 [7]. The proportionality between secondary carbon number and activity implies that activation of alkanes occurs predominately at the secondary carbon atom. Table 2 compares the reaction data for butane oxidation over supported and unsupported catalysts. Supporting the precursor on silica caused a decrease in maleic anhydride selectivity, but an increase in the conversion at both 425~ and 485~ However, with the addition of extra phosphorus to the supported catalyst, the selectivity could be nearly recovered with no loss in the increased conversion. Figure 4 shows the XRD spectra for both the bulk catalyst and the supported catalysts. The spectrum of the bulk catalyst shows a broad feature centered around 20=22 ~ indicating the presence of material amorphous to XRD. The crystalline portion produced a peak at 20=23.9 ~ which can be assigned to the overlap of the two major peaks of VO(PO3)2, 20=23.1 o and 24.2 ~ [1,6]. Two remaining peaks at 20=14.5 ~ and 29.2 ~ can be attributed to 0~l-WOPO4 (20=14.3 ~ and 28.8~ although the peak at 20=21.6 ~ is missing [8]. Alternatively, since the XRD was collected in ambient conditions after reaction, these peaks could be attributed to VOPOao2H20 (20=28.8 ~, 24.0 ~ and the hydration peak at 20<15 ~) [8]. The presence of amorphous material

457

TABLE 1: Activity and Selectivity for Alkane Oxidation over bulk activated VO(H2PO4)2 a Carbon Selectivity b, (%) Hydrocarbon ~

Temp (~

Conv (%)

MA

PA

Ar

C2

C3

CO

CO2

C3H8

450

4

0

0

6

7

0

74

13

C3H8

485

11

0

0

5

8

0

70

17

C4H10

450

11

52

0

5

2

trace

28

13

C4H10

485

22

51

0

4

3

~ace

31

11

C5H12

450

15

50

1

5

1

1

29

13

CsH12

485

29

53

1

1

1

1

32

11

a) Total flowrate = 50 cc/min, 0.80 g catalyst after reaction, P/V = 2.0 before reaction. b) MA - Maleic anhydride, PA = Phthalic anhydride, Ar = Acrylic acid. c) Hydrocarbon concentration in feed: 0.9% for C3H8, 1.5% for C4H10 and CsH12, HC/O2- 1/15, balance He.

TABLE 2: Butane Oxidation over Activated Silica-Supported VO(H2PO4)2a Carbon Selectivity b, (%) Catalyst

P/V

Catalyst Wt c (g)

Temp (~

Conv (%)

MA

Ar

C2

C3

CO

CO2

A

2.05

0.86

425

6

43

4

3

trace

39

11

A

2.05

0.86

485

25

48

1

3

trace

37

11

B

2.13

0.89

425

6

57

trace

1

1

31

10

B

2.13

0.89

485

23

50

trace

1

2

36

11

Bulk

2.05

0.82

425

2

62

6

2

trace

24

6

Bulk

2.05

0.82

485

12

57

3

2

trace

29

9

a) Total flowrate = 70 cc/min, C4H10/OJHe = 2/24/74. b) MA - Maleic Anhydride, Ar = Acrylic acid. c) After reaction.

458

c,:i r~

i,,,--i

,. ., ., .,

10

I ., .| .l .|

15

I .I .l .l .l

20

.I .| .l .l

I

25

I

I,,,T

|. .l .l .|

| | l |

30

35

40

2-Theta Figure 4: XRD of the catalysts after reaction with butane. (a) Bulk catalyst, (b) Catalyst A, (c) Catalyst B, o=LiF (intemal standard).

r,r

I

1100

I

1000

I

900

I

800

700

Wavenumber (cm "1) Figure 5" LRS of the catalysts during reaction with butane at 400~ Catalyst A, (c) Catalyst B.

(a) Bulk catalyst, (b)

459 was also reported in some of the catalysts in ref. 1. Interestingly, in that report, it appears that the activity was the highest for the catalyst with the most amorphous material. The LRS spectrum of the bulk catalyst during reaction at 400~ is shown in Fig. 5a. The peaks were broad, consistent with the presence of a significant amount of amorphous material in the sample. The dominant peak was around 950 cm ~, which could be due to overlap of the 957 cm -~(P-O stretch) band of VO(PO3)2 and 928 cm ~ (symmetric Vs(PO4-3) stretching) band of 0~I-VOPO 4 [6,9]. There might be fine structures at 947 cm l and 960 cm -1, but they were close to the noise level. A second dominant feature was the broad peak around 1040 cm -~, which is possibly due to the V-O-P coupled stretching in Ix1-VOPO 4 [9]. Interestingly, this 1040 cm -1 peak was also observed in some of the catalysts in ref. 1, and it was larger for the more active catalysts. The low intensity of the 928 cm l band relative to the 1040 cm l band was consistent with reports in the literature showing the 928 cm ~ band of 0~I-VOPO 4 decreasing in intensity relative to the 1040 cm -1 band as the temperature increased [2,9]. Room temperature spectra taken before and after the in-situ spectrum of Figure 5a showed the 928 cm -~ band having a higher intensity than the 1040 cm l band. The silica-supported catalysts A and B were amorphous to XRD and yielded featureless spectra (Fig. 4b and c). Figure 5 shows their LRS spectra during reaction. Similar to that for the bulk catalyst, they showed two broad peaks around 950 and 1040 cm -~ and exhibited the same relative intensity trends at 928 cm -~ and 1040 cm -~ depending on the temperature. The similar structural and catalytic properties of the SiO2-supported and unsupported samples prepared from the same precursor suggest that the same active surface is formed on both types of samples. The higher conversions obtained with the supported samples could be attributed to higher dispersions of the VPO compounds. The slightly lower maleic anhydride selectivity observed for catalyst A than B or the bulk catalyst could be due to some phosphorus atoms interacting with the silica surface rather than with vanadium atoms, such that the P/V ratio is less than two in the VPO compounds. Addition of phosphorus to catalyst B replenished this lost phosphorus. Previous studies of supported vanadium-phosphorus oxides have shown that some phosphorus atoms can be associated with the silica [2,8]. The catalytic properties of the supported samples as well as the LRS are similar to the SiO2-supported P/V=2 VPO samples prepared previously [2,3]. These earlier samples were prepared by adding H3PO4 to P/V=I samples synthesized by various synthesis routes. Thus, for the supported samples, the method of preparation is much less important than the composition. Reports in the literature have shown that bulk VO(PO3) 2 is inactive towards butane and butene oxidation [6,10]. The samples presented in this study showed activity towards butane oxidation but contained ~-VOPO4 and amorphous material in addition to VO(PO3) 2. This implies that the activity of these samples is related to the presence of o~1-VOPO 4 or a VOPO 4like compound. However, the samples in this study showed a higher selectivity to maleic anhydride than bulk 0~-VOPO 4 [11 ]. Possibly, the inactive VO(PO3) 2 acts as a barrier to separate and disperse nonselective and highly active ~I-VOPO4 moieties, thus decreasing the production of complete combustion products while increasing the production of partial oxidation products. Such a model agrees with the data from the VO(PO3)2-containing catalysts reported by Sananes et al. in which the activity could be correlated to the amount of amorphous material present and the relative intensity of the 1040 cm -~LRS band [ 1]. In both the bulk and the supported samples studied here, the presence of ~-VOPO4 was detected under reaction conditions. Thus, it

460 appears that the presence of 0~I-VOPO4 is important for high activity. This is consistent with the recent report in which the production of maleic anhydride was directly correlated with the V 5+concentration in the catalyst under transient conditions [12]. That is, ~l-VOPO4 is the active component to activate butane. 4. CONCLUSIONS In conclusion, a comparison of the results of this study and literature results has shown that catalytically equivalent supported catalysts can be formed either by using VO(H2PO4)2 as a precursor or by adding phosphorus to P/V=I compounds to increase the P/V ratio to about 2. Characterization data suggest that these two types of supported catalysts may be identical. Finally, alkane oxidation over a bulk catalyst derived from VO(H2PO4)2 may proceed via activation at the secondary carbon atoms, and small amounts of dispersed ~l-VOPO4 are responsible for the activation of butane. ACKNOWLEDGMENTS Financial support was provided by the U.S. Department of Energy, Basic Energy Sciences. J.M.C. Bueno acknowledges the support from CNPq (Brasil) Conselho Nacional de Desenvolvimento Cientifica e Tecnologico. K.A. Bethke assisted by collecting DRS data, and S.M. Babitz and J. Shen assisted by collecting ICP data. REFERENCES [1] M.T. Sananes, G.J. Hutchings and J.C. Volta, J. Catal. 154 (1995) 253-260. [2] K.E. Birkeland, S.M. Babitz, G.K. Bethke, H.H. Kung, G.W. Coulston and S.R. Bare, submitted (1996). [3] J.M.C. Bueno, G.K. Bethke, M.C. Kung and H.H. Kung, Presented at the XV Simposio Iberoamericano de Catalisis, C6rdoba, Argentina, September 1996; paper [B-4]. [4] G. Villeneuve, A. Erragh, D. Beltran, M. Drillon and P. Hagenmuller, Mat. Res. Bull. 21 (1986) 621-631. [5] S.A. Linde and Yu. E. Gorbunova, A.V. Lavrov and V.G. Kusnetsov, Dokl. Akad. Nauk SSSR 224(1979) 1411. [6] V.V. Guliants, J.B. Benziger, S. Sundaresan, I.E. Wachs, J.-M. Jehng and J.E. Roberts, Catal. Today 28 (1996) 275-295. [7] D. Patel, M.C. Kung and H.H. Kung, in "Proceedings, 9th International Congress on Catalysis, Calgary, 1988," (M.J. Philips and M. Ternan, Eds.), Vol. 4, 1554-1561, Chem. Institute of Canada, Ottawa, 1988. [8] K.E. Birkeland, Ph.D. thesis, Northwestern University, 1995. [9] F. Benabdelouahab, R. Olier, N. Guilhaume, F. Lefebvre and J.C. Volta, J. Catal. 134 (1992) 151-167. [10] E. Bordes and P. Courtine, J. Catal. 57 (1979) 236-252. [ 11 ] T. Shimoda, T. Okuhara and M. Misono, Bull. Chem. Soc. Jpn. 58 (1985) 21632171. [12] G.W. Coulston, S.R. Bare, H. Kung, K. Birkeland, G.K. Bethke, R. Harlow, N. Herron and P.L. Lee, Science 275 (1997) 191-193.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

461

T h e N a t u r e of t h e A c t i v e S i t e of t h e (VO)2P~O7 C a t a l y s t : A n I n v e s t i g a t i o n of t h e C h e m i c a l C o m p o s i t i o n a n d D y n a m i c s of t h e Catalyst Surface B.Kubias ~, F.Richter b, H.Papp b, A.Krepel a, A.Kretschmer c a Institut fiir Angewandte Chemie Berlin-Adlershof, D-12484 Berlin, Germany b Institut fiir Technische Chemie, Universit~it Leipzig, D-04103 Leipzig, Germany c Institut fiir Chemie, Humboldt-Universit~it zu Berlin, D-10115 Berlin, Germany

The oxidation of n-butane to maleic anhydride (MA) has been investigated in the presence and absence of water vapour in the gas phase to study the influence of water on activity and selectivity of an equilibrated (VO)2P207 catalyst. It is shown that water vapour strongly influences the catalytic properties by blocking the active sites (,adsorption effect") and an enrichment of phosphate anions on the catalyst surface (,,structural effect"). Both effects are reversible on different time scales. A migration and a loss of phosphorus were found using thermogravimetric experiments and steam containing gases. The reason for these effects is discussed in terms of XPS and ISS results. A dynamic model of the surface has been developed describing the active site as a thin layer of single VOG octahedra linked to the (V 9 core in a chain-like fashion. This structure is thought to comprise only a few subsurface layers.

1. I N T R O D U C T I O N Up to now, (VO)2P207 based catalysts are the most selective ones in the oxidation of n-butane to maleic anhydride (MA) [1]. The reason for this unique property has been studied intensively during the past decade [2]. However, to gain reliable information on the nature of the active site has proven to be difficult. As an important fact, an enrichment in phosphorus on the surface of these catalysts was found but it turned out that the precise determination of the surface P:V ratio by XPS is problematic. As recently emphasized again, the determination of the XPS sensivity factor is absolutely necessary [3]. Furthermore, during the operation of a technical reactor a loss in phosphorus of the vanadylpyrophosphate catalyst was observed. The valence state of vanadium in the surface layers of equilibrated catalysts is between +4 and +5 whereas its state in the bulk is near +4 [4,5]. So far, it has been assumed t h a t the selective butane oxidation occurs on the (100) surface plane of the (VO)2P207 crystal. As recently reviewed [2], cutting the crystal parallel to (100) and considering different terminations of the surface leads to different surface models. The most

462 impressive model of the active site was proposed by Thompson and Ebner [6] and is based on the termination of the surface by pyrophosphate groups. Edge-shared vanadyl double octahedra are thought to be the active and selective vanadium sites located in clefts and fenced by pyrophosphate groups. However, the structural models mentioned fail in the explanation of certain findings, e.g., higher P:V ratios than predicted by the pyrophosphate termination model [3] were, frequently, measured by XPS. The oxidation of V 4§ seems to be restricted to the surface region and not extended to the bulk for (VO)2P207 catalysts prepared from aqueous medium [5]. Especially, the known models cannot explain the finding that the area-specific rates of MA formation are in the same order of magnitude on both (VO)2P207 and VO(PO~)2 catalysts [7-9] which strongly differ in their structural properties. Furthermore, the driving force for the phosphorus enrichment on the (VO)2P207 surface is not clear yet. In the study of Arnold and Sundaresan [10] on the influence of water on the catalytic properties of a non-equilibrated VPO catalyst the hydrophilicity of phosphorus was supposed to be a reason for it but this assumption has not yet been proven. It was found t h a t the addition of water vapour to the feed lowers the overall activity towards butane oxidation and enhances the selectivity towards partial oxidation products. The addition of steam accelerated the formation of (VO)2P207 in the solid structure which originally contained approximately equal amounts of a-VOPO4 and (VO)2P207 crystalline phases [10]. However, the (VO)2P207 phase is the only crystalline phase in an equilibrated VPO catalyst with a P:V ratio of 1.1 as used in [10]. Recently Cavani and Trifir6 emphasized [2] that only a study of equilibrated catalysts provides precise information on the relationships between activity and selectivity and the structural properties of vanadylpyrophosphate catalysts. Therefore, it was the aim of this study to investigate the influence of water vapour on the chemical composition of the surface and on the activity and selectivity of an equilibrated (VO)2P207 catalyst intending to get a better u n d e r s t a n d i n g of the dynamics of the catalyst surface and the nature of the active site of this catalyst.

2. E X P E R I M E N T A L

The catalyst precursor VOHPO4.0.5 H20 was prepared from aqueous solution as described in [5] but without the addition of H2SO4. It was calcined in N2 (fresh catalyst) and conditioned in a butane(1.5 %)- air mixture [5] to obtain a (VO)2P207 catalyst which showed the characteristics of an equilibrated catalyst described by Ebner and Thompson [11]. The average valence state of vanadium a m o u n t e d to 4.02, the P:V ratio was 1.0 and the specific surface area was about 10 m2/g. It is assumed that V 5§ ions formed in the conditioning procedure of the

463 fresh (V4+O)2P207 are only situated in surface and subsurface layers [5]. The activity and selectivity as well as the chemical composition of the surface and bulk of this catalyst did not change perceptibly with time on stream indicating the state of an equilibrated catalyst. Calcination, conditioning and catalytic reaction under steady state and t r a n s i e n t conditions were carried out in a microcatalytic tubular reactor using different feeds containing up to 30 vol % water vapour. The reaction products were determined by on line gas chromatography (MA and n-butane) and nondispersive IR photometry (CO, CO2). In addition, experiments were performed with catalyst samples pretreated by mixtures of Nz with water vapour. The P:V ratios in the surface and subsurface layers were determined by XPS and ISS, respectively. The influence of H20 on the solids was measured by in situ TG. X-ray photoelectron spectroscopy (XPS) and ion scattering spectroscopy (ISS) were performed with a Leybold Heraeus spectrometer (LHS 10) [12]. The exact peak positions and areas were determined after smoothing and linear background substraction (Shirley) by fitting curves using gaussian symmetry. For XPS MgK(~ radiation at 240 W was used. Binding energies were referred to the C ls line at 284.6 eV. To determine a reliable sensitivity factor freshly powdered VPO glasses prepared according to [13] were used as standard materials having uniform compositions of surface and bulk [14]. The real surface P/V ratios of the catalysts were then calculated using the peak areas of P 2p and V 2p3/2 and the determined sensitivity factor. For the O ls region the sensitivity factor of Briggs and Seah [15] was used. The O ls peaks of all vanadyl pyrophosphate catalysts were decomposed into two peaks (0 2. of the oxide and OH groups [16,17]). The difference in the binding energies of 9 ls (02) and V 2p3/2 was used to determine the oxidation state of vanadium [18]. ISS m e a s u r e m e n t s were performed with 4He+ ions at 1000 eV and an ion current of 5.7 nA/mm 2 at a pressure of 2 9 10 .7 mbar. The IS spectra were recorded between 300 and 1000 eV kinetic energy, the scattering angle was 120 ~ The relative ISS sensitivity factor for P to V was derived from IS spectra of pure V205 and P205. Taking the different densities of the V and P surface species into account a relative sensitivity factor for the peak areas of P to V of 0.18 resulted. The FWHM was always about 80 eV for P, 4 8 e V for V and 55 eV for O. The thermogravimetric measurements were performed in microbalance systems (CAHN TG-121 and Sartorius 4433). The experimental conditions were the following: sample mass: 100-150 mg, sample temperatures: 300~ to 700~ t e m p e r a t u r e ramps: 2 and 5 K/min. The relative error of the thermogravimetric results was <10%. For the in situ-measurements in the presence of water vapour a water bubbler was used. It was possible to switch between the pure N2 stream and the HzO/N2 mixture. The gas streams (H20/N2 and He in counterstream) were adjusted so t h a t water condensation in the balance chamber was avoided.

464 3. R E S U L T S AND D I S C U S S I O N Firstly the influence of water vapour on the catalytic properties of the equilibrated (VO)2P207 catalyst is reported. Afterwards the effect of water vapour on the enrichment in phosphorus of the catalyst surface and the reversibility of this effect is described using thermogravimetric experiments in the presence and absence of water vapour and XPS as well as ISS measurements of the respective catalyst samples. 3.1. T h e e f f e c t o f w a t e r v a p o u r o n t h e c a t a l y t i c p r o p e r t i e s o f (VO)2P207 In order to assess the influence of water vapour on the activity and selectivity of (VO)2P207 the oxidation of n-butane was studied in a transient experiment adding 20 vol% water vapour to the feed. After 3 h the addition of water vapour was stopped and the conversion of n-butane and the selectivity of MA formation has been followed for further 3 h. The results are presented in Figure 1.

60

~

9 ,--4

.

>~409 r

(D

20-

+ o% 2o 0

i

I

!

60

120

180

--I i

240 t, min

'1

300

!

t

360

420

Figure 1. Dependence of the conversion of n-butane ([~) and of the MA selectivity (A) on time on stream Feed: 1.46 % C4H10- air, T= 452~ x = 1.6 sec

As shown in Figure 1, the addition of water caused a marked decrease in butane conversion and also in MA selectivity. After the exposure to the wet butane-air feed both activity and MA selectivity increased but did not reach the original levels and remained constant at these lower levels for at least 3 h. Two effects of water vapour on the activity and selectivity of (VO)2P207 can be identified, having different time scales: a readily reversible effect occuring within a few seconds and a second effect which seems to be non-reversible in a time scale of hours. In agreement with Arnold and Sundaresan [10], the first effect is interpreted as an ,,adsorption effect" of water molecules blocking the active

465 v a n a d i u m sites. The decline in MA selectivity which is in contrast to their results in [10] may be due to the fact t h a t the authors investigated a non-equilibrated, (~VOPO4 containing catalyst [10]. The lower selectivity may be caused by an enhanced participation of adsorbed nucleophilic oxygen species in the conversion of n-butane leading to an increase in total oxidation as was shown by TAP experiments [19]. The second effect of w a t e r vapour is explained by a structural change of the catalyst surface (,,structural effect"). To prove if an enrichment in phosphorus is the reason for this effect, different catalyst samples were exposed to wet feeds containing different levels of water vapour and, thereafter, the chemical composition of the catalyst surfaces was determined. Figure 2 shows the change in P:V ratio as determined by XPS versus the duration of water vapour t r e a t m e n t . 1.55

1.6 t

~9 1.45

~9 1.5 -

1.35

~1.4 -

9

1.25

I

0

120

I

!

240 360 tH20, min

I

480

1.3 600

Figure 2. Dependence of the P/V ratio (XPS) on the duration of the water vapour t r e a t m e n t feeds: 1.46% C4H10- air +H20 480~ 0% H20(~); 480~ 5%H20('); 452~ 30%H20(A); 480~ 20%H20(m); 480~ 20%H20(O); 480~ 20%H20(X)

--D

I 0

I

1000 2000 'tN2, min

3000

Figure 3. Dependence of the P/V ratio (XPS) on the duration of the t r e a t m e n t with dry nitrogen T: 450~ p r e t r e a t m e n t : 8h 30%H20 in N2, 120~

The figure clearly reveals t h a t a prolonged exposure to w a t e r vapour resulted in an increase in the P:V ratio of the surface layers of the different samples. The presence of a small w a t e r vapour concentration of 5 vol% formed by the oxidation of b u t a n e in the conditioning led to an enhanced P:V ratio, too, in accordance with the finding in a previous study [5]. It is a s s u m e d t h a t the e n r i c h m e n t in phosphorus is driven by the gain in energy of the system due to the h y d r a t a t i o n of acidic phosphate groups. If this a s s u m p t i o n is true the e n r i c h m e n t m u s t be reversible . In the t r a n s i e n t e x p e r i m e n t mentioned above no reversibility of the a s s u m e d structural effect was observed certainly due to the w a t e r vapour formed by the oxidation of the originally dry feed. Therefore, the following experiment was performed: a (VO)2PzO7 catalyst with a surface enriched in phosphorus by a suitable p r e t r e a t m e n t with w a t e r vapour was exposed to a s t r e a m of dry nitrogen at elevated t e m p e r a t u r e for two days. During this t r e a t m e n t the P:V ratios in different samples t a k e n from the catalyst bed were determined. The results obtained are depicted in Figure 3. Indeed, a decline in the P/V ratio was found

466 depending on the time of treatment with N2. This result supports the above s t a t e m e n t concerning the reversibility of the P enrichment. For further confirmation, a transient experiment similar to t h a t in Figure 1 was carried out with the difference that, additionally, after 320 and 570 min time on stream, the butane-air feed was replaced by a stream of dry nitrogen for 20 hours. As shown in Figure 4, this procedure restored the original levels of activity and MA selectivity.

80 o ~60-

9 ,-4 CD

.

>

~.~ o ~40C9

~9

~ ~20-

-

N2 (20h)

+ 20% H20

0

I

I

120

240

I

N2 (20h) I

360 480 t, min

I

I

600

720

Figure 4. Dependence of the conversion of n-butane (I I) and of the MA selectivity (A) on time on stream Feed: 1.46 % C4H10- air, T= 450 ~ ~ = 1.6 sec; after 320 min and 570 min t r e a t m e n t with butane containing feed: 20 h Nz (GHSV: 67 h a (NPT), T= 450~

3.2. T h e m o b i l i t y o f p h o s p h a t e a n i o n s in (VO)2P207 c a t a l y s t s In order to get more information on the factors influencing the mobility of phosphate anions and causing the loss of phosphorus in (VO)2P207 catalysts in situ-thermogravimetric experiments were performed. Firstly, the adsorption of water molecules on the surface of CVO)2P207 was studied. Selected results are depicted in the Figures 5 and 6. Figure 5 shows the water uptake of the sample versus t e m p e r a t u r e on changing stepwise the composition of the gas phase from dry to wet nitrogen and vice versa (see also Figure 7). With increasing temperature the amount of adsorbed water molecules decreased up to 300~ and then remained nearly constant at higher temperatures. Figure 6 shows the dependence of P:V ratios determined by ISS (uppermost layer) and by XPS (about 50% of the XPS signal intensity stems from first 6 layers) on the temperature of the water vapour treatment.

467

~0.6

~4.0

%.

55 55

~.3o

~0.4 r

=0.2 ~: 0

t

t

200

400

C~l. 0 600

T,~ Figure 5. Dependence of the mass of H2Oads on the temperature, CH2O: 20% in N2, m e a t . : 120mg

I

100

I

300 T,~

,

I

500

Figure 6. Dependence of the P/V ratios (XPS (A), ISS (E~i))on the temperature of the water vapour treatment, CH2O: 30% in N2

Both the ISS- and the XPS-P:V ratios decrease with increasing temperature, i.e., with decreasing amounts of adsorbed water molecules. This dependence illustrates again the driving force of the P enrichment, i.e., the gain in energy by hydratation of P-OH groups. After conventional butane oxidation (450~ dry feed) the ISS-P:V ratio of the surface layer in used catalysts exceeds 2.0 indicating both an excess of phosphorus and a marked presence of vanadium in the utmost layer. This is in contrast to the pyrophosphate termination model of Thompson and Ebner [6] from which primarily the presence of phosphorus on the surface follows. P:V ratios ranging from 1.3 up to 1.4 were measured by XPS in samples which were obtained after conventional butane oxidation. These ratios are higher than those of O k u h a r a et al. [21] (1.10 _+ 0.04) and Coulston et al. [3] (1.08) who also used reliable sensitivity factors of XPS obtained by calibration. The reason for this difference might be that both the precursor synthesis and the catalyst preparation of the authors differ from the methods applied in the present work. From the XPS and ISS results obtained the following conclusion may be drawn. If the P:V ratio in the surface layer is about 2:1 then the P:V ratio in the next subsurface layers should be of similar magnitude. This follows from an electron attenuation length in the sample 00 of about 8 layers roughly estimated according to [20] and the lattice data of (VO)2P207 and VO(PO3)2 [22,23]. This conclusion is s u p p o r t e d by the s t a t e m e n t of C o u l s t o n et al [3] t h a t at m o s t 25% of the X P S s i g n a l is expected to originate from the o u t e r 5A of the (VO)2P207 s a m p l e . As Shown in Figure 7, a total weight loss of 0.11% occurs during the alternate t r e a t m e n t with wet and dry nitrogen for 6h at 480~ This is confirmed by other adsorption studies in the temperature range from 100~ to 600~ which reveal weight losses in the same order of magnitude. In contrast, no weight loss was observed if dry nitrogen was dosed over a dry calcined fresh catalyst.

468 122.3 ~ .

wet

wet Figure 7. Dependence of the mass of the (VO)2P207 catalyst on the duration of the alternate t r e a t m e n t with wet (20% H20) and dry N2, T: 480~

~122 1 drY-drY" 121.9

I 0

~y I

120

y I

240 360 t, min

I 480

In order to elucidate the reasons of this effect the weight of (VO)2P207 samples was m e a s u r e d during the treatment with nitrogen streams containing different levels of water vapour. Typical results are presented in Figures 8 and 9. 0.4 0.3 0.2 0.1

0.5 ~Y h~

0

-0.5

o o0 oQ

-0.1 -O.2 -0.3

I 0

120

-1.5

-2

I '

60

-1

180

t, min Figure 8. Influence of water vapour on the mass of a (VO)2P207 catalyst CH2Oin N2 : 0 % (O) ,10 % (?~),20 % (A), 30 % (X), T = 400 ~

0

60 120 t , min

180

Figure 9. Influence of water vapour on the mass of a (VO)2P207 catalyst at different t e m p e r a t u r e s C H 2 0 : 30 % in N2, T = 300 ~ (':)), 400 oc (D), 550 oc (A)

Figure 8 shows that in the first 30 min the mass of (VO)2P207 rose with increasing concentration of water vapour in the feed due to the adsorption of H20. During treatment with 20% and 30% H20 in N2 a subsequent weight loss was observed which was more pronounced with the higher water content. Figure 9 exhibits that the weight loss strongly depends on the t e m p e r a t u r e s of the water vapour treatment. The weight of the sample remained constant at 300~ whereas a strong mass loss was observed at 550~ The chemical analysis showed the presence of phosphoric acid and, surprisingly, also traces of vanadium containing species in the condensate. This finding is explained by decomposition of vanadyloligophosphate structures on the catalyst surface which are assumed to be formed under the influence of adsorbed w a t e r molecules. It is assumed that this also occur during the technical process but on a much longer time scale. However, after some hundreds of hours time on s t r e a m of an industrial reactor the loss in phosphorus leads to a decline in

469 selectivity and to an enhancement in activity. The addition of volatile phosphorus containing compounds to the feed is required to retain the catalytic performance. 4. C O N C L U S I O N S As shown by the results presented the surface layer of an equilibrated (VO)2P20: catalyst consists of a VPO structure with a P:V ratio of about 2 and in the next subsurface layers also higher P:V ratios than one have to be assumed. To elucidate the structure of the active site of the (VO)2P207 catalyst first of all some characteristics of the known active site models are given: usually the active center is described in terms of the bulk structure sometimes terminated in pyrophosphate groups stressing the role of edge-shared VO6 double octahedra (e.g. [2,6,24]). However, the existence of these VO6 double octahedra on the surface plane parallel to (I00) has not been proven up to now for the equilibrated (VO)2P207 catalyst, the surface of which is enriched in phosphorus due to the formation of water in the butane oxidation. Furthermore, the dynamic character of the surface is not considered in the published models although Pepera et al. [25] have demonstrated a high mobility of surface oxygen and that only a few surface layers take part in the butane oxidation. Additionally, the special properties of the surface in contrast to the bulk of a solid catalyst must be considered. This surface must be seen as a two-dimensional phase with special s t r u c t u r a l and energetic properties (see e.g. [26]). Taking these ideas into account, in the following a model is proposed which is able to explain our findings as well as the results of other authors who investigated the catalytic properties of VO(PO3)2 [7,8]. The active site of an equilibrated (VO)2P207 catalyst is assumed to consist of a thin layer of si~gle VOw; octahedra linked to the (VO)2P207 bulk in a chain-like fashion. This chain s t r u c t u r e is thought to be similar to the one of VO(PO3)2 and the VO6 octahedra to be isolated from each other by oligo- and monophosphate anions resulting in a P:V ratio of the uppermost layer of about 2. This structure should comprise besides the surface layer only few subsurface layers. The valence state of v a n a d i u m in this layer is between +4 and +5 according to previous results [5] and both phosphate and oxygen ions are very mobile (see [25]). This model explains additionally the result t h a t the area-specific rates of the MA formation are similar on (VO)2P207 and VO(PO3)2 as shown by Morishige et al. [7], S a n a n e s et al. [8] and by us [9]. The remarkable resistance of the bulk to oxidation under the usual reaction conditions as shown by earlier TG m e a s u r e m e n t s [5] reveals the active surface to be confined to a few layers on an oxidation resistent (VO)2P207 core. Furthermore, applying this model it is also possible to explain the enhancement in MA selectivity of a fresh catalyst p r e p a r e d from aqueous medium in the conditioning procedure [5]: During this process the original VO6 double octahedra of the surface region are splitted by migratory acidic phosphate anions to corner-shared single octahedra chains linked to the bulk double octahedra chains. This leads to a structure with isolated vanadyl chains which act more selectively according to the well known

470 ,,principle-of-the-density-of-the-oxidizing-sites" [27]. The same result is obtained when the precursor VOHPO~-0.5H~O is calcined in the presence of water vapour which results in a selective catalyst without the need of any conditioning procedure [28]. REFERENCES

1. 2. 3. 4. 5.

G. Centi, Catal. Today, 16 (1993) 1 F. Cavani, F. TrifirS, Catalysis, 11 (1994) 246 G.W. Coulston, E.A. Thompson, N. Herron, J. Catal., 163 (1996) 122 D. Ye, A. Satsuma, T. Hattori, J. Murakami, Appl. Catal., 69 (1991) L1 B. Kubias, M. Meisel, G.-U. Wolf, U. Rodemerck, Stud. Surf. Sci. Catal., 82 (1994) 195 6. M. R. Thompson, J. R. Ebner, Stud. Surf. Sci. Catal., 72 (1992) 353 7. H. Morishige, J. Tamaki, N. Miura, N. Yamazoe, Chem. Lett., (1990) 1513 8. M. Sananes, G.J. Hutchings, J.C. Volta, J. Catal., 154 (1995) 253 9. F. Hannour, A. Martin, B. Kubias, B. Liicke, E. Bordes, P. Courtine, Catal. Today, submitted 10. E. W. Arnold, S. Sundaresan, Appl. Catal., 41 (1988) 225 11.J.R. Ebner, M.R. Thompson, Catal. Today, 16 (1993) 51 12.H. Papp, F. Richter, G.-U. Wolf, Th. G5tze, Surf. Interface Anal., in preparation 13.F.R. Landsberger, P.J. Bray, J. Chem. Phys., 53 (1970) 2757 14.T. Okuhara, T. Nakama, M. Misono, Chem. Lett.,(1990) 1941 15.D. Briggs, M.P. Seah, Practical Surface Analysis, John Wiley and Sons, Chichester 1983, 512 16.T.L. Barr, J. Phys. Chem., 82 (1978) 1801 17.J.Ph. Nogier, M. Delmar, Catal. Today, 20 (1994) 109 18.F. Garbassi, J.C.J. Bart, R. Tassinari, G. Vlaic, P. Lagarde, J. Catal. 98 (1986) 271 19.U. Rodemerck, B. Kubias, H.-W. Zanthoff, M. Baerns, Appl. Catal., in press 20.R.M. Friedman, Silicates Ind., 39 (1973) 247 21.T. Okuhara, M. Misono, Catal. Today, 16 (1993) 61 22.S. Linde, L. Gorbunova, Dokl. Akad. Nauk, SSSR, 245 (1979) 584 23.V.V. Krasnikov, S.A. Konstant, Inorganic Mat. (Russ.) 15 (1979) 2164 24.J.Ziolkowski, E. Bordes, P. Courtine, J. Mol. Catal., 84 (1993) 307 25.M.A. Pepera, J.C. Callahan, M.S. Desmond, E.C. Milberger, P.R. Blum, N.J. Bremer, J. Am. Chem. Soc., 107 (1985) 4883 26. J. Haber, 8th Internat. Congr. on Catal., Berlin, 2.-6.7.84, Proc. Vol. 1, 85 27.J.L.Callahan, R.K.Grasselli, AIChE J., 9 (1963) 755 28.B. Kubias, unpublished result This work was supported by the Deutsche Forschungsgemeinschaft (project No Ku 957/1-4) and by the Bundesministerium fiir Bildung, Wissenschaft, Forschung und Technologie of the FRG (project No 03D0024).

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

471

Partial oxidation of C5 hydrocarbons to phthalic and maleic anhydrides over suboxides of vanadia" Use of dicyclopentadiene as a probe molecule U.S. Ozkan*, G. Karakas, B.T. Schilf, and S. Ang Department of Chemical Engineering, The Ohio State University, Columbus, Ohio 43210, U.S.A.

Partially reduced vanadium oxide catalysts have been examined in the selective oxidation of pentane and pentene to phthalic and maleic anhydride. The anhydride selectivity has been shown to be a strong function of catalyst pre-reduction and reaction temperatures. Controlled-atmosphere, postreduction and post-reaction surface characterization experiments have shown the most selective catalyst surface to_be comprised of V6013, V409, and VO 2 species. In this phase of the research, dicyclopentadiene has been used as a probe molecule to elucidate the reaction network for the fQrmation of phthalic and maleic anhydrides. 1. INTRODUCTION There is growing interest in the partial oxidation of the C 5 fraction of the hydrocarbon stream from naphtha steam crackers since there is no real market for them at the present time. Furthermore, the partial oxidation of lower alkanes and alkenes continues to pose challenging problems for catalysis researchers. In the case of C 5 hydrocarbon oxidation to form phthalic anhydride, the challenge is even greater since the catalyst needs not only to insert oxygen selectively, but also promote the formation of C-C bonds in an oxidative medium. In recent years, several studies have been reported in the literature, focusing on C 5 oxidation using catalysts such as supported vanadia, VPO catalysts, and molybdates [1-11] . In this study, suboxides of vanadia catalysts were used in pentane, pentene, dicyclopentadiene and cyclopentane oxidation reactions. In the previous phase of the work [12,13], the role of alkali promoters on the catalyst selectivity was examined. The catalysts were reduced in situ at different temperatures and the effect of pre-reduction temperature was investigated. C~176176 characterization of pre-reduction, post-reduction, and post-reaction catalysts were performed using X-ray diffraction, X-ray photoelectron spectroscopy, laser Raman spectroscopy and temperatureprogrammed desorption experiments. The objectives of this study were to determine the activity and selectivity of different suboxides of vanadia in

472 maleic and phthalic anhydride formation reactions and to examine the effect of alkali promoters on the product distribution. In this article, we report the results from the transient response and temperature programmed desorption experiments performed using dicyclopentadiene and phthalic anhydride as probe molecules. 2.

EXPERIMENTAL

Crystalline V20 5 was prepared by the decomposition of a m m o n i u m metavanadate (NH4VO 3 from Aldrich) with oxygen using a t e m p e r a t u r e programmed heat treatment. V20 5 was promoted with alkali promoters in the form of carbonates. The suboxides were prepared by reducing the catalysts with hydrogen in a temperature range of 300-600 ~ The catalysts were characterized by krypton BET surface area measurement, X-ray diffraction, controlled-atmospher e Raman spectroscopy, controlled-atmosphere X-ray photoelectron spectroscopy, temperature programmed reduction (TPR), and temperature programmed desorption (TPD)using n-pentane, 1-pentene, and dicyclopentadiene as adsorbates. The steady-state partial oxidation reaction experiments were carried out in a quartz fixed-bed, flow reactor with 10 mm I.D. and a length of 50 mm. The analytical system consisted of two parts, Gas Chromatography (GC) and High Performance Liquid Chromatography (HPLC). GC was used to analyze on-line the reaction products that were in the gas phase. The GC (HP-5890)was equipped with a Chromosorb PAW 23% SP-1700 column (30 i~. 1/8") for FID, a Hayesep D column (15 ft. 1/8") and a Molecular Sieve 5A column (10 i~. 1/8") connected serially to TCD. Reverse-phase HPLC with a Spherisorb 5 ODS-2 (1 ft. 1/4")column was used off-line to analyze the products that were solid or liquid at room temperature. The details of the experimental procedures have been described elsewhere [12,13]. The transient response experiments outlined in this article were performed pulsing feed mixtures of dicyclopentadiene, nitrogen, and oxygen through the catalyst bed which was connected on-line to the mass spectrometer (M.S. Engine, Hewlett Packard). The feed mixture had a composition of 0.036% dicyclopentadiene (H.P. Gas Products Inc.), 0.280% oxygen, 71.650% nitrogen and balance helium. In oxygen-free runs, oxygen was replaced by helium. The pulse size was 1 cm 3. The major ions monitored through the selected ion mode of mass spectrometer operation were 44 for CO2, 32 for 02, 18 for H20 , 16 and 29 for CH4, 54 and 98 for maleic anhydride, 104 for phthalic anhydride and phthalic acid, 105 for phthalic anhydride, 66 for dicyclopentadiene, 68 for methyl-maleic anhydride, 34 for pentadiene, 46, 88, and 116 for maleic acid. Similar experiments were performed by pulsing mixtures of phthalic anhydride through the catalyst bed. In the TPD experiments, the catalyst was reduced in situ and was degassed in vacuum. The adsorption of dicyclopentadiene was done at room temperature for 2 hours. After 30 minutes of flushing the system, the temperature program was started with a heating rate of 15~ The effluents were monitored by the quadrupole mass spectrometer using the s a m e m/e ratios listed above. In the TPD experiments, m/e=28 was used to monitor CO since there was no nitrogen present in the system.

473 3.

RESULTS AND DISCUSSION

In our previous studies [13], we have found the product distribution and the anhydride yields to be a strong function of reduction t e m p e r a t u r e in 1-pentene oxidation reaction. Table 1 s u m m a r i z e s the yields of major reaction products over the pre-reduced vanadia catalysts at different reduction and reaction temperatures. Only the results from the unpromoted catalysts are included in this table. The yield, YA for any product A is defined as YA =100 X (moles of A in the product stream x 7A)/(moles of reactant in the feed) where 7A is the ratio of the number of C atoms in product A to the n u m b e r of C atoms in the reactant molecule. Table 1 Yield of major products in 1-pentene oxidation Reduction temp. (~ Reaction temp. (~ M.A.

P.A. CO CO2 Me-M.A.

450

400

600

350

375

400

350

375

400

350

375

400

0.2 0.3 2.5 3.5 0.3

0.8 0.5 4.5 8.6 3.0

5.7 7.7 34.2 37.1 7.3

4.3 0.5 6.6 9.2 10.6

6.3 0.9 12.5 14.1 12.8

5.8 1.1 16.6 19.7 10.2

0.3 0.1 2.9 3.6 1.3

0.2 0.1 2.9 4.8 0.8

0.3 0.1 4.1 6.1 0.8

M.A.- maleic anhydride, P.A.- phthalic anhydride, Me-M.A.- methyl maleic anhydride As seen in this table, the highest anhydride yield is obtained with a prereduction t e m p e r a t u r e of 400 ~ and at a reaction t e m p e r a t u r e of 400 ~ The other reaction products are acetone, acetaldehyde and formaldehyde. W h e n the catalysts pre-reduced at 400 ~ were characterized by controlledatmosphere X-ray photoelectron spectroscopy, the surface was found to be comprised of three different suboxides. The surface composition was found to be 64% V409, 5% V6013, and 31% VO 2. While our studies for identification of the active/selective phase continues, in a parallel effort, we have conducted some experiments to assess the role of dicyclopentadiene, one of the intermediates proposed in the literature [7]. The objective behind this phase of the study was to examine the routes that originate from dicyclopentadiene and result in phthalic anhydride v e r s u s maleic anhydride formation. The catalyst used in this set of experiments w a s v a n a d i u m pentoxide t h a t was pre-reduced in situ at 400 ~ with hydrogen.

474 Figure 1 shows signals when the feed mixture consisting of dicyclopentadiene, oxygen, and nitrogen was pulsed through the catalyst bed at room temperature. What is worth mentioning about this figure is the dicyclopentadiene signal which shows an initial spike, that coincides with the nitrogen signal and a second broader feature, which exhibits a tail. This provides a clear indication of the adsorption/desorption interaction between the catalyst surface and the dicyclopentadiene molecule. The time scale between the on-set of the signal and the end point of the tail was about 3.9 minutes.

m/e=66 o

i~ 3.9 min.

"o t..Q

~1

<

m/e=32

.

.

.

.

IL

_

_

,

_

m/e=28

Time

Figure 1.

Variation of reactant signals in dicyclopentadiene-nitrogen-oxygen pulses at room temperature.

475 Figure 2 shows the results obtained when the same feed mixture was pulsed through the catalyst bed at a reaction temperature of 250~ At this temperature, the dicyclopentadiene was almost completely converted. The major products obtained from these pulses were maleic anhydride, methyl maleic anhydride, and CO 2. The formation of CO could not be detected since the major ion m/e=28 for CO overlapped with that of nitrogen. There was no phthalic anhydride observed in these experiments. An interesting feature of the maleic anhydride and methyl maleic anhydride signals is that, at each pulse, there are two maxima. The first maximum in methyl maleic anhydride signal coincides with those of oxygen and nitrogen. The second m a x i m u m , which is much stronger in intensity than the first one occurs 63 seconds later. The maleic anhydride, on the other hand, exhibits two m axi m a 26 seconds apart and the second m a x i m u m coincides with the second m a x i m u m of the methyl maleic anhydride signal. This observation suggests that there may be two distinct routes leading to the formation of maleic and methyl maleic anhydride. The continuous tailing feature of the CO2, as opposed to the doublet feature, on the other hand, suggests several multiple steps contributing to the C O2 signal.

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..Q

m/e=54

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m/e=66

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Time Figure 2.

Variation of product signals in dicyclopentadiene-nitrogen-oxygen pulses at 250 ~

476 When similar pulse experiments were repeated using an oxygen-free feed mixture, the results were significantly different from those obtained in the pulses with oxygen present. The results are presented in Figure 3. The experiment was continued for 20 pulses although only the first eleven are presented in the figure. This time, dicyclopentadiene signal was seen to grow with increasing pulse number and reach a steady value after 12 pulses. The major products were again maleic anhydride, methyl maleic anhydride and CO 2 . However, what was different about these pulses was the fact that phthalic anhydride formation was also observed as shown in the figure. This was verified by both m/e=104 and m/e=105 signals. There were also m/e=88 and 116 signals which are likely to be due to maleic acid. The m/e=28 signal showed an increase with the pulse number, which is likely to be due to CO signal being superimposed on the nitrogen signal. In these pulse experiments, it was demonstrated that dicyclopentadiene was highly reactive even at temperatures as low as 250 ~ and some phthalic anhydride was formed over the catalyst if there was no oxygen in the gas feed. However, maleic anhydride and methyl maleic anhydride were produced in larger quantities in the presence as well as in the absence of oxygen, suggesting that there may be alternative pathways that lead to maleic anhydride formation. One of the possibilities is C-C bond cleavage of dicyclopentadiene before oxygen insertion step(s), which lead to maleic and/or methyl maleic anhydride. The other possibility is the ring opening and C-C bond cleavage steps of phthalic anhydride, which again may lead to methyl maleic anhydride and, upon removal of the methyl group, to maleic anhydride. In order to examine the reactivity of phthalic anhydride over these catalysts, pulse experiments were performed using a feed mixture of phthalic anhydride, oxygen, and nitrogen. Experiments were repeated after removing the oxygen from the mixture. The phthalic anhydride concentration was approximately 0.4%, estimated from vapor pressure correlations. The pulses sent through the catalyst bed at 200 ~ resulted in essentially complete conversion of phthalic anhydride. There were trace quantities of maleic anhydride (m/e=98), substantial amounts of maleic acid (m/e=ll6), CO 2 (m/e=44), and water (m/e=18). There was also a strong signal at m/e=70, which can be due to a C~ hydrocarbon. These experiments showed phthalic anhydride to be very reactive, even at temperatures as low as 200 ~ They also pointed out the possibility of the formation of maleic anhydride (or maleic acid, which can be formed from the anhydride readily in the presence of water) from phthalic anhydride. When the results obtained from the oxygen-containing pulses are reconsidered in the light of this new observation, it appears that the first and the second maxima seen in the maleic anhydride and methyl maleic anhydride signals could be due to the formation of the two anhydrides from a dicyclopentadiene derivatives and from ring opening the dealkylation of phthalic anhydride, respectively. Temperature programmed desorption profiles obtained after roomtemperature adsorption of dicyclopentadiene over the catalyst which was prereduced at 400 ~ are shown in Figure 4. These profiles showed two temperature maxima. The first one took place at 460 ~ and gave signals for the m/e=28 and 29, which are likely to be due to carbon monoxide and m e t h a n e , especially since there was a peak at the same temperature at m/e=16.

477

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m/e 54

< "~

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m

10 min.

m/e=66

~

Time

Figure 3. Variation of product signals in dicyclopentadiene-nitrogen pulses at 250 ~

478 The other ions which showed very weak signals at this temperature were m/e=66, 68, and 98, suggesting the presence of dicyclopentadiene, methyl maleic anhydride, and maleic anhydride. The second temperature m a x i m a were seen after the catalyst temperature was raised to 600 ~ The ions detected at this temperature were m/e=34, 44, 46, signaling the presence of pentadiene, CO 2 , and possibly maleic acid. The fact that dicyclopentadiene signal was almost negligible implies a strong surface-adsorbate interaction. Also, the peaks that emerged at 600 ~ show that the adsorbed species desorb only as complete oxidation or cracking products.

m/e=68&98

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30

40

50

60

100

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Figure 4. Dicyclopentadiene temperature-programmed desorption profiles.

479 4. CONCLUSIONS Dicyclopentadiene, an intermediate proposed in the C~ partial oxidation literature [7], was used as a probe molecule in this phase of the study. In our previous studies, we have reported the strong dependence of anhydride selectivity on pre-reduction temperature of V20 S catalyst. Although an optimization procedure was not followed to maximize the yield, a pre-reduction temperature of 400 ~ coupled with a reaction temperature of 400 ~ gave the best results with a combined anhydride yield (maleic+phthalic) of about 22 % under the conditions used in these experiments. When we examined the dicyclopentadiene reactivity using the same pre-reduction temperature, it was shown to be adsorbing strongly and reacting readily to produce phthalic anhydride, maleic and methyl maleic anhydride, and carbon oxides. Experiments conducted using phthalic anhydride as a feed molecule showed that phthalic anhydride was not very stable, yielding maleic acid or anhydride as a result of C-C bond breakage, leading to complete oxidation to CO and CO 2. The observations made in this phase of our study lead us to suggest that there are multiple routes that go through the dicylcopentadiene intermediate. One of the routes involves a pre-oxidation ring opening and dealkylation of the molecule followed by oxygen insertion steps. This route can give C 4 or C 5 anhydrides (or acids in the presence of water). The second route involves direct oxygen insertion into the dicyclopentadiene molecule, leading to phthalic anhydride formation. It is more likely, however, that phthalic anhydride itself can undergo partial ring opening and dealkylation steps, resulting in the maleic and methyl-maleic anhydrides, and, with further oxidation, in carbon oxides. ACKNOWLEDGEMENTS Financial support provided for this work by the National Foundation through Grant CTS-9412544 is gratefully acknowledged.

Science

REFEttENCES ,

2. 3.

.

,

N.S. Butt and A. Fish, J. Catal., 5 (1966) 205. G. Centi, J.L. Nieto, and C. Iapalucci, App. Catal., 46 (1989) 197. K.E. Birkeland, W.D. Harding, L. Owend, and H.H. Kung, in S.A. Bradley and J. Stencel (Editors) Chemistry and Characterization of Supported Metal Catalysts, Division of Petroleum Chemistry, Inc., American Chemical Society. 38, (1993), 880 C. Fumagalli, G. Golinelli, G. Mazzoni, M. Messori, G. Stefani, and F. Trifir6, in V. C. Corber~n and S.V. Bell6n (Eds.) New Developments in Selective Oxidation, Elsevier Science Publishers B.V., Amsterdam, 1994, p.221. G. Centi, J.T. Gleaves, G. Golinelli, and F. Trifir6, in B. Delmon and P. Ruiz, (Eds.) New Developments in Selective Oxidation, Elsevier Science Publishers B.V., Amsterdam, 1992, p. 231.

480 ,

7. 8. 9. 10. 11. 12. 13.

F. Trifir6, Catal. Today, 16 (1993) 91. G. Centi and F. Trifir6, Chem. Engr. Sci., 45(8) (1990) 2589. G. Busca and G. Centi, J.Am.Chem.Soc., 111 (1989) 46. [15] G. Centi, J.T. Gleaves, G. Golinelli, S. Perathoner, and F. Trifir6, in C.H. Bartholomew, and J.B. Butt (Eds.), Catalyst Deactivation, Elsevier Science Publishers B.V., Amsterdam, 1991, p. 449. G. Golinelli and J.T. Gleaves, J. Mol. Catal., 73 (1992) 353. U.S. Ozkan, R.E. Gooding and B.T. Schilf, in B.K. warren and S. T. Oyama (Eds) Heterogeneous Hydrocarbon Oxidation, ACS Symposium Series 638, p. 178. U.S. Ozkan, T. A. Harris, and B.T. Schilf, Catalysis Today, 33(1996) 57. B.T. Schilf and U.S. Ozkan, Applied Catalysis, to be submitted for publication.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

481

Role of h o m o g e n e o u s r e a c t i o n s i n t h e c o n t r o l of t h e s e l e c t i v i t y to m a l e i c a n d p h t h a l i c a n h y d r i d e s in t h e o x i d a t i o n of n - p e n t a n e Z. Sobalik 1 , P. Ruiz 2 and B. Delmon 2 1j. Heyrovsky Institute of Physical Chemistry, Academy of Sciences of the Czech Republic, Dolejskova 3, 182 23 Prague 8, Czech Republic 2 Unit~ de catalyse et chimie des mat~riaux divis~s, Universit~ catholique de Louvain, Place Croix du Sud, 2 bte 17, B-1348, Louvain-la-Neuve, Belgium. Oxidation of n-pentane on vanadium phosphate catalysts was followed in a mixed heterogeneous-homogeneous regime and the effect of the modification of both reaction regimes on the ratio of phthalic (PA) to maleie anhydride (MA) was evaluated. The extent of the oxidation under homogeneous conditions was controlled either by varying the volume of the void section of the reactor or by adding methyl-ter-butylether (MTBE). Results show that the homogeneous reaction represses preferentially PA formation, i n c r e a s i n g the MA/PA ratio. MTBE inhibits both the homogeneous and the heterogeneous reaction. Such behaviour is reversible, however recovery of the PA formation was substantially slower after stopping MTBE addition. It could be concluded that, in the study of the activation of n-pentane (and alkanes), the interaction of both (homogeneous and heterogeneous) processes must be taken into account. 1. INTRODUCTION The present work was undertaken in the context of the present attempts to use saturated alkanes to produce valuable functionalized hydrocarbons. This general approach is presently one of the major challenges for the petrochemical industries in search for better utilisation of light alkanes. Catalytic oxidation seems to be the most promising route. But as has been pointed out recently [1], a much broader approach has to be taken compared to that used in former investigations for achieving high selectivity. Among others the interrelation between the homogeneous and heterogeneous processes which may occur simultaneously in light hydrocarbons oxidation has to be taken into account. Its role in the selectivity control should be investigated carefully, this particularly because it appears that the gas phase reaction may change selectivity, possibly in a favourable way in some cases. The idea to use a combination of homogeneous and heterogeneous reactions to change the selectivity of oxidative dehydrogenation of propane

482 was described by Burch et al. [2]. They found that the radical reactions may influence the product distribution of the catalysed reaction. We illustrate the importance of such an approach in the case of npentane oxidation over a vanadium phosphate catalyst (VPO). Both phthalic and maleic anhydrides can be obtained in this reaction. The formation of PA is highly specific for VPO catalysts. The corresponding reaction is complex. It involves the abstraction of 10 H-atoms, and several oxygens should be provided for full oxidation of two carbon atoms to COs, insertion of 3 oxygen atoms and condensation-cyclisation of two molecules of n-pentane. As has been already underlined, the control over the PA/MA ratio has a high practical impact if the functionalisation of n-pentane through selective oxidation is contemplated [3]. With n-pentane the homogeneous reaction can occur in conditions identical or very close to those of the heterogeneous oxidation. Two methods of controlling the homogeneous reaction of the mixed homogeneous- heterogeneous process were used. In the first method, we varied the void volume of the reactor. In the second, we blocked the propagation of the oxidation reaction in the homogeneous phase. It is well known t h a t there are promoters (e.g. N02) and inhibitors (e.g. tetraethyl lead) for the combustion [4]. In particular, MTBE used as a substitute of tetraethyl lead should block the spontaneous ignition in the fuel/air mixture in the engines. Accordingly, we have speculated that MTBE could influence the homogeneous oxidation under our experimental conditions. Our strategy was to use various combinations of homogeneous and heterogeneous regimes and to control the extent of the homogeneous reaction by these two methods. The speculation was based on the fact that, in some mechanism models of n - p e n t a n e oxidation (hydrogen a b s t r a c t i o n ) , oxidative dehydrogenation is also the starting point. 2. EXPERIMENTAL The VPO catalyst was prepared by a method described previously [5]. The bulk molar P/V ratio of the catalyst was 1.26 and the surface area was 44 m2/g. The catalyst was equilibrated for over 80 hours at the n-pentane/air reaction mixture at the temperature of about 390~ before measurements. Silicium carbide (Prolabo, purity > 99 %; grain fraction 0.35 mm), n-pentane (Aldrich, purity > 99 %; impurity identified as 2-pentenes) and methyl-tertbutyl ether (MTBE) (Aldrich, purity > 98 %) were used without further purification. Oxidation of n-pentane tests were made in a U-tube glass reactor with a glass frit layer (thickness of about 2 mm) and the inner diameter of the tube 4 mm and 9.3 mm before and aider the frit layer, respectively. The length of the void section of the reactor was controlled by inserting a plug of quartz wool in the tube upstream of the frit layer. In the catalytic experiments about 0.2 g of the VPO catalyst (0.5-0.8 mm particles) was placed on the frit layer. The rest of the reactor was filled with silicium carbide. The reactor was inserted into an electrically heated furnace providing for a flat temperature profile (e.g. at 400 ~ the temperature along the 6 cm length of the central zone of the reactor varied for less than + 2.5 ~

483

jb

Figure 1. Scheme of the U-tube reactor for homogeneous and mixed regime oxidation. Parts filled by SiC(a); thermocouple (b); place for a catalyst layer (c); frit layer (d); void space(e); plug of a glass wool(f).

The gas composition (volume) was: n-pentane 0.6-1.3 %; oxygen 0-20 %; helium 5 %; balance nitrogen. Total flow 30 c m 3 . min-1 ( H o u r l y Space Velocity 6000 h-l). Brooks mass flow controllers were used to control air and helium stream. MTBE was introduced into the np e n t a n e / a i r m i x t u r e using a mechanical injector and a tube evaporator kept at about 150 ~ The delay for the step-wise change of the MTBE feed to reach the reactor zone was about 2 minutes. The outlet of the reactor prior to analysis was kept at between 140 and 150 ~ so the c o n d e n s a t i o n of the products at the concentration levels produced at the catalytic test was satisfactorily prevented.

The inlet as well as the outlet streams were analysed by an on-line gas chromatography, equipped with a TCD detector. Two columns, one with TENAX (Alltech Associates, Inc.) for determination of MA and PA and other oxygenates products, and a Porapak Q (Alltech Associates, Inc.) for npentane were used, both stainless steel of about 2.5 m length and 1/8" in diameter. Helium was used as a carrier gas at a flow of 22 ml/min in both columns. Two loops of 1.69 and 9.50 ml were used for injection on the Porapak Q and the TENAX columns, respectively. Injection loops as well as the 6-port valve were kept at a temperature of about 175~ A temperature program was used on both columns, i.e. 2 minutes at 70 ~ and then increase of the temperature at a rate of about 20 ~ up to 190 ~ for the column with Porapak Q and, for Tenax, 2 minutes at 120 ~ and then an increase up to 240 ~ Conversion of n-pentane is defined as the number of moles of n-pentane converted to the number of moles of n-pentane feed to the reactor (in %tool). Selectivities to maleic and phthalic anhydrides are expressed as a fraction of moles of n-pentane converted into MA and PA (in %mol), respectively: SMA = ((CMA)/( CC5))X 102, SpA = ((2 x CpA)/(CC5)) x 102. CC5 is the amount of n pentane transformed, i.e. moles of n-pentane in the inlet minus moles o~ n-pentane in the outlet (mole n-pentane . ml-1), CMAis the amount of MA produced (mole PA per ml of the outlet gas mixture), CpA is the amount of PA produced (mole PA per ml of the outlet gas mixture).

484

3. RESULTS 3.1. Oxidation of n-pentane u n d e r homogeneous conditions Two methods of control of the homogeneous reaction were evaluated: i) varying the residence time in the empty space of the reactor by changing the length of the void tube up stream of the catalyst; or ii) trapping free radicals by addition of methyl-terc-butyl ether (MTBE), and thus blocking the propagation of the oxidation reaction in the homogeneous phase. 3.1.1. Control by a void volume

Homogeneous oxidation of n-pentane presented as conversion vs. reaction t e m p e r a t u r e under various n-pentane/oxygen ratios displayed typical bell-shape curves over the temperature region of 300 - 500 ~ [3], as illustrated in Fig. 2. The n-pentane conversion is near to zero at 5 % of oxygen and increases with oxygen content. For a given n-pentane/oxygen ratio, the n-pentane conversion varies in a regular manner (see Fig. 3), increasing with temperature and residence time of the reaction mixture in the empty space of the reactor.

Cb

"6 2 E ~L

50

C

cb

b

C~

E :3

25

0

1

(D

0

~

200

Ci

0

250

300

350

400

450

500

Temperature ~

Figure 2. Homogeneous oxidation of n-pentane (1 vol.%) and oxygen/npentane ratio of 10 (a) or 20 (b); length of the void part of reactor 7 cm.

0

'

0

1

O

2 Residence time, s

Figure 3. Consumption of n-pentane by homogeneous oxidation at n-pentane (0.6 vol.%) / oxygen (19 vol%) mixture, by varying the residence time in the e m p t y space of the r e a c t o r for temperatures of 350 ~ (a), 370 ~ (b), and 390 ~ (c).

485

3.1.2. Control by MTBE addition Addition of MTBE substantially suppressed the n-pentane oxidation, with the onset of np e n t a n e conversion shifting by about 50 ~ to a higher temperature. The typical result of such experiment is depicted in Fig. 4. Accordingly by MTBE addition the homogeneous oxidation of npentane could be fully suppressed up to about 360 ~

(b .~

50

~ 40 c (b ~ 30 *~ 20 ,~ c /o

o i

250

300

350 400 Temperature, ~

Figure 4. Homogeneous oxidation of n-pentane (1.5 vol.%) by oxygen (20 vol.%) before (a) and after addition of 0.35 vol.% (b) or 0.6 vol.% (c) of MTBE; length of the void part of the reactor 7 cm.

3.2. Oxidation of n-pentane under mixed h o m o g e n e o u s - h e t e r o g e n e o u s conditions. 32,1. Control of the homogeneous oxidation by void volume. The homogeneous reaction occurred in the void section of the reactor and then the reaction mixture passed through the frit and the catalyst layer. R e s u l t i n g n - p e n t a n e conversion (CToT, %, void section + catalyst), homogeneous conversion (CHoM, %; void section without catalyst), and selectivities to MA and PA (void section + catalyst) are shown in Table I. Table I Influence of the homogeneous reaction on the oxidation of n-pentane over VPO catalyst. void CHOM~ CTOT o~ SMA • SpA % section(cm) 0 0 27 62 34 0.2 12 33 37 20 0.5 27 60 22 6.5 1.5 35 68 20.4 5.0 E x p e r i m e n t a l conditions" pentane 0.7 vol.%; oxygen 20%; reaction temperature 375 ~ By increasing the void volume of the reactor, the homogeneous reaction diminishes the selectivity to partially oxidised products and, at the same time, increases the n-pentane conversion and significantly modifies the maleic to phthalic anhydride selectivity (see also Fig. 5).

486

3.2.2. Control of the h o m o g e n e o u s r e a c t i o n by MTBE additiorL It has been shown in the previous part that addition of MTBE substantially suppressed the h o m o g e n e o u s oxidation. But (see Table II), co-feeding of MTBE to the npentane/oxygen reaction mixture also blocked the heterogeneous reaction.

3

2

o

o

lO

20

30

40

Chom, %

Figure 5. SMA/SpA vs. homogeneous conversion of n-pentane.

Table II Influence of MTBE addition on the oxidation of n-pentane c (MTB E, n-pentane conv., S(MA) S(P %vol) % A) 0 52 10.5 9.4 0.6 2 58 14 Experimental conditions: n-pentane 1.5 vol.%, oxygen 20 vol%.; reaction temperature 365 ~ 1.o ~

0--0-43

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b T

0.6

T112

c

0.4 0.2 0.0 9 i

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!

20

40

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60

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80 100 120 Time, minutes

Figure 6. Response curves of n-pentane oxidation (a) and formation of maleic (b) or phthalic (c) anhydrides after stopping MTBE addition. (Conversion of n-pentane, and the amount of MA or PA formed are normalised to the values before MTBE addition). Experimental conditions: n-pentane 1.5 vol.%, oxygen 20 vol.%, and 0.6 vol.% of MTBE; reaction temperature 365 ~ void section 7 can. ~1/2 is half time, namely the time necessary after stopping of MTBE to reach one half of its steady-state value before the MTBE was injected.

487 Such behaviour has been shown to be reversible and, after stopping the MTBE injection, the n-pentane conversion and selectivity to MA and PA return to their former values. Nevertheless, the analysis of the response curves indicates a substantial difference between the transient changes of MA or PA formation. Results characterising the transition behaviour of the system after starting and stopping the MTBE injection are depicted in Fig. 6. 4. DISCUSSION

The behavior of the interplay between homogeneous and heterogeneous processes of n-pentane activation and change of catalyst oxidation state during dynamic experiments is discussed. The results suggest that by taking advantage of the interaction between the homogeneous and heterogeneous oxidation processes, and controlling the conditions for activation of n-pentane in homogeneous phase, the MA/PA ratio could be varied substantially. By varying the void volume and n-pentane/oxygen ratio, the homogeneous reaction could be controlled conveniently. Under such conditions the non-selective homogeneous reaction increases the n-pentane conversion and influences also substantially the ratio of maleic/phthalic anhydrides selectivity. By varying the concentration of n-pentane in the feed, the observed relative decrease of the PA/MA formation cannot be explained by a decrease of the n-pentane content in the reaction mixture entering the catalyst layer following the void section. In fact, the increase of the MA/PA value caused by varying the n-pentane content in the feed under reaction conditions without the void section brings about less than one-third of the effect found in experiments where the same changes in the feed were produced by n-pentane consumption in the homogeneous oxidation. This means that the presence of some of the products of homogeneous oxidation suppressed preferentially the formation of PA. The explanation of such an observation is difficult. At this stage of our investigation, it can be suggested that the decrease of PA formation could be correlated with a change in the nature of the reaction mixtures, for example by the presence of some unsaturated species. Results presented in the literature showing a decrease of PA formation by the addition of olefins to n-pentane/oxygen mixtures [6] support this explanation. Such unsaturated species are usually observed among the products of oxidative dehydration of alkanes in the homogeneous phase [2]. Results of the MTBE co-feeding experiments show a new and complex phenomenon. Change in the homogeneous and in the heterogeneous reaction have been observed. Obviously addition of a substance with an established ability to suppress initiation of radical reaction protect n-pentane against an oxygen attack in homogeneous phase, but at the same time suppress n-pentane activation on the surface of VPO catalyst. Suppression of the heterogeneous reaction could be explained by the fact that MTBE can readily produce some isobutylene, which will very efficiently compete for npentane active surface sites. The competition will be much in favor of the olefin, thus bringing about the observed decrease in paraffin activation. However, it must be stressed that the subsequent reaction steps of the complex reaction mechanism leading to MA formation are not influenced,

488 as indicated by a high selectivity to MA under MTBE addition. On the other hand, the PA/MA ratio is substantially decreased, quite opposite to a general tendency of the n-pentane VPO system producing usually more PA at low npentane conversion levels [7]. The difference in the response of the MA and PA formation in the system due to MTBE addition is further stressed by the shapes of transition curves for n-pentane conversion and MA and PA formation after stopping MTBE injection. The recovery of n-pentane conversion is the same in both homogeneous and heterogeneous regimes and fully parallel with the MA formation, showing xl/2 of about 5-7 minutes. On the other hand the recovery of the PA formation is much slower (~1/2 - 50 rain). The explanation of such phenomena is also complex and needs more investigation in order to understand in more detail this phenomenon. But what is necessary to underline here is that the selectivity to MA can be modulated by adding small amounts of another component that inhibits the homogeneous and a part of the heterogeneous reaction. 5. CONCLUSION Using the example of n-pentane oxidation on VPO catalyst it has been shown t h a t homogeneous reaction can play an i m p o r t a n t role in the oxidation of n-pentane. The occurrence of the homogeneous reaction in parallel with the heterogeneously catalysed one might modify the selectivity. It could be speculated that by using such approach and controlling the conditions for activation of n-pentane in homogeneous phase, MA/PA ratio could be varied substantially. In this work, we have shown that homogeneous-heterogeneous reactions are interdependent in the oxidation of n-pentane on VPO. Results presented in the literature show that the activation of other alkanes involves similar phenomena. Thus, in order to progress in the knowledge of the activation of alkanes, the interaction between homogeneous and heterogeneous processes must be taken into account. These two processes depend on the nature, shape and size of the catalysts, and the form of the reactor. They bring about a change of the overall reaction kinetics. 6. ACKNOWI~DGEMENTS The Service de Programmation de la Politique Scientifique (Belgium) is gratefully acknowledged for its Concerted Action grant, especially for the support of Ing. Z. Sobalik and Dr. P. Ruiz. R~'ERENCES 1. B. Delmon P. Ruiz, S.R.G. Carrazan, S. Korili, M.A. Vicente Rodriquez, and Z. Sobalik, Catalysis in Petroleum Refining and Petrochemical Industries 1995, M. Absi-Halabi et al. (Editors), Elsevier,Amsterdam, 1996, p. 1. 2. R. Burch and E.M. Crabb, Appl. Catal., 100 (1993) 111.

489 3. G. Centi, J.T. Gleaves, G. Golinelli and F. Trifiro, Stud. Surf. Sci. Catal., 72 (1992) 231. 4. H.M. Spiers, F.I. Ceram (Eds.), Technical Data on Fuel, London 1962, p.263. 5. Z. Sobalik, S. Gonzales, P. Ruiz, Stud.Surf.Sci.Catal. 91 (1995) 727. 6. G. Centi, J.L. Nieto, D. Pinelli, F. Trifiro and F. Ungarelli, New Developments in Selective Oxidation (G. Centi, F. Trifiro, Eds.) Elsevier, Amsterdam 1990, p. 635. 7. J.T. Gleaves and G. Centi, Catal. Today, 16 (1993) 69.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

Catalytic Oxidation of Alkanes at Millisecond Contact T i m e s Lanny D. Schmidt and Christian T. Goralski, Jr. Department of Chemical Engineering and Materials Science, University of Minnesota, Minneapolis MN 55455 Catalytic oxidation reactions on noble metal surfaces are sufficiently fast and exothermic that they can be operated at contact times on the order of one millisecond with nearly adiabatic temperatures of--1000~ At short contact times and high temperatures complete reaction of the limiting feed is observed, and highly nonequilibrium products are obtained. We summarize experiments where these processes are used to produce syngas by partial oxidation of methane, olefins by partial oxidation of higher alkanes, and combustion products by total oxidation of alkanes. The former are used to produce chemicals, while the latter is used for high temperature catalytic incineration of volatile organic compounds. 1. CATALYTIC OXIDATION The subject of catalytic oxidation of alkanes goes back to the early 19th century when Sir Humphrey Davy and Michael Faraday found that a Pt wire would glow spontaneously in air when exposed to fuels such as H2, NH3, and natural gas. They correctly identified this as a surface catalytic reaction in which the heat of the exothermic oxidation reaction heated the catalyst to high temperatures where rates are extremely high. Wilhelm Ostwald exploited this process in the early 1900s for the oxidation of NH3 using wires woven into gauzes, and this process is still the basis of commercial oxidation of NH3 to NO in the reaction: NH3 + 5/402 ~ NO + 3/2H20 and then on to HNO3 in nitric acid synthesis. Modem reactors operate with 5 to 20 layers of Pt- 10%Rh woven gauze in layers 5 to 20 feet diameter with 10% NH 3 in air at pressures up to 10 arm, gauze temperatures of--800~ and gas velocities of --1 m/sec, for contact times of --1 msec. In the 1940s Leonid Andrussow added CH4 and removed some 02 to produce HCN in the overall reaction: CH4 + NH3 + 3/202 ---) HCN + 3H20. This process takes place over Pt-10% Rh gauze in a catalyst and conditions essentially identical to that used for HNO3 synthesis in an atmospheric pressure adiabatic reactor operating at --1100~ with contact times of ~1 msec. This process is the basis of HCN synthesis in Nylon and methylmethacrylate polymers. In spite of the commercial successes of these two millisecond reactors, few processes other than NH3 oxidation (a superoxidation) and HCN synthesis (oxidative dehydrogenation or ammoxidation) have been carried out on a large scale. In the automotive catalytic converter, contact times over noble metals on wash coated extruded ceramics are used with contact times of ~0.1 sec with temperatures of ~400oc. This research was supported by NSF, DOE, and the Gas Research Institute.

491

492 In catalytic incineration with conventional technology, noble metal particles on high area

T-A1203is used with contact times o f - 1 sec at temperatures of 200 to 600~ There has been considerable recent interest in partial oxidation of alkanes for new routes to chemical synthesis from light alkanes, and this research has been summarized extensively[ 1-3]. While much of this research uses dilution and low flow rates to attain temperatures between 200 and 600~ another mode of operation involves very short contact times with no dilution to increase the temperature to --1000~ and decrease the catalyst contact time to --1 millisecond. Recent research with millisecond adiabatic reactors have been carried out by Green et al[4], Lunsford et al[5], Choudhary et al[6], and many other research groups. We have also examined these processes extensively, and this paper will summarize this research to produce syngas[7-9], olefins[10], and oxygenates[11]. In excess 02 similar processes produce primarily CO2 and H20, and these processes are very important to reduce pollution from NOx, CO, and unburned hydrocarbons in combustion to produce heat, radiation, and for abatement of volatile organic compounds (VOCs) in air. Many chemical synthesis processes are carded out in excess fuel to promote formation of partial oxidation products, while combustion to CO2 and H20 is carded out in excess 02. These species are the dominant products in excess 02, and the high heat of combustion of alkanes creates adiabatic temperatures o f - 2 0 0 0 ~ for feed compositions near the stoichiometric composition for CO2 and H20 (t~=l). Beyond t~=l, CO forms instead of CO2 and, at even higher ~, H2 forms instead of H20 to produce syngas which can be used to produce methanol, synthetic diesel fuel, or H2. For even richer mixtures, oxidative dehydrogenation to produce olefins begins, and then oxidative addition to produce oxygen containing products[ 11]. Millisecond reactors generally operate with a low area porous monolith catalyst sealed in a tube as sketched in figure 1. Reacting gases near room temperature flow at high velocities through the monolith, and, after lightoff, almost complete reaction of the limiting reacting reactant can be obtained for ~=0.3 to >2.0 with total contact times on the catalyst less than 10 milliseconds. In results reported here we used a metal coated ceramic foam 45 ppi t~-A1203 monolith 10 mm thick and 18 mm in diameter sealed in a quartz tube reactor as sketched in figure 1. Radiation

Shields

/\ Mass Flow Controllers

/ Catalyst

Incinerator

Gas Chromatograph

Figure 1. Sketch of reactor configuration used for catalytic oxidation on monolith reactors at millisecond contact times. Gases slightly above atmospheric pressure flow at high velocities through porous ceramic monoliths coated with Rh or Pt. All metals were deposited from metal salt solutions (H2PtC16, Rh(NO3)3, and SnC12), followed by calcination and reduction as described elsewhere [7-10]. The total area of these

493 catalysts was--0.1 m2/gram. The catalyst was wrapped in alumina cloth insulation prior to insertion into the tube to prevent gases from bypassing the catalyst and to prevent radiation losses in the radial direction. Radiation shields made of uncoated (x-Al203 monoliths were placed both up and downstream of the catalyst to prevent radiation losses in the axial direction. This region of the reactor was then insulated with alumina fiber so that experiments were run within 50 ~ of adiabatic temperature. Catalyst temperatures were measured by a Pt/Pt-13% Rh thermocouple placed on the downstream face of the catalyst. As sketched in Figure 1, gases were fed from fuel, oxygen, and nitrogen cylinders with mass flow controllers to control the flow rates to the reactor, a valve to control the pressure in the reactor and a gas chromatograph with both thermal conductivity and flame ionization detectors to measure the composition of the outlet gases from the reactor. All experiments were run at a pressure of--20 psia in order to provide enough pressure to overcome the pressure drop in the gas chromatograph. The outlet gases from the GC were then passed through an incinerator and vented. 2. M E T H A N E TO S Y N G A S When Pt on (:t-alumina was used as a catalyst for methane oxidation, only 60 to 80% H2 selectivity was obtained with 20 to 40% selectivity to H20. However, when Rh was used instead of Pt, the H2 selectivity increased to at least 90%. In all of these experiments the 02 conversion was --99%, and the CH4 conversion was 80 to 95%, depending on the CH4/O2 ratio and the preheat temperature. Typical results are shown in figure 2. 1

_,2, 0.8 ._>

CO

-

=

=

.

,

1

Rh

H2

_>, 0.8

._ ._

._

(1.) 03

0.6

HO C~2u ~

T

_

l

f

1.7

_

~ .

0.9 e-

~

i

1.8

CH 4 / 02 Ratio

1.9

~

.

2

E o

0.4

-r O

0.2

a)

0

2.1

-

1.5

H20 ~

=

..

,,

c%,

;

~

;

1.6

1.7 OH 4

9

_

---...

.s tO

_-

~9 0.6

1.6

L_

.=

H 2 -----~-

03

E o 0.4

-1o 0.2

CO.=

9

9

9

;

1.8 1.9 / 0 2 Ratio _-

._

b)

2

2.1

-

Pt

0 2

02 0.9

--.....

-

E

.o

0.8

>~ 0 . 8 -

8

CH 4

o

o

0.7-

0.7

CH 4

0.6

1.5

16 1.

17, 18 i 1. 1. 1.9 CH 4 / 02 Ratio

I 2

c) 2.1

0.6

1.5

11.6

a7 1.

18 I 1. 1.9 CH 4 / 02 Ratio

I 2

d) 2.1

Figure 2. Conversions and selectivities to CO and H2 over Pt (fight) and Rh (left) on o~alumina monolith catalysts at atmospheric pressure at catalyst contact times o f - 1 millisecond.

Rh is far superior to Pt in producing syngas; and these results can be explained by the lower stability of adsorbed OH on Rh compared to Pt[12]. We have modeled these processes m detail using the known reaction steps from surface science experiments.

494

Recent experiments using higher flow velocities and with other metals and support materials have also been used to characterize this process in detail[8,9, 13]. We find that the reactions on Rh can be successfully operated at pressures up to 5.5 atm, which suggests that it should be possible to produce syngas by catalytic oxidation commercially at the 20 to 30 atm pressures necessary for methanol or hydrocarbon synthesis. 3. ETHANE TO ETHYLENE In the above experiments with CH4 oxidation, Pt produces large amounts of H20, while Rh produces typically a 10/1 H2/H20 ratio. With higher alkanes Rh also produces largely syngas, but with Pt, a major product is olefins. For all fuel/O2 feed composition above the syngas ratio, solid graphite is thermodynamically stable at all temperatures, and at lower temperatures it is an equilibrium product even at the syngas composition. Therefore, the reaction: C2H6 + 1/202 ~ C2H4 + H20 should not produce appreciable ethylene in any reactor which is allowed to come to thermodynamic equilibrium at any temperature. At the 2/1 feed ratio, thermodynamics predicts mainly CH4 as the carbon containing product, but over one-half of the carbon at equilibrium should form solid graphite. This is of course totally unacceptable for a catalytic reaction process because the carbon would cover the catalyst surface and eventually plug the porous monolith. 0.75 .Z"

0.9

0o7 .=

"~ 0.65 f./) .r.~. r 0.60.55

1.4

(J o"

I 1.5

116

i8 19 11.7 1. 1. C2H6 / O 2 Ratio

I 2

211

2.2

0.6

.,

lls

1.'6

117 1.'8

119

C2H 6 / 02 Ratio

0.75

0.55

~

b)

211 2.2

Pt-Sn

0.7 + (n t-

0.5

:I:~ r 0.45

"5 ._ .->

0.4 0.35

0.7

0.5

a)

0.6

-t3 ~9 >-

o.6

1.,

15

1.

I

19

1'.~ 1.~ 1'.8 1.

C2H 6 / 02 Ratio

I

~

~'.1

c)

~.~

--....

0.65

0.6

0.55

~.,

16

11~ 1.

17

~.

1'.8 1'.9

I

~

~'.~

d)

~.~

C2H6 / 0 2 Ratio

Figure 3. Conversion, selectivity, and yield of ethylene formation by ethane oxidation on Pt coated alumina ceramic foam monoliths. Up to 70% C2H4 selectivity is obtained at-70% C2H6 conversion for a single pass yield o f - 5 5 % . Addition of Sn to Pt increases the selectivity and alkane conversion significantly. However, when a 2/1 mixture of C2H6 and 02 are passed over a Pt containing ceramic foam catalyst at velocities between 0.1 and 1 rn/sec, the primary carbon containing product is ethylene, and no detectable carbon buildup is detected on the catalyst even after many

495 hours of operation. Thus, at short residence times with Pt, these catalysts produce a highly nonequilibrium products with <5% CH4 and no carbon formed. Typical results for C2H6 oxidation on Pt coated alumina foam monoliths are shown in fgure 3. The 02 conversion is >99% at all C2H6/O2 ratios from 1.0 to 2.0, and the C2H6 conversion goes from 0.9 to -0.7. The carbon atom selectivity is shown in figure 3b. Approximately 65% goes to C2H4, with 20% CO and 10% CO2. These data can be reproduced on any foam ceramic monoliths with Pt loadings from 0.1 to 20% by weight. Also, the conversions and selectivities do not vary significantly with flow rate for residence times from 0.1 m/sec to > 10 rn/sec. This product distribution is predicted by a [3 elimination of H from the adsorbed C2H5 intermediate in the reaction sequence: C2H6(g) ~ C2H5(s) ---~C2H4(g) rather than the further decomposition reactions of adsorbed ethyl: C2Hs(s) ~ C2Hx(s) ~ CHx(s) ~ C(s) ---~CO(g). Since the heat of adsorption of adsorbed C2H5(s) has been estimated to be -25 kcal/mole, the predicted adsorption lifetime and times for any of these processes is estimated to be -10 -9 sec. The key step is clearly the removal of adsorbed C2H4 from the surface and its escape from the monolith before it decomposes. We have done considerable modeling of these processes using data for the elementary steps in the reaction[ 14]. These models confirm that processes occur very quickly so that coverages of all species are extremely low except for O(s) near the entrance of the monolith where the 02 partial pressure is high, and C(s) after all 02 has been consumed. Thus we predict that the reaction is probably initiated by C2H6 reacting with O(s) to form C2Hs(s) and OH(s). The adsorbed C2H5 can dehydrogenate to form C2H4 which desorbs. We suggest that the C2H4 can escape the hot monolith in spite of many collisions with the surface because the later regions of the catalyst are covered with nearly a monolayer of carbon which is less reactive in adsorbing and decomposing C2H4 than would a clean Pt surface. The absence of any multilayer buildup of carbon we attribute to an above equilibrium formation of H20 which promotes the steam reforming reaction C(s) + H20 ~ CO + H2, and this reaction purges the surface of adsorbed carbon which must be formed continuously by decomposition of some of the C2H4, and thus maintains a steady state carbon coverage of less than one monolayer. Figure 3 also shows conversion and selectivity to C2H4 with Sn added to the Pt catalyst[ 15]. Both the alkane conversion and the selectivity to olefins increase significantly with added Sn. X-ray diffraction and XPS of the Pt-Sn catalyst indicate intermetallic compound formation rather than fcc metal, and this surface evidently increases the alkane conversion and reduces the decomposition of olefins. 4. CYCLOHEXANE TO OLEFINS We have also examined olefin formation from higher alkanes. Propane and butane also produce up to 70% selectivity to olefins on Pt monolith ceramic foam monoliths at nearly 100% 02 conversion with alkane conversions of typically 80% at comparable flow rates and catalyst temperatures to those used for C2H6. However, olefins from these higher alkanes exhibit considerable cracking, with C2H4 the dominant product except at low temperatures and excess fuel. However, isobutane produces primarily isobutylene and C3H6 with little C2H4. These product distributions can be explained by a surface reaction mechanism in which the alkane first adsorbs on the Pt surface by breaking a C-H bond to form the adsorbed alkyl. The adsorbed alkyl then can undergo 13 elimination (breaking of a bond on the

496

second carbon atom from the Pt-C bond). If this is a C-H bond, the olefin from the parent alkyl is produced, but, if a C-C bond is broken, cracked olefins should be produced. Since the C-C bond is generally weaker than the C-H bond, one expects cracked olefins as expected. However, for alkanes larger than C4H10, somewhat different behavior is observed[ 16]. First we note that these are more difficult experiments to run because the alkane must be vaporized and mixed with 02 in the reactor. The higher alkanes are more flammable, and this produces considerable problems with flames in the tube, both before and after the catalyst. This can be overcome by careful mixing, but the reaction conditions for successful catalytic formation of olefins are narrower, and we typically used air rather than pure 02 to reduce problems with flames. Cyclohexane produced a different pattern of olefins than the linear and branched alkanes. Figure 4a shows plots of conversions of cyclohexane and 02 and the selectivity to olefins versus the C6H12/O2 ratio. The temperature also decreases as the C6H12/O2 ratio increase, as expected. The total olefin selectivity is -65% which is typical of other alkanes. However, it is seen that the 02 conversion with cyclohexane is no longer complete, and up 1

850 =L

.-->" 0.8

02 "

' "

0.4 Ca

800

"

Olefinsl

0.3

"~ 0.6 03 ~ 0.4

650 c3

O 0.2

600

0

0'6 0'7 0'8 0'9

I

I

""

0.1

550 1.2

0

__ 0.5

0.6

0.7

CeH~2 / 02 Ratio

0.3

._ ._

0.1 0.05

1.1

1.2

C 40lefins

o~,o.1

~, o.15

03

1

1,3--C4H 6

0.15 _>,

0.2

0.8 0.9 C6H~2 / O 2 Ratio

0.2

C 60lefins

0.25

I

.___5_'

(1) 03 _-. . ~ .

C6H6

0.05

/~'~6He

o: - - - - ~ , , 0.5 0.(5 0.7

, 0.8 0.w Cell2 / O 2 Ratio

, 1

~

,-----" 1.1 1.2

o[ 0.5

I~4H8 0.6

"=

,-~

0.7

= 0.8

27,H~ r 019

C8H12102Ratio

1

-

1.1

1.2

Figure 4. Conversions and selectivities of cyclohexane oxidation. Up to 70% selectivity to total olefins are observed with cyclohexene, butadiene, and ethylene as the major products. At high 02 and higher temperatures ethylene and butadiene dominate, while at lower 02 and lower temperature, cyclohexene dominates. Much more cyclohexene is obtained than benzene, and much more ethylene and butadiene are produced .than propylene. to -10% 02 breakthrough in the product. The 02 conversion can be increased by using pure 02 rather than air and by preheating (although both increase the occurrence of flames), although complete 02 conversion cannot be obtained in these experiments. We have no definitive explanations for O2breakthrough from larger alkanes, but we suspect it is due to larger coverages of carbon which inhibits 02 adsorption and allows its breakthrough.

497 The most interesting feature of cyclohexane oxidation is the types of olefins which are obtained. At low temperatures and higher C6H12/O2 ratios, we observe primarily cyclic C6 olefins, as shown in figure 4b. Figure 4c shows that the C6 olefins are primarily cyclohexene with much smaller amounts of cyclohexadiene and benzene. Thus, the singly dehydrogenated olefin dominates over the multiply dehydrogenated products, even though thermodynamic equilibrium predicts benzene over cyclohexene. At higher temperatures and lower C6H12/O2 ratios, more cracked products are observed. However, as shown in Figures 5b, 5c, and 5d, the dominant products are C2 and C4 rather than C3. Examination of the C4 products shows that the dominant species is butadiene, which is formed with four times the selectivity of all butylenes. The patterns of olefins from cyclohexane thus has a distinctive characteristic which shows that the products are either cyclohexene or ethylene plus butadiene, with very little propylene. Further, the ethylene is approximately equal to the butadiene. This suggests that these products are produced in the simple reactions: cyclohexane ---) C2H4 + butadiene or cyclohexane --->cyclohexene. It is interesting that cracking splits the cyclohexane into C2 and C4 fragments rather than into two C3 fragments. If the adsorbed species were singly bonded cyclohexyl, then [3 scission should produce a secondary alkyl, and a second 13 scission would produce propylene and leave adsorbed isopropyl which should yield a second propylene molecule. The formation of ethylene and butadiene seems to argue that the intermediate which leads to these products is a di-cy adsorbed cyclohexane fragment, from which two 13 scission events would liberate ethylene and leave adsorbed C4 which could desorb as butadiene. 5. C A T A L Y T I C C O M B U S T I O N Complete oxidation of alkanes to carbon dioxide and water can also be achieved in millisecond contact time catalytic reactors by operating in the fuel lean regime. Short contact time combustion has many potential applications[ 17] including power generation in gas turbines, incineration of waste containing air streams, and the development of radiant heat sources. We have demonstrated greater than 99.5% conversion of methane, ethane, propane, and butane to complete oxidation products over platinum coated ceramic monoliths at contact times as short as 2 ms at compositions ranging from 20 to 100% excess air [ 18,19]. These experiments showed that, as long as the catalyst was in the ignited state, very high fuel conversions (>99%) were realized. It was found that the catalyst extinction temperature for methane, ethane, and propane are approximately 1150, 1050, and 800 ~ respectively, for room temperature feeds in air at a residence time of 5 ms. Extinction with these fuels occurred at 5.5, 2.2, and 1.3% fuel in air for room temperature feeds with is much leaner than the 9.5, 5.7 and 4.0 for ~=1. We have investigated the use of this type of reactor as an incinerator and found that it is possible to achieve very high destruction efficiencies by adding methane to an air stream containing volatile organic compounds and passing the resulting gas stream over an ignited monolith catalyst. Methane addition is necessary in order to increase the fuel value of the stream to maintain a sufficiently high catalyst temperature so that extinction of the catalyst does not occur. Figure 5 shows a plot of the data obtained for the incineration of toluene containing air streams at varying residence times and toluene inlet concentrations. For these experiments, methane was added to air streams containing 500-2000 ppm of toluene until the methane concentration was 6.75% (~=0.7). This mixture was then passed over an ignited catalyst at residence times ranging from 3 to 10 ms.

498

1O0

2

g "

8c-

1.5

1

Fm s f

99.98

~ o

o_--. ! O" 1200

5 ms//f 10 ms

g

99.94 .(~ ~

0.5

5"

99.92 =

8 '

500

3 ms ~..._...~.-.---~ 5 ms ~---

7 ms 99.96 ,~ o~ ~ 1000

O

0

1400

,a

?lit

_

~ I

~

I

r

lOres ~

&&

I

II

800

? '

1000 1500 2000 Inlet Concentration Toluene (ppm)

99.9

600 I

500

I

/

t

1000 1500 2000 Inlet Concentration Toluene (ppm)

Figure 5. Toluene destruction by methane oxidation in air. For 50 to 2000 ppm toluene the toluene breakthrough is less than 0.5 ppm for a destruction of greater than 99.5%. The temperature of the catalyst is - 1100oc, and the velocity before the monolith is ~ 1 rn/sec for a residence time of several milliseconds. Figure 5a shows both the concentration of toluene at the reactor outlet as well as the overall conversion of toluene to carbon dioxide and water as a function of the inlet concentration of toluene and the reactor residence time. This plot shows that the outlet concentration of toluene remains constant at -0.3 ppm regardless of residence time or inlet composition. This results in calculated toluene conversions from 99.94 to 99.99%. Any scatter in the conversion data can be attributed to experimental error since no general trend of conversion with residence time is observed. Figure 5b shows how the catalyst temperature varies with residence time and toluene composition. There is a small dependence of temperature on the toluene composition as would be expected as the fuel value of the stream increases. The figure also shows a large increase in temperature with decreasing residence time. This occurs because at higher flow rates (shorter residence times), heat losses from the reactor to the surroundings become less important. The adiabatic temperature for the compositions examined in figure 5b is approximately 1400oc. It is also important to note that theaddition of toluene allows the catalyst to operate below the extinction temperature for pure methane. We have performed similar experiments using other compounds such as chlorobenzene, acetonitrile, and thiophene in simulated VOC streams and see similar high conversions (>99%) and no apparent catalyst deactivation due to C1, S, or N. The most important variable in these experiments is temperature. High conversions are realized for all cases when the catalyst is ignited, although sufficient fuel must be supplied to the catalyst to maintain the temperature above the extinction temperature. 6. C O N C L U S I O N S Catalytic partial oxidation at very short contact times is a promising route to new chemicals and to catalytic destruction of volatile organic compounds. Conversions of the limiting reaction are >99% at residence times less than 1 millisecond, and low concentrations of undesired products are observed. The results discussed in this paper are summarized in Table 1 below.

499 Table 1 Conversions and Selectivites in Millisecond Oxidation Reactions Reaction

Selectivity

Fuel Conversion

Yield

Syngas

> 0.9

> 0.9

-- 0.9

C2H 6 tO C2H 4

_<0.7

< 0.9

-" 0.6

Alkanes to Olefins

< 0.7

< 0.8

< 0.6

Toluene to CO 2

0.999

> 0.99

> 0.99

C H 4 to

In partial oxidation of CH4 >90% syngas is routinely obtained on Rh while for higher alkanes up to 70% olefins are produced on Pt. Olefins are a nonequilibrium product, and carbon is predicted under all conditions while none is observed. These catalysts have very low area and are very stable over many hours of operation because sintering and poisoning are not problems at these high temperatures. The features of these processes which are essential to successful operation are high mass transfer in the monolith structures and the high heat transfer of the solid which prevents blowout even at very high flow rates by gas preheating through high heat conduction in the solid. In all of these processes the possibility of homogeneous reaction steps must be considered[20]. Partial oxidation of methane-rich flames does produce syngas, and steam cracking is a homogeneous process which produces olefins at comparable selectivities to those described here. We have argued that these results can be explained as occurring by purely surface reaction steps. The main argument against the'presence of significant homogeneous reaction in these systems is that the product distributions are somewhat different than would be predicted by homogeneous free radical reaction steps and that no carbon buildup is observed over many hours of operation. Homogeneous reactions require a buildup of radical species which propagate chain reactions, and these require times of many milliseconds. In the experiments of these studies, the gases go from reactants near 20~ to products containing no 02 in times much less than 1 millisecond, and this does not appear to allow time for significant homogeneous reaction. However, it is possible that the olefin cracking observed is due to homogeneous pyrolysis. Recent experiments with butane oxidation on a single gauze reactor produce high conversions, and primarily butylene and oxygenates. This result can be explained by surface initiated homogeneous reactions with cold gas quenching the products and eliminating continued decomposition of these unstable products. Further experiments designed to exploit homogeneous-heterogeneous processes to produce chemicals with high selectivity have great promise as new routes to chemicals from alkanes.

REFERENCES 1. S. Albonetti, F. Cavani, F. Trifiro, Catalysis Reviews-Science & Engineering 38(4) (1996) 413. 2. H. H. Kung, Advances in Catalysis 40 (1994), 1. 3. S. C. Tsang, J. B. Claridge, M. L. H. Green, Cat. Today 23(1) (1995) 3.

500 4. A.T. Ashcroft, A.K. Cheetham, J.S. Foord, M.L.H. Green, C.P. Grey, A.J. Murrell, P.D.F. Vernon, Nature 344 (1990) 319. 5. E. Morales, J.H. Lunsford, J. Catal., 118 (1989) 255. 6. V. R. Choudhary, et al., J Catal 163(2) (1996) 312-318; J Catal 157(2) (1995) 752-754. 7. D. A. Hickman and L. D. Schmidt, Catalysis Letters 17 (1993) 223; Science 259 (1993) 343. 8. P. Tomiainen and L. D. Schmidt, J. Catalysis 146 (1994) 1. 9. P. Witt and L. D. Schmidt, J. Catalysis 163 (1996) 465. 10. M. Huff and L. D. Schmidt, J. Phys. Chem. 97 (1993) 11815; Catalysis Today 21 (1994) 443: J. Catalysis 149 (1994) 127; J. Catalysis 155 (1995) 82. 11.

D. Goetsch and L. D. Schmidt, Science 271, 1560-1562, March 15 (1996); Heterogeneous Hydrocarbon Oxidation, ACS Symposium Series 638, 124-139 (1996).

12. D. A. Hickman and L. D. Schmidt, A.I.Ch.E Journal 39 (1993) 1164. 13. A. Dietz III and L. D. Schmidt, Catalysis Letters 33 (1995) 15. 14. M. Huff and L. D. Schmidt, AIChE Journal, to be published. 15. C. Yokoyama, S. Bharadwaj, and L. D. Schmidt, Catalysis Letters 38 (1996) 181. 16.A.G. Dietz llI, A. F. Carlsson, and L. D. Schmidt, "Partial Oxidation of C5 and C6 Alkanes at Millisecond Contact Times", submitted. 17. L.D. Pfefferle, W.C. Pfefferle, Catal. Rev.mSci. Eng., 29(2&3) (1987) 219. 18. C. T. Goralski Jr. and L. D. Schmidt, Catalysis Letters 42 (1996) 47. 19. C. T. Goralski Jr. and L. D. Schmidt, "Incineration of VOC Containing Air Streams at Very Short Contact Times", to be published. 20. D. G. Vlachos, L. D. Schmidt, and R. Aris, AIChE Journal 40 (1994) 1005.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

501

C a t a l y t i c O x i d a t i v e D e h y d r o g e n a t i o n o f I s o b u t a n e in a Pd M e m b r a n e R e a c t o r T. M. Raybold and M. C. Huff Center for Catalytic Science and Technology, Department of Chemical Engineering, University of Delaware, Newark DE 19716

The production of isobutylene from isobutane over a Pt/a~-A1203 catalyst in a palladium membrane reactor has been investigated at relatively short contact times (10 -1 seconds). The Pd membrane is being used for the selective removal of H2. Molecular oxygen has been added to the reaction side of the membrane module to produce favorable kinetics for the production of isobutylene and to reduce the external heat load required to maintain the reactor temperature of 400 - 650~ With oxygen addition (iC4H10:O2 = 1.0 to 2.0), the isobutane can react through both the endothermic dehydrogenation reaction and the exothermic oxidative dehydrogenation reaction, potentially creating an autothermal process. The continuous removal of hydrogen through the Pd membrane relieves the equilibrium limitation of the dehydrogenation reaction allowing increased isobutylene yields.

1.

INTRODUCTION

Isobutylene is typically produced by catalytic dehydrogenation over a Cr203-A1203 catalyst [1, 2] in the absence of free oxygen. iC4Hlo --~

iC4H8 + H2

(1)

Typically, this reaction is conducted at high temperature to achieve appreciable conversion in a reasonably short reaction time. However, under these conditions where the reaction kinetics are fast, this reaction tends to be thermodynamically limited. Therefore, it has been studied in a Pd membrane reactor [ 1, 3, 4] where the H2 is continuously removed. Since the rate of the dehydrogenation reaction is slow relative to the rate of hydrogen removal through the Pd membrane, these membrane systems remove the thermodynamic limitation and are instead kinetically limited [5]. Oxidative dehydrogenation is kinetically fast and may relieve this limitation [6]. Previous research has shown that substantial amounts of olefins are formed by catalytic oxidative dehydrogenation in a high temperature (1000~ short contact time (1-10 milliseconds) reactor for small alkanes [7, 8]. 1

C2H6 + ~ 02 ---) C2H4 + H20

(2)

This reaction is in direct competition with the partial oxidation to synthesis gas, complete combustion, and carbon formation on the surface with the olefin essentially and intermediate en route to more stable products. Thermodynamic equilibrium calculations under these severe reaction conditions, predict the formation of carbon, CO, and H2, and only trace amounts of C2H4. Clearly, thermodynamic equilibrium is not obtained since the system is kinetically

502

CO. Escape ~I~.~------~-H2 ~ to 02 # )y sweep J

~

9 iC4Hlo ~ ~ T ~ H20

~,~

i hermodynamic equilibrium iC4H8

Figure 1: Schematicof the oxidative dehydrogenation of isobutane reaction system. limited at these short contact times and olefin yields in excess of the maximum thermodynamic predictions are achieved. However, the larger alkanes (e.g., iC4H10) are more difficult to activate [6, 9] and longer reaction times are required for high conversions. 1 iC4H10 + ~ 02 ----) iC4H8 + H20

(3)

Partial oxidation to syngas also occurs which produces significant quantities of H2. iC4Hlo + 2 02 ----)4 CO + 5 H2

(4)

These products (iC4H8, CO, and H2) are more reactive than the iC4HlO, so they behave as intermediates and continue to react and approach thermodynamic equilibrium. Thus, the co-production of isobutylene and hydrogen in oxidative dehydrogenation causes this system to be thermodynamically limited. This reaction system is presented schematically in Figure 1. By using a membrane reactor, we continuously remove the co-produced H2 while still benefiting from the fast oxidation kinetics. 2. 2.1

EXPERIMENTAL

Reactor Design The oxidative dehydrogenation of iC4H10 has been studied in a Pd foil, radial flow membrane reactor using a supported Pt catalyst. A schematic of the reactor is shown in Figure 2. A 0.075 mm thick Pd foil is sealed between two copper gaskets in a gas tight stainless steel module to separate the reaction and sweep sides. The reaction side is loaded with 18.5 grams of 1/8 inch catalyst pellets, which contain 0.5 wt. % Pt/t~-A1203 (eggshell). The reactants enter the catalyst bed at the axis and move outward and upward to exit through the annulus at the axis of the module. The sweep side contains a cone of 4 mm Pyrex beads. The beads lead to a more uniform distribution of flow of the sweep gas, while the cone directs the flow toward the membrane and allows for quite efficient removal of the H2 that has diffused through the Pd membrane. The total flow rate on the reaction side is held constant at 1.00 slpm by mass flow controllers with accuracies for each component (iC4H10, 02, and N2) of + 0.01 slpm. The flow rate on the sweep side (N2) was maintained at 4 + 0.1 slpm. It is important to note that both Pt and Pd are active catalysts for the reactions of hydrocarbons, particularly in the presence of 02. Ideally, we would like to use the catalytic properties of the Pt without also seeing the catalytic effects of the Pd on the product distribution, since Pd

503

Reactants

Thermocouple

Products

Reaction side

~J~i!~i~::~!~i~~

/

~i~."~i

Sweep side

Pt/ot-AI203 catalyst

t

!

~

S w e e p ou

Sweep in Figure 2: Schematic of the radial flow Pd membrane reactor used for these studies. A circular Pd foil (.075 mm thick) separates the reaction side from the sweep side. The reaction side is loaded with 18.5 grams of 0.5 wt. % Pt/o~-A1203 1/8 inch pellets.

tends to lead to carbon formation. Therefore, although we are trying to use the Pd foil simply for its H2 permeability and not its catalytic activity, it is impossible to separate these actions in the current reactor design.

2.2

Reactor Operation

Prior to each experimental trial, the reactor is externally preheated to 200~ with inert flow on both the reaction side and the sweep side of the membrane at the desired total flow rate of the experimental trial. The reactants are then introduced to the reaction side maintaining the desired total flow rate. The net reaction that occurs is exothermic causing the reactor temperature (measured by an alumel-chromel thermocouple) to increase. At steady-state, the reactor temperature measures between 400 and 650~ depending on the iC4H10:O2 feed ratio and the level of dilution. This is not an isothermal reactor; at steady state, sizable temperature gradients exist within the catalyst bed. The temperature reported here is the temperature at the axis of the reactor where the feed stream meets the catalyst bed. After each run at a particular iC4H10:O2 feed ratio and the level of dilution, the catalyst and membrane are regenerated by de-coking in preparation for the next trial. This is accomplished by passing air at 1 slpm through the reactor while the reactor is still hot until COx formation ceases. This process takes 15-30 minutes and leads to reproducible data.

2.3

Product Analysis

The compositions of the effluent from both the reaction side and the sweep side of the module are measured by an on-line HP 6890 gas chromatograph. All of the chemical components except H20 are measured relative to calibration standards. The fraction of H20 in the product is calculated from a material balance on oxygen atoms. The remaining material balances (C and H) typically close to within 15%.

504 Typical data is shown in Figure 3. Oxygen is the limiting reactant and oxygen conversions are always greater than 97%. The selectivities for the carbon containing species are calculated on a carbon atom basis and the selectivities to H2 and H20 are calculated on a hydrogen atom basis. The major products are indicated in Figure 3. Other products include C3H8, C2H6, and C2H4 (< 4% selectivity each). For figure clarity, these components are not plotted. 3. RESULTS DISCUSSION 50

'

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O x y g e n in the Feed Isobutane conversion, CO selectivity, and reaction temperature are plotted in CO Figure 3 as a function of the iC4H10:O2 ratio in the feed for :~20 the oxidative dehydrogenation of _....._---.iC4H 10 over a 0.5 wt. % Pt/ocoo10 OH 4 A1203 catalyst loaded in a Pd CH . I . . , 13, ,6, I ,. , I ,, , I ,. , I, membrane reactor. The feed 1 1.2 1.4 1.6 1.8 2 contains 30% N2 diluent, has a 50 " l ' ' ' l ' ' ' , ' ' ' l ' ' ' l ' ' ' l ' total flow rate of 1 slpm and is Rxn. H2 maintained at a pressure of 2 psig. The sweep gas is N2 v f l o w i n g at 4 slpm and maintained at a pressure of 1 psig. Under these reaction c o n d i t i o n s , the d o m i n a n t m products are iC4H8, CO, CO2, oo10 Sweep H2 H 2 0 , and H2 (the latter being present on both the reaction side and the sweep side). All data 1 1.2 1.4 1.6 1.8 2 'l' 600 FTo . , . ' ' , ' ' ' l ' ' ' l ' ' ' , ' ' was collected using a freshly regenerated catalyst and c- 60 O membrane after 48 minutes on550~ E 50 line. >~ 40 The reaction temperature 500~ tracks the iC4H10 conversion O 30 and both decrease at the higher 20 450Cj iC4H10:O2 ratios, further away = 1 0 from the flammability limits. At ..Q o these higher iC4H10:O2 ratios, =1 i , , I , . , I , , , I , , , I , , , I, 400 0 1 1.2 1.4 1.6 1.8 2 less cracking of the iC4H 10 occurs and the selectivity to iCeHlo / 02 iC4H8 increases. However, this Figure 3: Selectivity,iC4H10 conversion, and reaction temperature for selectivity to iC4H8 is somewhat the oxidative dehydrogenation of iC4HI0 over Pt/oc-A1203 in a Pd less than the typical selectivity in membrane reactor as pictured in Figure 2 as a function of the iC4H]0:O2 either dehydrogenation in a ratio. The reaction side contained 30% N2 dilution, had a total flow rate of 1 slpm, and was maintained at a pressure of 2 psig. The sweep side membrane reactor [ 1, 3, 4] or (N2) had a total flow rate of 4 slpm and was maintained at a pressure of oxidative dehydrogenation in a 1 psig. monolithic reactor [6]. Oxidative .

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dehydrogenation in the monolithic reactor occurs at a somewhat higher temperature (800~ and at contact times two orders of magnitude shorter. At the comparably long contact times used in the current membrane reactor module (10 -1 seconds), it is likely that iC4H8 is decomposing before leaving the reactor. Second generation reactor designs will allow for shorter residence times. Figure 3 shows the selectivity to H2 split between the fraction that remains on the reaction side and the fraction that is swept a w a y after p a s s i n g through the Pd membrane. For example, at an iC4H10:O2 ratio o~" 25 of 2.0, the total selectivity to H2 >,, 20 (on an H atom basis) is 42% (38 > ,~,,.,-.-''"~ F)o = 1.45 + 4) with 10% of the total H2 O 15 transferred to the sweep. . . . . .

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3.2 T r a n s i e n t B e h a v i o r of the C a t a l y s t / M e m b r a n e ~ 5 Figure 4 shows the changes in conversion and selectivity for 0 m 0 this reactor as a function of time 0 10 20 30 40 50 ~ 7o on-line for two iC4H10:O2 (F/O) ratios each with 10% N2 dilu60 C .I'" F/O- .oo tion. For simplicity, only the 0 "~ 50 iC4H 10 conversion, reaction t emperature, and iC 4H 8 > 40 cO selectivity are shown. High F/O = 1.45 0 30 conversion and low selectivity = 20 are observed for F/O = 1.00 while low conversion and high 10 .,Q selectivity are observed for F/O 0 0 = 1.45. In both cases, the con10 20 30 40 50 0~'850 version and selectivity both inL... crease with time on-line. The re.~600 action t e m p e r a t u r e is very :3 closely linked to the conversion ~550 and it increases with time as O. well. Although there is some E ~500 change in the selectivities to each E of the products during this o450 "warm-up" stage, the amount of O H2 removed on the sweep side ~400 .... I .... I .... ! .... I .... of the module does not indicate 0 10 20 30 40 50 any loss in permeability of the Time Online ('min) Pd membrane. After about 30 Figure 4: Selectivity,iC4Hlo conversion, and reaction temperature for minutes, there is little change in the oxidative dehydrogenation of iCnHlo over Pt/t~-A1203 in a Pd c o m p o s i t i o n of the product membrane reactor as a function of time on stream. The reaction side contained 10% N2 dilution, had a total flow rate of 1 slpm, and was stream or the reaction temperamaintained at a pressure of 2 psig. The sweep side (N2) had a total flow ture. All of the data presented in rate of 4 slpm and was maintained at a pressure of 1 psig. Data is the other figures corresponds to shown for iC4H10:O2 ratios of 1.0 and 1.45. E

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506

data collected after 48 minutes on-line. 3.3

Hydrogen Removal Trials both with and without H2 removal were conducted under the same experimental conditions in the reactor shown in Figure 2. For trials without H2 removal, the Pd foil was backed by a solid copper disk which 25 blocked H2 removal from the re.i 20action side of the module. The , I> 0 Pd foil was present in the reactor 0 -~ 1 5 to account for any catalysis that may be occurring on the surface 0 r- 1 0 of the foil. As shown in Figure 0 >., 3, the for almost all composi:~ 5 tions, --10% of the total H2 pro..Q 0 duced diffuses through the Pd -0 membrane to the sweep side of 20 10 30 the reactor. For membrane reaco~" 70 tors, this is a very small amount t-- 6 0 o of H2 removal. However, resi.m -.,-.,-i,.. 50 dence times in more traditional O > membrane reactors are typically ,-- 4 0 o much longer than the residence 0 30times used in this study. c 20Figure 5 shows the effect of H2 removal on iC4H8 selectiv..Q 10 1.015 o ity, iC4H10 conversion, and re0 action t e m p e r a t u r e for two 20 10 30 iC4H10:O2 ratios, with varying 6-650 t._,. levels of N2 dilution. In all o cases, the selectivity to iC4H8 is t,,,,. = 600 greater at the higher iC4H10:O2 I... ratios and the iC4H10 conversion 0 -550 is greater at the lower iC4H 10:02 0 ratios. At all compositions, hyt-

5]

ro 500

.i

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n - 9450

drogen removal leads to increased iC4H8 selectivity and increased iC4H I O conversion.

Therefore, by removing as little as 10% of the produced H2, this reactor design has shifted the Figure 5: Selectivity, conversion, and reaction temperature with and p r o d u c t d i s t r i b u t i o n t o w a r d without a Pd membrane for continuous H2 removal. The striped bars olefin formation. correspond to the data for the membrane reactor (H2 removal). The solid bars correspond to data collected when the membrane is replaced by an 3.4 D i l u t i o n of t h e Reacimpermeabledisk (without H2 removal). In all cases, the total flow rate t a n t s on the reaction side was 1 slpm at a pressure of 2 psig. Data is shown for iC4H10:O2 ratios of 1.0 and 1.5 at 10%, 20%, and 30% N2 dilution. F i g u r e 5 also shows the For the trials with the Pd membrane, the flow rate on the sweep side effect of N2 dilution on reaction was 4 slpm of N2 at a pressure of 1 psig. system. In all cases, increased 10

20 30 N i t r o g e n Dilution (%)

507 level of N2 dilution decreases both the iC4H10 conversion and the iC4H8 selectivity. This is largely a temperature effect since the presence of the diluent also reduces the reaction temperature. The differences with dilution are much less apparent when additional heat is added to the system to compare dilution levels at equivalent reaction temperatures. 4.

CONCLUSIONS

The kinetics of oxidative dehydrogenation in a membrane reactor are much more favorable than the kinetics of catalytic dehydrogenation (in the absence of oxygen). The goal of this research is to balance the reaction rate and the rate of H2 removal from the reactor by adjusting the residence time and availability of oxygen to create a highly efficient membrane reactor for the production of isobutylene. We have achieved over 65% isobutane conversion with 13% selectivity to isobutylene at compositions just outside the flammability limits, using a Pd foil, radial flow membrane reactor and a Pt/~-A1203 catalyst. Selectivities were low, mainly due to the decomposition of the product, which can be attributed to the relatively long contact times (I: -- 1 sec) of the reactor. We have also shown that, although the current reactor and membrane configuration only allowed for a 10% hydrogen removal, the removal of hydrogen substantially increased both isobutane conversion and isobutylene selectivity. REFERENCES

1. 2. 3. 4. 5. 6. 7. 8. 9.

T. Matsuda, I. Koike, N. Kubo, E. Kikuchi, Appl. Catal. 96 (1993) 3-13. S. Udomsak, R. Anthony, Ind. Eng. Chem. Res. 35 (1996) 47-53. J. Deng, J. Wu, Appl. Catal. A 109 (1994) 63-76. E. Gobina, R. Hughes, J. Membr. Sci. 90 (1994) 11-19. B.A. Raich, PhD Thesis Chemical Engineering, University of Delaware, Newark 1995. M. Huff, L. D. Schmidt, J. Catal. 155 (1995) 82-94. M. Huff, L. D. Schmidt, J. Phys. Chem. 97 (1993) 11815-11822. M. Huff, L. D. Schmidt, J. Catal. 149 (1994) 127-141. H. Armendariz, G. Aguilar-Rios, P. Salas, M. A. Valenzuela, I. Schifter, H. Arriola, N. Nava, Appl. Catal. A 92 (1992) 29-38.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

509

Chemoselective catalytic oxidation of polyols with dioxygen on gold supported catalysts Laura Prati and Michele Rossi Dipartimento di Chimica Inorganica Metallorganica e Analitica, UniversitY. di Milano e Centro C.N.R., via Venezian 21, 1-20133 Milano, Italy

Supported gold catalysts in the presence of dioxygen have shown high selectivity toward polyols monooxygenation. In fact, commercially useful products such as glycolic and lactic acids from, respectively, ethane-l,2-diol and propane-l,2-diol, can be advantageously obtained. Studies on the influence of the support and preparation methods are reported along with comparison tests involving commercial Pd and Pt catalysts.

1. INTRODUCTION The oxidation of the alcohols to carbonylic or carboxylic derivatives is of great practical interest in the synthesis of commercially useful products, as is evident from the abundant literature that has appeared in recent years [ 1]. Particular interest has been shown in the catalytic oxidation of polyols with dioxygen using supported platinum-group metals as the catalysts. The most studied metals have been palladium and platinum which are, however, often affected by deactivation problems [2]. The introduction of cocatalysts such as bismuth or lead represents an enhancement in the use of these catalysts having the double effect of increasing catalytic activity and improving catalyst life [3]. One important application is in the field of carbohydrate transformations: the catalytic oxidation of D-glucose to D-gluconic acid represents an economically competitive route with respect to biochemical oxidation [4]. This new process is the result of extensive studies on the selective C 1-hydroxyl group oxidation in the presence of 0 2 using a Pt or Pd catalyst modified with cocatalysts [5]. In spite of the abundant literature dealing with carbohydrate transformations, there is relatively little on lower polyols [6]. In the case of glycerol particular attention has been paid to the influence of cocatalysts, like Bi, on the selectivity of Pd and Pt catalysts, which changes the hydroxyl group oxidation from primary to secondary producing dihydroxyacetone with 70 80% selectivity [7]. Industrial and academic investigations have dealt with the catalytic oxidation of C2 and C3 diols. In ethane-l,2-diol oxidation, a few studies indicate a limited application of platinum and palladium catalysts because of the overoxidation that normally occurs with C-C bond

510 cleavage, forming CO 2 via HCOOH. In fact, the patent literature reports mostly catalytic processes on protected subtrates such as methoxy glycol and polyethylene glycol [8]. The same type of limitation was encountered in the use of a copper-zinc catalyst in oxidizing ethane-1,2diol to glycol aldehyde where high selectivity (90%) was obtained only at low conversion (23%) [91 . Other catalytic systems have been employed to achieve high selectivity with more acceptable conversion. Thus, an Ag/Si-C catalyst was used for glyoxal production with 73% selectivity at 96% conversion [10] whereas, to the best of our knowledge, the most notable result in glycolic acid production from ethane-1,2-diol was claimed by using an Ir on carbon catalyst operating at 10 atm and 80~ (87% selectivity at 98% conversion) [ 11 ]. Also in the case of the C3-diol catalytic oxidation there is a lack of literature concerning the selective oxidation of the hydroxyl groups. According to a recent patent the selective oxidation of propane-1,2-diol has been claimed for either the primary or the secondary hydroxyl group, but no examples were reported for primary ones [6b]. Palladium on carbon has been reported to oxidize propane-l,2-diol in a non selective manner [12] whereas a patent claims 86% selectivity in propane-1,3-diol oxidation to 3-hydroxy-propanoic acid using a commercial Pd/C catalyst [6a]. This paper reports our studies on the selective oxidation of ethane-l,2-diol and propane1,2-diol using new gold-based catalysts. As an extension, commercial Pd and Pt catalysts were tested under similar reaction conditions to compare their activities and selectivities in the synthesis of important commercial products, such as glycolic acid and lactic acid.

2. EXPERIMENTAL

2.1 Reagents and apparatus Ethane-1,2-diol and propane-1,2-diol were of the highest purity from Fluka and were used without any further purification. NaOH was 99.9% pure from Merck and stored under nitrogen. Gaseous oxygen from SIAD was 99.99% pure. The gold metal was of the highest purity grade from Fluka. Commercial 5% Pd/C was supplied by Montecatini Tecnologie (MPT5 catalyst), 5% Pt/C by Engelhard. Activated carbon (5-100 kt) had a specific area of 1200 m2/g,; AI20 3 (type 507C 100-125 mesh), SiO 2 (220-440 mesh) and CeO 2 (>99%) were from Fluka. Reactions were carried out at the appropiate temperature in a thermostatted glass reactor (30 cm 3) provided with an electronically controlled magnetic stirrer connected to a large reservoir (5,000 cm 3) containing oxygen at 2 atm (310 KPa). The oxygen uptake was followed by the use of a mass flow controller connected to a PC through an A/D board, plotting a flow/time diagram.

2.2 Oxidation procedure

2.2.1 Oxidation of ethane-l,2-diol The substrate (0.50 g, 8.06 mmol), NaOH (0.96 g, 24 mmol) and the catalyst ( diol/metal = 1000) were mixed in distilled water (total volume 10 ml). The reactor was pressurized at 2 atm of 0 2 and thermostatted at the appropriate temperature. The mixture was stirred and the samples analyzed at various times by HPLC and 13C-nmr.

511

2. 2.2 Oxidation of propane-1, 2-diol The substrate (0.518 g, 6.8 retool), NaOH (0.82 g, 20.4 mmol) and the catalyst ( diol/M - 100 or 1000) were mixed in distilled water (total volume 10 ml). The reactor was pressurized at 2 atm of 0 2 and thermostatted at the appropriate temperature. The mixture was stirred and the samples analyzed at various times by HPLC and/or 13C_nmr.

2.3 Catalyst preparations 1% Ir/C was prepared as previously described [ 11 ] using IrC14 solution ( Ir 44.8 mg/ml). Gold catalysts were prepared using a HAuCI4 0.1M solution obtained by dissolving 1.97 g of metal gold in a minimum amount of a 3:1 (v/v) mixture of concentrated HC1 and HNO 3 and then diluted to 100 ml with distilled water. After reduction all the catalysts were filtered and then washed with hot water to obtain chloride ion free catalysts. The catalysts were used in wet form.

2.3.1 GoM on oxides 1% Au on Al203 and SiO2 were prepared by suspending the support (2 g) in water (3 ml) and then by adding the gold solution (1 ml). The slurry was evaporated to dryness and then calcined under H 2 at 250~ for 3h. 1% Au/otFe20 3 was prepared by the coprecipitation method as previously reported [ 14 ]. 1% Au/CeO 2 was prepared using the incipient wetness impregnation method (see below). 2.3. 2 GoM on carbon a) Impregnation method Carbon (2 g) was suspended in 10 ml of distilled water and while stirring the solution of gold (1 ml) was added. The mixture was allowed to stand at room temperature for lh then heated to 70~ and reduced by adding HCHO 37% (1.5 ml) dropwise. b) Incipient wetness impregnation The support (2 g) was impregnated with 1 ml of 0.1M HAuCI 4 diluted with distilled water to a volume equal to the pore volume of carbon. The suspension was mixed for 20 min then put in a hot solution of HCOONa ( 20 ml of water and 200 mg of formate). c) Precipitation method The solution of HAuC14 0.1M (1 ml) was diluted with distilled water (10 ml) and a saturated solution of Na2CO 3 was added until pH8 was reached. To the stirred solution carbon (2 g) was added. The slurry was allowed to stand for lh then heated to 70~ and reduced by adding HCHO 37% (1.5 ml) dropwise. Alternative reductive agents used were HCOONa (200 mg) or NaH2PO 2 (200 mg).

2.4 Analysis of products The products were identified by comparison with authentic samples. Quantitative analyses were performed by either HPLC or 13C-nmr analyses, using an internal standard (propionic acid). 2.4.1 HPLC analysis Analyses were performed on a Varian 9010 instrument equipped with a Varian 9050 UV detector (210 nm) using a Merck Lichrospher 100 RP18 (5~tm) column. An aqueous nBu4NHSO 4 5mM (1 ml/min) as the eluent was used for ethane-l,2-diol reaction products

512 and NaH2PO4/H3PO 4 0.15M (pH 2.47) ( 0.7 ml/min) for propane-l,2-diol reaction products. Samples of reaction mixture (0.1 ml) were diluted (10 ml) by using the eluent after adding the standard.

2. 4. 2 13C_NMR analysis Spectra were recorded in water on a Varian 200 MHz. Samples of reaction mixture were neutralized with HC1 12N before adding the standard.

3. RESULTS AND DISCUSSION Gold is reported to weakly interact with molecular oxygen, inhibiting subsurface oxygen diffusion and some studies (surface enhanced Raman spectroscopy, microgravimetric and temperature desorption techniques) indicate that molecular oxygen adsorption is preferred to the atomic one [13]. However most theoretical studies have been conducted on smooth gold surfaces or large gold particles. On the other hand it has been reported that the surface chemistry and reactivity change drammatically for supported gold particles [ 14]. The catalytic behavior of gold has recently been reviewed by Haruta who highlights the strong dependence of gold activity on the type of support, the preparation methods and the size particles during CO oxidation, hydrocarbon combustion and other selective partial oxidation [ 15]. With the aim of exploring the activation of 0 2 toward the oxidation of vicinal diols we tested different supported-gold catalysts under mild conditions. By working in neutral aqueous solution of ethane-1,2-diol up to 100~ and 2 atm of 0 2 in the presence of 1% Au supported on active carbon there was no oxidation, whereas in alkaline solution a smooth oxygen uptake at 50-90~ was observed. HPLC and 13C-NMR analyses of the reaction products showed quite good slectivity toward monooxygenation. In fact, operating with a substrate/metal ratio of 1000 (Table 1, entry 1) very high selectivity in glycolate production was reached (90%) at very high conversion (94%). The catalyst was recycled 10 times without loss of activity, there being a slight decrease in selectivity (2-3 %). The attractive result obtained with the gold catalyst prompted us to compare its performance with more convemional catalysts (Pd/C and Pt/C) and with others suggested by previous studies, such as Ir/C [ 11 ] and Cu/C [ 16 ]. Platinum and palladium on carbon were found to be more active but less selective than gold, being affected by overoxidation (entries 2 and 3) to C1 products (HCOOH). By reducing the temperature from 70 to 50~ the selectivity improved, rising to about 70% (entries 6 and 7). Copper on carbon produced relevant quantities of formate even at low conversion (entry 4) owing to the known activation toward C-C bond cleavage [ 16]. Under our reaction conditions the Ir on carbon catalyst, prepared according to the literature [ 11], was inactive (entry 5). Thus, due to the different experimental conditions no comparison between our and the previously reported ethane-1,2-diol oxidation with Ir/C [ 11] was possible.

513 Table 1 Oxidation of ethane-1,2-diol with various catalysts entry

catalyst

t(h)

T(~

GLA

OXA

FOR

(mol )

(mol%)

(tool%)

1%Au/C

3

70

85

5%Pt/C

2

70

50

6

5%Pd/C

2

70

57

6

I%Cu/C

6

70

5.4

l%Ir/C

5

70

5%Pt/C

2

5O

68

5%Pd/C

2

50

74

.

,

.

.

0.5

.

conv.%

selec.% GLA

5

94

90

10

100

50

3

100

57

16

54

10

95

71

100

74

,

.

10.5

Reaction conditions: ethane-l,2-diol/M = 1000; NaOH 2.6M; NaOH/ethane-l,2-diol = 3 GLA-glycolate; OXA=oxalate; FOR=formate. Na2CO 3 was also formed.

The influence of the support on the gold catalyzed oxidation of ethane-l,2-diol is very relevant. In fact a comparison of the results in Table 2 shows carbon (entry 1) to be peculiar in its high activity, with respect to A120 3, CeO 2, SiO 2 and ot-Fe20 3 (entries 2-5).

Table 2 Influence of support entry

r'

catalyst

t(h)

GLA

OXA

FOR

(tool%)

(~mol%)

(mol%)

conv.%

selec.% GLA

1

I%Au/C

3

85

5

94

90

2

1%Au/Al20~

6

36

23

70

51

3

1%Au/CeO 2

6

22

6

35

63

4

l%Au/SiO 2

6

-

5

1%Au/otFe203

6

17

29

50

34

Reaction conditions: ethane- 1,2-diol/M = 1000; NaOH 2.6M; NaOH/ethane- 1,2-diol = 3; T -- 70~ GLA=glycolate; OXA=oxalate; FOR=formate. Na2CO 3 was also formed. With a view to optimizing the activity and selectivity of the Au/C catalyst, a short screening on deposition and activation methods was carried out. Comparing, in Table 3, three different deposition methods, namely alkaline precipitation (entry 1), absorption from diluted solution of HAuC14 (entry 2) and incipient wetness impregnation (entry 3), followed by reduction of gold to metal, we observed that the first method performed the best. Although all the methods show

514 similar selectivity the precipitation one had the best activity, resulting in, for the given time (2 h), the highest conversion. On the contrary the effect of the different reductive agents (HCHO, HCOONa, NaH2PO2) was negligible. In any case, the activated catalyst had always to be dechlorinated to avoid the known poisoning effect of chloride ion [ 13]. The dependence of activity on the preparation methods observed for Au/C catalysts is presently under investigation. According to the literature data, larger gold particles (10 nm) can be expected with the adsorption method while the other two methods result in smaller particles (2-5 nm) [ 13]. Thus, a linear correlation between gold dispersion and activity seems to be ruled out. Table 3 Influence of preparation methods on the activity of 1% Au/C catalyst

GLA

preparation

(mol%)

1

a

75

-

4

2

b

58

-

2

3

c

64

-

5

entry

OXA

FOR

conv.%

selec.%

(mol%)(mol%)

GLA 93

81 65

,,

70

,

89 91

a. precipitation method b. absorption from diluted solution c. incipient wetness impregnation Reaction conditions: ethane- 1,2-diol/M = 1000; NaOH 2.6M; NaOH/ethane- 1,2-diol = 3; T = 70~ t = 2 h. GLA=glycolate; OXA=oxalate; FOR=formate.

As an extension of the ethane-1,2-diol oxidation, we investigated the catalytic oxidation of propane-1,2-diol. In this case the problem of chemoselectivity, arising from the presence of a primary and a secondary alcoholic function, is of great interest as both hydroxyacetone and lactic acid are products of synthetic importance (Fig. 1) Figure 1. Oxidation products of propane- 1,2-diol

~

OH

OH

[~

OH,-~COOH

0 hydroxyacetone

OH hctic acid

By considering the results reported in Table 4 we can outline that almost total selectivity has been shown by the I%Au/C catalyst toward lactic acid production under mild conditions (70~ 2 atm of 02) (entry 1). However this result Canbe achieved using a low substrate/ metal ratio (100) because at higher values (1000) the selectivity drops to 90% (entry 2).

515 By using palladium on carbon, from previous studies on C3-diol oxidation, poor selectivity can be expected in the case of propane-l,2-diol [12] and good selectivity in the case of propane-l,3-diol (86% of 3-hydroxy-propanoate) [6a]. In our experiments, however, oxidizing propane-l,2-diol we obtained good selectivity in lactic acid using either Pd/C or Pt/C (Table 4, entries 4 and 6). In these cases, higher substrate to metal ratio (1000) produced higher selectivity. Acetate was the main byproduct and a small amount of pyruvate was observed only with Pd/C catalyst. Table 4 Oxidation of propane-1,2-diol with various catalysts entry

catalyst

S/M

t(h)

ratio

LA

AC

PYR

(mol%)

(mol%)

(mol%)

conv.%

selec.% lactic a.

1

1%Au/C

100

1

98

-

-

100

98

2

1%Au/C

1000

2

76

8

-

84

90

3

5%Pd/C

100

0.5

68

16

2

86

79

4

5%Pd/C

1000

1

81

11

1

94

86

5

5%Pt/C

100

1

90

10

-

100

90

6

5%Pt/C

1000

2

5

-

83

94

i

78

Reaction conditions: NaOH 2.6M; NaOH/propane-l,2-diol = 3; T = 70~ LA=lactate; AC=acetate; PYR=pyruvate. Among the products hydroxyacetone was not detected.

4. CONCLUSIONS As has already been noted by Haruta [14] the tunable reactivity of gold catalysts by controlling the particle size, the type of support and the different preparation methods widens the potential of such catalysts in different fields of applications. The results presented in our studies point to the synthetic use of supported gold as the catalyst in oxidation reactions of industrial interest. In fact the selective synthesis of important products like lactic and glycolic acids could make the catalytic route competitive with others currently used.

REFERENCES 1. a) R.A.Sheldon, J.K.Kochi Metal Catalyzed Oxidations of Organic Compounds, Academic Press, New York, 1981 b) R.ASheldon Heterogeneous Catalysis and fine chemicals II M.Guisnet, J.Barrault, C.Bouchoule, D.Duprez, G.Perot, M.Maurel and C.Montassier Eds., Elseviers, Amsterdam, 1991, pages 33-54 c) S.T.Oyama, J.W.Hightower Eds. Catalytic Selective Oxidation A.C.S.,Symposium Series, Washington, 1993

516 d) M.Hudlicky Oxidations in Organic Chemistry A.C.S.Monograph 186, Washington, D.C., 1990 2. a) H.E.van Dam, A.P.G.Kieboom, H.van Bekkum Appl.Catal. 33 (1987) 361 b) P.J.M.Dukgraaf, H.A.M.Duisters, B.F.M.Kuster, K.van der Wiele J.Catal. 112(1988)337 3. T.Mallat, A.Baiker Catalysis Today 19 (1994) 247 and references cited therein 4. a) K.Deller, B.Despeyroux Chem.Ind (Dekker) 47 (1992) 261 b) M.Besson, F.Lahmer, P.Gallezot, P.Fuertes, G.Fleche J.Catal. 152 (1995) 116 c) M.Wenkin, R.TouiUaux, P.Ruiz, B.Delmon, M.Devillers Appl.Cat.A:General 148 (1996) 181 5. a) H.E.van Dam, A.P.G.Kieboom, H.van Bekkum Appl.Catal. 33 (1987) 361 b) P.J.M.Dukgraaf, H.A.M.Duisters, B.F.M.Kuster, K.van der Wiele J.Catal. 112(1988)337 c) C.Broennimann, Z.Bodnar, P, Hug, T.Mallat, A.Baiker J.Catal. 150 (1994) 199 d) C.Broenniman, T.Mallat, A.Baiker J.Chem.Soc.Chem.Commun. 1377 (1995) 6. a) A.Behr, A.Botulinski, F.J.Carduck, M.Schneider U.S.Patent 5,321,156 (1994) b)H.Kimura, K.Tsuto U.S.Patent 5,274,187 (1993) 7. a)H.Kimura, K.Tsuto, T.Wakisaka, Y.Kazumi, Y.Inaya Applied Cat.A:General 96(1993)217 b)R.Garcia, M.Besson, P.Gallezot Applied Cat.A:General 127 (1995) 165 c)H.Kimura, K.Tsuto U.S.Patent 5,274,187 (1993) 8. a) M.Nozue Jpn.KOKAI TOKKIO KOHO JP 63,211,251 (1988) b)M.B.Libman, V.F.Shvets, Yu.P.Suchov Khim.Prom.-st. 9 (1988) 520 c)M.Nozue Jpn.KOKAI TOKKIO KOHO JP 04,342,559 (1992) 9. T.Seto, M.Odagiri, M.Imanari Jpn.KOKAI TOKKIO KOHO JP 03,279,342 (1991) 10.P.GaUezot, S.Tretjak, Y.Christidis, G.Mattioda, A.Schouteeten J.Catal. 142 (1993) 729 11.T.Oku, Y.Onda, H.Tsuneki, Y.Sumino Jpn.KOKAI TOKKIO KOHO JP 07,112,953 (1995) 12.T.Tsujino, S.Ohigashi, S.Sugiyama, K.Kawashiro, H.Hayashi J.Mol.Cat. 71 (1992) 25 13.D.I.Kondarides, X.E.Verykios J.Catal. 158 (1996) 363 14.M.Hamta, N.Yamada, T.Kobayashi, S.Iijima J.Catal. 115 (1989) 301 15.M.Haruta Catalysis Today, inpress 16.M.Lanfranchi, L.Prati, M.Rossi, A:Tiripicchio J. Chem. Soc.Chem.Commun. 1698 (1993)

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama,A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

517

P r o m o t i n g e f f e c t s of b i s m u t h in c a r b o n - s u p p o r t e d b i m e t a l l i c P d - B i c a t a l y s t s for the selective o x i d a t i o n of glucose to g l u c o n i c acid M. Wenkin a, C. Renard a, p. Ruiz b, B. Delmon b and M. Devillers a, * Universit6 Catholique de Louvain, a Laboratoire de Chimie Inorganique et Analytique, place Louis Pasteur, 1 b Unit6 de Catalyse et de Chimie des Mat6riaux Divis6s, place Croix du Sud, 2 B-1348 Louvain-la-Neuve, Belgium

Experiments are carried out to improve the u n d e r s t a n d i n g of the behaviour of Bi-promoted P d / C catalysts during their use in the selective oxidation of glucose to gluconic acid by 02. Supported Bi(5 wt.%)-Pd(5 wt.%)/C catalysts are prepared by deposition from a suspension of several carboxylate precursors in heptane, followed by thermal degradation under N2 at 773 K. The catalysts are characterized by XRD, XPS and SEM-EDX. Because significant amounts of bismuth are leached from the catalysts under the reaction conditions, recycling experiments are performed to evaluate the influence of this process on the catalyst lifetime. Whereas the Bi losses are essentially restricted to the first few catalytic runs, the gluconic acid yield, normalized with respect to the catalyst mass, remains constant. Catalytic tests are also conducted in the presence of diethylenetriaminepentaacetate, which is a stronger chelating agent than the gluconate ions, to remove the major part of dissolved Bi from the solution. The behaviour of the bimetallic catalyst is also compared with that of a commercial trimetallic Pd(5 wt.%)-Pt(1 wt.%)-Bi(5 wt.%)/C catalyst.

1. INTRODUCTION Bismuth is known for displaying very attractive properties as promoting element in numerous heterogeneous oxidation catalysts and namely, in its association with noble metals like palladium and platinum for the selective oxidation of alcohols or aldehydes by molecular oxygen in aqueous solutions [17]. However, the actual origin of the promoting role of Bi and the question whether Bi-Pd alloys are present and do play, or not, a significant role in these catalysts is still under discussion. In addition, the fact that significant amounts of bismuth are leached from these catalysts during their functioning remains a critical point which impedes their broader use. In previous works [8,9], bimetallic Bi-Pd catalysts supported on activated carbon and characterized by various Bi/Pd molar ratios ((Pd+Bi)=10 wt.%) were prepared from the thermal degradation of Bi and Pd acetate-type precursors under nitrogen at 773 K. Because several binary Bi-Pd alloys were heavily suspected in the supported catalysts, three intermetallic compounds, Bi2Pd, BiPd and BiPd3 were also prepared from the same precursors, according to the same

518 procedure, upon thermal heating under nitrogen in appropriate conditions described elsewhere [9]. The catalytic performances of the supported and unsupported catalysts characterized as having the same overall composition were measured and compared for the selective oxidation of glucose to gluconic acid. At constant palladium weight, Bi2Pd was found to be the most active phase, whereas the BiPd3 alloy displayed no activity. For the carbon-supported catalysts, the highest performances were observed for Bi/Pd=l. When comparing the pure alloys with the supported catalysts, the highest absolute yields in gluconic acid were therefore found for different compositions, suggesting the multiphasic nature of the supported catalysts. Furthermore, bismuth was found systematically to dissolve in the reaction medium during the catalytic tests, the losses being significantly more extensive from the monometallic Bi/C than from the bimetallic PdBi/C catalysts. Bi2Pd is the phase that loses Bi to the largest extent, while the promoting element does not dissolve from the inactive phase BiPd3. Glucose and gluconate in solution were shown to be both responsible for Bi dissolution. However, we demonstrated that there was no simple relationship between the extent of Bi losses and the performances of the supported catalysts. The present work reports on further experiments devoted to the role played by dissolved bismuth in the catalytic reaction and its consequences on the deactivation process. Particular interest will be paid to the recycling capabilities of these catalysts. Because the complexing properties of gluconate ions were proved to influence the dissolution process and are therefore suspected to modulate the catalytic behaviour, experiments aimed at trapping the solubilized fraction of the promoting element were carried out by adding a stronger complexing agent than the gluconate ion (KD [Bi2(C6H1007)2(OH)] + = 10-10 [10]). Diethylenetriaminepentaacetic acid (H5dtpa), which is known as a strong chelating agent for Bi3+ (KD [Bi(dtpa)] 2- = 10-30 [11]) was therefore added at various amounts to the reaction mixture in the presence of the bimetallic catalyst. The behaviour of the Bi-Pd/C catalysts used in the present work was also compared with that of a well-established commercial trimetallic Pd(4 wt.%)Pt(1 wt.%)Bi(5 wt.%)/C from Degussa. Because the latter was shown to loose smaller amounts of bismuth than the bimetallic Pd-Bi/C catalysts tested so far [8], this comparison was aimed at finding and understanding the relationship between the catalytic behaviour and the extent of Bi losses. In addition, scanning electron microscopy coupled with energy-dispersive energy analysis (SEM-EDX) has been implemented as a complementary analytical tool for various purposes : (i) to estimate the mean particle size of the metallic particles and look at the eventual influence of the used precursors on these characteristics ; (ii) to investigate more deeply the composition and dispersion anomalies detected by XPS on certain catalysts; (iii) to find experimental evidence for bismuth redeposition on the catalyst surface after use.

519 2. EXPERIMENTAL 2.1. Preparation of the catalysts Carbon-supported (Activated carbon SXplus supplied by NORIT, SBET = 750 m2.g -1, particle size : 0.2-0.1mm) bimetallic and monometallic catalysts were prepared by deposition from a suspension of carboxylate particles in n-heptane chosen as inert organic solvent. Precursors used for the incorporation of the metals were either, palladium(II) acetate (ACROS) and bismuth(III) oxoacetate, BiO(O2CCH3) (synthesized as described elsewhere [8]), or diammine(pyrazine-2,3dicarboxylato-N,O)palladium(II) [12] and tris(monohydrogenopyrazine-2,3dicarboxylato)bismuth(III) (noted Bi(2,3-pzdcH) 3) [13]. Monometallic Pd(5 wt.%)/C and Bi(5 wt.%)/C catalysts : the selected carboxylate precursor of palladium or bismuth (see above) was dispersed in the presence of 7.2 g of the activated carbon in 250 ml n-heptane under ultrasonic stirring for 30 min. The hydrocarbon was evaporated very slowly at room temperature under vacuum and the precursor deposited on the support was subsequently decomposed upon heating under nitrogen at 773 K during 18h. Bimetallic Pd(5 wt.%)Bi(5 wt.%)/C catalyst : The n o n - d e g r a d e d monometallic P d / C catalyst was dispersed in n-heptane under ultrasonic stirring for 30 min in the presence of the bismuth carboxylate containing the same ligand as the precursor Pd compound. After slow evaporation of the hydrocarbon, the bimetallic catalyst was activated upon thermal heating under nitrogen at 773 K d u r i n g 18h. D e p e n d i n g on the p r e c u r s o r used, acetate (Ac) or pyrazinedicarboxylate (pzdc), these catalysts were noted (Ac.PdBi/C) and (pzdc.PdBi/C). The pure BiPd3 alloy was prepared from the acetate precursors according to the deposition procedure described above. The carboxylates were decomposed upon thermal heating under nitrogen at 1173 K during 18h. 2.2. Catalytic measurements 2.2.1. Reaction conditions and analysis of the reaction products The selective oxidation of D-glucose into gluconic acid was selected as catalytic test reaction. The reactor vessel and the experimental conditions were described in detail elsewhere [8]. The pH of the reaction mixture was kept at a constant value in the range 9.25-9.45 by adding a 20 wt.% aqueous solution of sodium hydroxide with an automatic titrator (Star Titrino 718) from METROHM. The base consumption was recorded as a function of time. The glucose solution (72 g glucose in 400 ml water) was heated in the reactor to 50~ Once the temperature was stabilized, the catalyst (mcat = 54-200 mg) was added to the solution and the oxidation reaction started by introducing oxygen (flow rate : 0.4 1.min -1) in the stirred (1000 rpm) slurry. Measurements performed with 50 to 100 mg catalyst at different stirring rates (in the range 10001800 rpm) confirmed the absence of diffusional limitations under these conditions. After 4 hours reaction, the oxygen inlet was turned off and the

520 catalyst was removed from the reaction mixture by filtration. The catalyst was washed with water, ethanol and ether and dried under vacuum at 30~ before being analyzed by XPS, XRD and SEM-EDX. For the recycling experiments, the catalysts were dried under vacuum at room temperature before being reused in a further catalytic test with 400 ml of a fresh 1 mol.1-1 glucose solution. The composition of the reaction mixture was determined by HPLC and 13C-NMR spectroscopy. The bismuth and palladium losses from the catalysts in the reaction mixture during the catalytic tests were determined by analyzing the collected filtrates by atomic absorption spectrometry. Analytical conditions were described elsewhere [8].

2.2.2. Expression of the catalytic results Because the 13C-NMR analyses showed that gluconic acid was the only carboxylic acid generated in the reaction medium, the yields in gluconic acid (YGLU, %) were calculated directly from the NaOH consumption. The main side product is fructose due to isomerisation in the presence of oxygen and appears at an extent between 2.6 and 4.6 % yield when YGLU is larger than 10 %.

2.3. Catalyst characterization techniques 2.3.1. X-ray diffractometry (XRD) Powder X-ray diffraction patterns were obtained with a SIEMENS D-5000 diffractometer using the Ks-radiation of a copper anode. The samples were analyzed after deposition on a quartz monocrystal sample-holder supplied by Siemens. The crystalline phases were identified by reference to the ASTM data files.

2.3.2. X-ray induced photoelectron spectroscopy (XPS) X-ray photoelectron spectroscopy was performed on a SSI-X-probe (SSX100/206) spectrometer from FISONS, using the AI-K~ radiation (E = 1486.6 eV). The energy scale was calibrated by taking the Au 4f7/2 binding energy at 84 eV. The C ls binding energy of contamination carbon fixed at 284.8 eV was used as internal standard value. The analysis of bismuth and palladium were based on the Bi 4f7/2 and Pd 3d5/2 photopeaks. The intensity ratios I(Bi4f7/2)/I(Bi4fs/2) and I(Pd3d5/2)/I(Pd3d3/2) were fixed at 1.33 and 1.5 respectively.

2.3.3. Electron microscopy Scanning electron microscopy was carried out on a LEICA stereoscope $260 equipped for energy dispersive X-ray analysis (EDAX 9100). Samples are deposited on a Cu-AI support from a slurry in acetone.

521

3. RESULTS 3.1. Characterization of the supported bimetallic catalysts The XRD characterization results of this type of catalysts have already been discussed in a previous paper [8]. The main observations consisted in the presence of metallic palladium in the monometallic P d / C catalyst, and the obtention of poorly resolved XRD spectra for the bimetallic catalysts, in which the formation of one of several binary BixPdy alloys was nevertheless suspected. Representative XPS results are listed in table 1 for fresh and used catalysts. Table 1 XPS atomic intensity ratios in fresh and used catalysts Catalyst (a)

P d / C (%)

Bi/Pd (b)

Bi/C (%) (b)

Other data

(Ac.Pd/C)f

0.69

-

-

(Ac.Pd/C)tl

0.65

-

-

(Ac.Pd/C)tl - 5g H5dtpa

0.65

-

-

(Ac.PdBi/C)f

0.72

1.10

0.80

(Ac.PdBi/C)tl

1.12

0.50

0.56

(Ac.PdBi/C)tl - 5g H5dtpa

0.77

0.39

0.30

(Ac. PdBi / C)t5

1.19

0.37

0.44

(Ac.PdBi/C)t13

0.90

0.37

0.33

(PdPtBi / C-Degussa)f

1.06

1.92

2.05

P t / C = 0.24 %

(PdPtBi/C-Degussa)tl

1.30

1.00

1.30

P t / C = 0.12 %

N / C = 0.4 %

N / C = 0.5 %

P t / C = 0.15 % (PdPtBi/C-Degussa)t5 1.02 0.90 0.92 (a) f: fresh catalyst tn: catalyst engaged in n consecutive tests (b) theoretical(bulk) values : Bi/Pd = 0.52 ; Pd/C (xl00) = 0.62 ; Bi/C (xl00) = 0.32 X P S : Bismuth and palladium appear in the metallic (Pd 0, Bi 0) and the oxidized form (Pd 2+, Bi3+). The binding energy values associated with the Bi 4f7/2 photopeak lie in the range 157.7-158.7 eV for Bi0 and 159.1-160.2 eV for Bi3+ and, for the Pd3d5/2 line, in the range 335.6-335.9 for Pd 0 and 337.3-338.0 eV for Pd 2+. The experimental B i / P d ratio in the fresh bimetallic catalyst is higher than the theoretical value calculated from the bulk composition of the catalyst, indicating a partial coverage of palladium by bismuth. This observation is in agreement with the sequential incorporation of Pd first, then Bi, during the preparation of this catalyst, but also with the propensity to segregation in Pd-Bi alloys. The experimental values for the atomic ratios P d / C and Bi/C in the fresh catalysts are higher than the theoretical values, indicating that a d e q u a t e d i s p e r s i o n is achieved for both metals on the surface. The B i / P d molar ratio in the used bimetallic catalysts decreases to reach the value of 0.5 after one run. This value

522 was already observed in previous experiments with bimetallic catalysts of several Bi-Pd compositions [9]. This value decreases slightly further up to about 0.4 after a series of 5 consecutive tests and remains constant afterwards. This decrease in the Bi/Pd but also in the Bi/C ratio after the catalytic tests is in line with the bismuth losses previously observed during the catalytic oxidation of D-glucose in Dgluconic acid. The observation of nitrogen compounds on the surface of the catalysts tested in presence of H5dtpa and the unusually intense lines in the Cls spectral region of those catalysts indicate that this complexing agent is adsorbed on the surface of the catalyst. Electron microscopy : The use of pyrazinedicarboxylate-type precursors was found to generate the largest metal particles with a maximum size of 20 ~tm, whereas the catalysts prepared from acetate precursors are so small that they are not detected by backscattered electron microscopy. EDX analysis of the (pzdc.PdBi/C) catalyst demonstrated the bimetallic nature of the metal particles. The SEM-EDX spectra of the used monometallic Ac.Pd/C catalyst tested in the presence of added dissolved bismuth, according to the procedure described elsewhere [8], gave evidence for bismuth deposition on the support and on Pd particles. 3.2. Catalytic results 3.2.1. Influence of dissolved bismuth Preliminary experiments described elsewhere allowed us to check the absence of any catalytic activity in the presence of the support alone, but also in the presence of the monometallic Bi/C catalyst [8]. The experiments performed in this work (fig. 1) indicate that the catalytic performances of the inactive phase BiPd3 are enhanced in the presence of a monometallic catalyst Ac.Bi(10 wt.% Bi)/C (after 4 h reaction, YGLU (50 mg BiPd3) = 2 %, to compare with YGLU (50 mg BiPd3 + 54 mg Ac.Bi/C) = 9.6 %). The same observation was made when two separate monometallic catalysts Ac.Pd/C and Ac.Bi/C were engaged simultaneously [8]. As 70 to 80 % of the initial bismuth content was found to dissolve from this monometallic catalyst in the reaction m e d i u m during the catalytic tests, the presence of this dissolved bismuth in solution could be responsible for the activity increase. Under conditions that need to be investigated further, the presence of bismuth in solution is therefore a sufficient condition to improve the catalytic activity of a monometallic catalyst P d / C , in disagreement with previous observations [8]. Complementary experiments in the presence of a chelating agent (Hsdtpa) were conducted to improve the understanding of this peculiar point. The gluconic acid yields obtained after 4h and the bismuth losses are listed in table 2. A decrease of the gluconic acid yield was observed when the mass of engaged H sdtpa in the catalytic operations was globally increased. This means that complexation with dtpa induces a decrease in activity because it removes complexation of dissolved Bi3+ by the gluconate ions generated by the catalytic

523 reaction. However, the observed loss of activity could also be assigned to a poisoning effect of the catalyst by the complexing agent itself. This is supported by the deactivation of a monometallic catalyst Ac.Pd/C in the presence of the same complexing agent (YGLU (Ac.Pd/C) = 11.9 %, to compare with YGLU (Ac.Pd/C + 5g H5dtpa) = 3.4 %) and also by the XPS results showing that nitrogen is present on the catalysts after these experiments. Further experiments should therefore be performed in the presence of an insoluble complexing agent able to trap the dissolved bismuth and to remove it physically from the solution, in order to avoid poisoning effects due to adsorption phenomena.

40

~Cy-0 30-

20-

10-

o

.~

,~

.................

,? .. .1. .1 "~

0

b

j

-

9............ ._

!

!

c _

Is Ii

d

_

I

I

I

!

1

2

3

4

time (h)

Figure 1 : Yield in gluconic acid vs reaction time in the presence of a) a mixture of two monometallic catalysts (Ac.Pd/C + Ac.Bi/C) b) a mixture of an inactive intermetallic phase and a monometallic catalyst (BiPd3 and Ac.Bi/C) c) the monometallic catalyst (Ac.Pd/C) and d) the intermetallic phase BiPd3 Table 2 Influence of the mass of H5dtpa engaged on the catalytic performances and Bi losses of the bimetallic catalyst Ac.PdBi/C m H5dtpa (mg)

YGLU (%)

Bi losses (%)

0

45.1

12

50

26.5

18

100

24.6

15

500

20.4

15

5000

10.9

23

524

3.2.2. Recycling Two series of tests were carried out to investigate the deactivation process of the bimetallic catalyst Ac.PdBi/C; 200 and 100 mg of the bimetallic catalyst were respectively engaged in 13 and 5 successive tests. Because the absolute amounts of catalyst engaged in consecutive tests was regularly decreasing, gluconic acid yields were normalized with respect to the catalyst mass. The absolute and normalized gluconic acid yields and the bismuth losses after each catalytic operation are listed in tables 3 and 4. When normalized with respect to the mass of engaged catalyst, the gluconic acid yield remains constant during the successive operations while significant Bi losses are observed only after the first one or two experiments. This is in agreement with the fact that constant Bi/Pd atomic ratios were observed by XPS on the surface of the bimetallic catalyst after 5 to 13 successive experiments (see Table 1). Table 3 Catalytic performances and Bi losses of the bimetallic catalyst Ac.PdBi/C engaged in 13 successive tests (180-200 mg catalyst - 50~ - 1000 rpm - pH 9.25-9.45 - 4h diffusional limitations) Test

YGLU(%)

YGLU/mcat (%.mg -1)

Bi losses (%)

1

89.4

0.45

23 + 3

2

79.7

0.40

<2.5

3

74.8

0.40

<2.5

4

75.5

0.40

<2.5

5 6

74.0 76.7

0.37 0.39

<2.5 <2.5

7

73.3

0.40

<2.5

8

70.0

0.39

<2.5

9

73.2

0.41

<2.5

10

57.8

0.32

<2.5

11

67.4

0.38

<2.5

12

71.1

0.40

<2.5

13

68.9

0.39

<2.5

Similar results were obtained with a commercial trimetallic Pd-Pt-Bi/C catalyst engaged in 5 successive tests; the normalized gluconic acid yield measured after each catalytic test remains constant. Bi losses are less extensive during the first operation but proceed further throughout the first five experiments (table 5). Perhaps the presence of Pt in the catalyst can account for this slightly different behaviour under the reaction conditions.

525 Table 4 Catalytic performances and Bi losses of the bimetallic catalyst Ac.PdBi/C engaged in 5 successive tests (90-100 mg catalyst - 50~ - 1000 rpm - pH 9.25-9.45 - 4h no diffusional limitations) Tests

YGLU(%)

YGLU/mcat (%.mg -1)

Bi losses (%)

1

65.0

0.65

15

2

65.3

0.72

<5

3

65.7

0.70

<5

4

62.0

0.65

<5

5

61.9

0.67

<5

Table 5 Catalytic performances and Bi losses of the commercial trimetallic catalyst Pd(4%)Pt(1%)Bi(5%)/C from Degussa engaged in 13 successive tests (130-140 mg catalyst- 50~ - 1000 r p m - pH 9.25-9.45 -4h) Tests

YGLU(%)

YGLU/mcat(%.mg -1)

Bi losses (%)

1

66.7

0.48

4.0

2

65.0

0.47

6.3

3

66.4

0.50

6.6

4

70.2

0.52

4.6

5

70.9

0.51

<3.5

These results suggest that part of the promoting element initially present in the catalyst is not essential to allow the reaction to proceed and that the useless Bi is removed from the catalyst surface during the first experiments. However, the dissolution process involving a small quantity of the initial Bi might be critical to control the performances of these catalysts by setting a chemical equilibrium between reduced and oxidised Bi species at the catalyst surface.

4. CONCLUSIONS The major conclusions drawn from the present work are" (i) C o m p a r a t i v e experiments carried out in the presence of mixtures of monometallic catalysts or when dissolved Bi is added to a P d / C catalyst suggest that the d i s s o l u t i o n - r e d e p o s i t i o n p h e n o m e n a i n v o l v i n g Bi are critical parameters to understand the nature of the promoting effect in these catalysts.

526 (ii) The experiments carried out in the presence of diethylenetriaminepentaacetate, a stronger complexing agent than the gluconate ion, resulted in a decrease in the gluconic acid yield that could be assigned to the complexation of the dissolved Bi by dtpa but also to the poisoning of the catalyst by these species. (iii) When a bi- or tri- metallic catalyst was engaged in series of 13 or 5 successive tests, the measured gluconic acid yield, normalized with respect to the mass of engaged catalyst, remained constant. Significant Bi losses were observed only after the first experiments, suggesting that most of the dissolved fraction of the initial Bi represents an excess amount of promoting element which is not necessary to monitor the oxidation reaction. However, Bi dissolution might be a critical feature for the operation of these catalysts, by ensuring an appropriate Bi3+/Bi 0 balance at the catalyst surface and by making gluconate desorption from the surface easier.

ACKNOWLEDGEMENTS

The authors greatly acknowledge financial support from the Belgian National Fund for Scientific Research (F.N.R.S.. Brussels) for the programme concerning the selective oxidation of glucose. The authors are grateful to NORIT for supplying the carbon support, to Dr. R. Touillaux and J.F. Statsijns for their assistance in the analytical part of this work, and to the F.R.I.A. Brussels and the Catholic University of Louvain for the fellowships allotted to M.W. REFERENCES

1. T. Mallat and A. Baiker. Catal. Today 19 (1994) 247. 2. C. Br6nimann, Z. Bodnar, P. Hug, T. Mallat and A. Balker, J. Catal. 150 (1994) 199. 3. M. Besson, F. Lahmer, P. Gallezot, P. Fuertes and G. Fl~che, J. Catal. 152 (1995) 116. 4. A. Abbadi and H. van Bekkum, Appl. Catal. A 124 (1995) 409. 5. B.M. Despeyroux, K. Deller and E. Peldszus, Stud. Surf. Sci. Catal. 55 (1990) 159. 6. H.E.J. Hendriks, B.F.M. Kuster and G.B. Marin, Carbohydrate Res. 204 (1990) 121. 7. R. Garcia, M. Besson and P. Gallezot, Appl. Catal. A 127 (1995) 165. 8. M. Wenkin, R. Touillaux, P. Ruiz, B. Delmon and M. Devillers. Appl. Catal. A, 148 (1996) 181. 9. M. Wenkin, C. Renard, P. Ruiz, B. Delmon and M. Devillers, Proceedings of the 4th International Symposium on Heterogeneous Catalysis and Fine Chemicals (Basel, Switzerland, sept. 1996), in press. 10. D. T. Sawyer, Chem. Rev. 69 (1964) 633. 11. V.I. Kornev and A.V. Trubachev, Russ. J. Inorg. Chem. 32 (1987) 1419. 12. M. Wenkin, M. Devillers, B. Tinant and J-P. Declercq, Inorg. Chim. Acta, in press. 13. M. Wenkin, R. Touillaux and M. Devillers, submitted for publication.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

527

O x i d a t i v e d e h y d r o g e n a t i o n of glycolic acid to g l y o x y l i c acid over Fe-P-O catalyst M. Ai a and K. Ohdan b aDepartment of Applied Chemistry and Biotechnology, Niigata Institute of Technology, 1719 Fujihashi, Kashiwazaki 945-11, Japan bUbe Laboratory, UBE Industies Ltd., 1978-5 Kogushi, Ube 755, Japan Glyoxylic acid was found to be produced by a vapor-phase oxidative dehydrogenation of glycolic acid over iron phosphate catalysts with a P/Fe atomic ratio of 1.2. The best results were obtained with iron phosphates freshly calcined at 400 to 450~ Reduced iron phosphates showed a markedly lower activity. The optimum reaction temperature was about 240~ The selectivity to glyoxylic acid was 74 tool% up to the glycolic acid conversion of about 70%; the highest yield of glyoxylic acid was 56.5 mol% at the conversion of 80 %.

1. INTRODUCTION Glyoxylic acid is a raw material for various chemicals. It is generally produced by enzymatic or nitric acid oxidation of glyoxal, or electrolytic reduction of oxalic acid. It is also known that alkyl esters of glyoxylic acid are obtained by a vaporphase oxidation of corresponding alkyl esters of glycolic acid [1]. The yield of ester reached 69 mol% at the conversion of 94%. HOCH2-COOR + 0.5 02

9 OHC-COOR + H20

However, it seemed still difficult to obtain glyoxylic acid directly from glycolic acid in a vapor-phase process. In the preceding studies [2-4], it was found that iron phosphate catalysts show a high selectivity in the oxidative dehydrogenation of compounds in which the carbon atom at the a-position of an electron-attracting group (X) such as -COOH, -CHO, or CN, is tertiary; for example, isobutyric acid, isobutyraldehyde, and isobutyronitrile, but that they are inactive for oxygen insertion reactions. R-CH2-CHR'-X + 0.5 02

9 R-CH=CR'-X + H20

528 It was also found recently that iron phosphate catalysts show a high selectivity both in the oxidative dehydrogenation of lactic acid to pyruvic acid and in the oxidative decarboxy-condensation of pyruvic acid to citraconic anhydride [5-7]. CH3-CH(OH)-COOH + 0.5 02 ; CH3-CO-COOH + H20 2 CH3-CO-COOH + 0.502 ; C5H403 + CO2 + 2 H 2 0 These findings led us to study the catalytic performance of iron phosphate in the oxidative dehydrogenation of glycolic acid to glyoxylic acid. HOCH2-COOH + 0.5 02

, OHC-COOH + H20

2. EXPERIMENTAL

2.1. Catalysts The catalysts used in this study were iron phosphates with a P / F e atomic ratio of 1.2, prepared according to the procedures described in the previous studies [2,4,8]. On the other hand, a V-P oxide catalyst with a P / V atomic ratio of 1.06 consisting of vanadyl pyrophosphate [(VO)2P207] was prepared according to a patent [9]. H3PMo12040 and HsPMol0V 2040 catalysts were prepared by supporting molybdophosphoric acid and molybdovanadophosphoric acid, respectively, on an equal weight of natural pumice between 10 to 20 in mesh size [10].

2.1. Reaction procedures The reaction was carried out with a continuous-flow system at atmospheric pressure. The reactor was made of a stainless steel tube, 50 cm long and 1.8 cm I.D., mounted vertically and immersed in a lead bath. The catalyst was placed near the bottom of the reactor and porcelain cylinders, 3 m m long and 1.5 m m I.D./3.0 m m O.D., were placed both under and above the catalyst bed. Air or a mixture of nitrogen and oxygen was fed in from the top of the reactor; an aqueous solution containing 100 g of glycolic acid in I 000 ml was introduced into the preheating section of the reactor with a syringe pump. Unless indicated otherwise, the feed rates of glycolic acid, oxygen, water, and nitrogen were 12.26, 10.0, 480, and 500 m m o l / h , respectively, and the reaction temperature was 240~ The amount of catalyst used was 10 g; the contact time was about 2.5 s. The extent of reaction was varied by changing the amount of catalyst used from 0.7 to 20 g, while fixing the feed rates. The effluent gas from the reactor was led successively into four chilled scrubbers to recover the water soluble compounds. The products were analyzed by GC and LC. The definitions of contact time, conversion, yield, and selectivity were the same as those in a previous study [11]. 3. RESULTS 3.1. Performance of metal phosphate catalysts Since the reaction is an "acid to acid type" transformation, it is predictable that the effective catalysts should be acidic [12]. The V-P oxide and heteropoly

529 compound catalysts, which show a good performance in oxidation producing acidic compounds such as maleic anhydride and methacrylic acid, were tested for the oxidative dehydrogenation of glycolic acid together with an iron phosphate catalyst. The results are summarized in Table 1. It is clear that the iron phosphate catalyst is much more effective than the V-P oxide and heteropoly compound catalysts. Table 1 Performance of metal phosphate catalysts Catalyst Temp Fe-P

V-P Mo-P Mo-V-P

02 feed

Conv

oc

mmol/h

%

GXAD

HCHO

Yield (mol%) HCCK)H

COx

mol%

S

240 240 240 240 265 240 250 240 270

11.3 12.5 16.2 7.5 7.5 11.0 11.0 10.0 10.0

53.0 59.0 70.4 54.4 93.7 66.0 83.7 70.5 86.9

40.0 45.0 53.8 15.1 16.0 16.6 15.5 19.5 24.4

3.9 4.0 4.1 10.0 30.0 9.1 10.6 2.5 11.2

0.7 0.9 1.6 2.0 4.5 10.0 10.9 5.7 11.0'

7.4 6.7 10.7 15.2 45.1 32.4 38.0 15.2 42.1

77 76 76 28 18 25 19 27 28

Conv: conversion of glycolic acid, GXAD: glyoxylic acid, S: selectivity to glyoxylic acid, Fe-P: iron phosphate, V-P: vanadyl pyrophosphate, Mo-P: H3PMo12040, Mo-V-P: H5PMol0V 2040.

3.2. Effect of calcination temperature on the catalytic performance It was found in previous studies [13,14] that when the temperature of calcination of freshly prepared iron phiosphate is raised, the surface area decreases markedly and the structure changes clearly, that is, amorphous phase FePO4 is transformed into tridymite type FePO4 at 400 to 500~ and the tridymite type FePO4 is then transformed into quartz type FePO4 at a temperature above 500 to 550~ The effects of the calcination temperature on the catalytic performance were studied. The results are summarized in Table 2. When the temperature of calcination was low; 300~ the catalyst consisted of amorphous FePO4. It showed a high activity for consumption of glycolic acid, but the main products were COx and formaldehyde; the selectivity to glyoxylic acid was low. On the other hand, when the temperature of calcination was high; 500~ the catalyst consisted of a mixture of tridymite and quartz FePO4. The catalyst showed a low activity, though the selectivity to glyoxylic acid was not bad. As a result, it is concluded that the optimum temperature of calcination of iron phosphate is in the range of 400 to 450~

530 Table 2 Effect of calcination temperature on the catalytic performance of iron phosphate Calcin Temp ~ 300 400

450

500

Reaction Temp c.t. ~ s 240 0.6 230 1.2 240 1.2 240 2.5 240 4.0 240 1.2 240 2.5 240 3.0 240 4.0 240 5.0 240 2.5 240 5.0

GAD Conv % 42.0 63.5 44.5 66.0 84.0 36.3 62.0 69.5 76.4 85.0 11.5 45.0

GXAD 13.8 9.2 31.0 46.2 55.8 26.8 45.0 51.2 55.0 56.3 6.7 31.8

Yield (mol%) HCHO HCOOH 8.2 14.3 3.2 6.0 6.0 2.1 4.0 4.0 5.4 5.4 1.8 3.2

1.6 2.3 0.7 3.0 5.5 0.8 1.9 2.9 3.2 6.3 0.0 0.5

COx 14.4 26.0 4.6 12.4 17.3 2.8 9.3 10.1 13.6 18.6 0.8 4.8

Select GXAD tool% 33 14 70 70 67 74 73 74 72 67 58 70

Calci: calcination of iron phosphate, Select: selectivity, c.t.- contact time, Cony" conversion, GAD: glycolic acid, GXAD: glyoxylic acid. 3.3. Effect of reduction of iron phosphate on the catalytic performance It is known from a previous study [13] that fresh iron phosphate catalysts consising of FePO4 (Fe3+), in which amorphous, tridymite type, and quartz type are included, are reduced to Fe2P207 (Fe 2+) via an intermediate compound consisting of Fe3(P207)2 (Fe233+) 9 In the oxidative dehydrogenation of lactic acid to pyruvic acid, the catalysts consisting of Fe3(P2C)7)2 phase show a higher selectivity than the FePO4 and Fe2P207 catalysts [5]. While in the oxidative dehydrogenation of isobutyric acid to methacrylic acid, the catalytic performance; activity and selectivity, were almost independent of the variation in both the structure and the degree of reduction [15]. These findings led us to check the effect of reduction of iron phosphate on the catalytic performance. The reaction was performed using three catalysts; an iron phosphate freshly calcined at 450~ the one which was partially reduced, and the one which was fully reduced by hydrogen. The results are summarized in Table 3. The partially and fully reduced iron phosphates are clearly lower in both the catalytic activity and selectivity than the iron phosphate consisting of tridymite FePO4.

3.4. Performance of fresh (non-reduced) iron phosphate catalyst In order to clarify in more detail the characteristic features of the oxidative dehydrogenation of glycolic acid to glyoxylic acid, the study was performed using a iron phosphatre catalyst freshly calcinbed at 450~

531 Table 3 Effect of oxidation states and structure on the catalytic p e r f o r m a n c e Catalyst structure

Reaction c.t. Temp s ~

[Fresh] FePO4 tridymite

2.5 2.5 4.0

Fe3(P207)2

Fe2P207

Yield (mol%) HCHO HCOOH

GXAD

240 250 240

62.0 77.0 76.4

45.0 54.2 55.0

4.0 6.1 5.4

1.9 3.9 3.2

9.3 15.8 13.6

73 70 72

2.5 2.5 4.0 4.0

240 250 240 250

33.0 47.5 48.3 72.6

25.0 32.7 32.3 45.0

3.3 4.4 4.8 7.8

1.0 1.4 0.9 3.8

3.5 6.9 6.3 18.2

75 68 67 62

2.5 2.5

240 260

35.2 70.0

25.0 40.0

3.7 11.0

0.7 2.1

5.0 17.6

71 57

COx

A b b r e v i a t i o n s are the s a m e as those in Tables 1 a n d 2

3.4.1. Product distribution 60

50

~.

40

3o

20 o

10

Select GXAD mol%

GAD Conv ~

|/I

f

0 L/ 0

9-" e ~~ , 20 30 40 50 60 10 C o n v e r s i o n of glycolic acid / %

Figure 1. P r o d u c t distribution

, 70

HCQOH 80

90

532 ~-

601 300

~ 50 40

~- 40 30

240 ~C

280oc

~ 30

0~

"~ 20 ~ 10

" i

0

0

"

", 0.2

,

,

,

,

,

0.4 0.6 0.8 1.0 1.2 Contact time / s

Figure 2. Effect of temperature on the rate

1020

, . . . . 40 50 60 70 80 9 0 Conversion of glycolic acid / % ,

30

Figure 3. Effect of temperature on the yield of glyoxylic acid

The reaction was performed under the fixed reaction conditions described in the Experimental section, while changing the contact time from 0.3 to 5.0 s. The main products detected were glyoxylic acid, formaldehyde, formic acid, and carbon oxides. The yields of these products are plotted as a function of the conversion of glycolic acid in Figure 1. The slope of lines from the origin indicate the selectivities to each product. The selectivity to glyoxylic acid remains unchanged at about 74 mol% with an increase in the conversion of glycolic acid up to about 70%. With a further incrase in the conversion, the yield of glyoxylic acid levels off; the m a x i m u m yield is 56.5 tool% at the conversion of 80%. Glyoxylic acid is relatively stable under the reaction conditions used. The side reactions are the formation of formaldehyde and formic acid, that is, C-C bond fission by oxygen insertion. 3.4.2.

Effect

of reaction

temperature

The reaction was performed by changing the reaction temperature from 200 to 300~ and the contact time from 0.3 to 5 s, while fixing the feed rates as described in the Experimental section. Figure 2 shows the conversion of glycolic acid as a function of the temperature. The results indicate that the rate of reaction becomes about double as the temperature is raised by 20~ The yield of glyoxylic acid obtained at different temperatures are plotted as a function of the conversion of glycolic acid in Figure 3. As the temperature is raised, the selectivity to glyoxylic acid decreases, especially at higher conversion regions. However, it should be noted that vaporization of glycolic acid becomes difficult at tempertures below 230~ because the boiling point of glycolic acid is high. It is therefore concluded that the optimum temperature is about 240~ under the reaction conditions used.

533 ~60 m

C

40

-----O'-

~ t~

U

~ ~,,,i

;>

O U

@ =

o;> 30 -

S

Feed of glycolic acid - 12.2 retool/h

olO

o

4:0

~

30

,- 20

=

E 5o

0

t

t3

I

L

I

I

//,

o 20 i

5 10 15 20 25 100 Oxygen feed rate / m m o l / h

Figure 4. Effect of oxygen on the rate

10

~, ~ a

Oxygen feed rate o 10 m m o l / h

-

a

a~ I

25 m m o l / h 9 100 m m o l o / h

I

I

I

l

l

I

20 30 40 50 60 70 80 Conversion of glycolic acid /% Figure 5. Effect of oxygen on the yield of glyoxlic acid

3.4.3. Effect of oxygen feed rate The reaction was performed by changing the feed rate of oxygen from 1.5 to 100 r e t o o l / h and the a m o u n t of catalyst used, while fixing the temperature and the feed rates of glycolic acid, water and nitrogen as described in the Experimental section. The conversions of glycolic acid obtained at a short contact time of 0.62 s (amount of catalyst used = 2.5 g) are plotted as a function of the feed rate of oxygen in Figure 4. The rate of reaction increases with an increase in the feed rate of oxygen, but the oxgen dependency is very low; far from first order dependency. Figure 5 shows the yields of glyoxylic acid obtained with different oxygen feed rates and at different contact times as a function of the conversion of gylcolic acid. W h e n the c o n v e r s i o n is not high; less than 60 %, the selectivity to glyoxylic acid is almost independent of the oxygen feed rate. However, with a further increase in the conversion, the selectivity decreases with an increase in the feed rate of oxygen. It is therefore concluded that the o p t i m u m oxygen feed rate is in the range of 10 to 25 m m o l / h under the reaction conditions used. 4. DISCUSSION Three reactions of glycolic acid may take place in parallel as follows: HOCH2-COOH + 0.5 02 > OHC-COOH + H20 HOCH2-COOH + 02 > HCHO + CO2 + H 2 0 HOCH2-COOH + 02 > HCOOH + CO + H 2 0 Mo and V based oxides and phosphates possess both acidic and redox functions. Therefore, they show a good performance as catalysts in many partial oxidations for producing especially acidic compounds [12]. However, they possess

534 double bond oxygen species, that is, M = O species, which promote the oxygen insertion reactions as well as the dehydrogenation. As a result, in the reaction of glycolic acid, they promote the C-C bond fission by oxygen insertion. On the other hand, iron phosphate possesses both acidic and redox functions as like Mo and V oxides and phosphates. However, it possesses no double bonded oxygen species unlike the Mo and V compounds. Therefore, its function to promote oxygen insertion processes is very weak. As a result, in the reaction of glycolic acid, the formation of formaldehyde and formic acid is suppressed, to a certain extent, and a relatively good performance for glyoxylic acid is obtained. As is seen in Figures 4 and 5, use of a low oxygen feed rate seems to be beneficial. However, it should be noted that under the reaction conditions described in the Experimental section, the catalytic activity was stable enough during the first 1 or 2 days, but after then the activity decreased slowly. The deactivated catalyst was found to be reduced during the reaction. As shown in Table 3, the catalytic activity for the reaction of glycolic acid is markedly decreased by the reduction of iron ions. As is seen in Figure 3, use of a low temperature of 240~ is beneficial to the selectivity to glyoxylic acid. The iron phosphate catalyst shows a relatively high oxidation activity even at a low temperature of 240~ This suggests that the redox cycles on the surface takes place rapidly. However, the rate of re-oxidation of reduced bulk iron phosphate should be very slow at a low temperture of 240~ It is therefore necessary to avoid the reduction of bulk in order to keep the activity. It is also noted that the deactivated catalyst was fully regenerated by heattreatment at about 450~ in stream of air. REFERENCES

1. P.R. Anantaneni and T.P. Li, US Patent No. 4 900 864 (1990). 2. M. Ai, E. Muneyama, A. Kunishige, and K. Ohdan, Bull. Chem. Soc. Jpn., 67 (1994) 551. 3. M. Ai, E. Muneyama, A. Kunishige, and K. Ohdan, Catal. Lett., 24 (1994) 362. 4. E. Muneyama, A. Kunishige, K. Ohdan, and M. Ai, J. Mol. Catal., 89 (1994) 371. 5. M. Ai and K. Ohdan, Chem. Lett., (1995) 405. 6. M. Ai and K. Ohdan, Chem. Lett., (1996) 247' 7. M. Ai and K. Ohdan, Stud. Surf. Sci. Catal., 101 (1996) 201. 8. E. Muneyama, A. Kunishige, K. Ohdan, and M. Ai, Appl. Catal., A. 116 (1994) 165. 9. M. Katsumoto and D.M. Marquis, US Patent No. 4 132 670 (1979). 10. M. Ai, J. Catal., 116 (1989) 23. 11. M. Ai, E. Muneyama, A. Kunishige, and K. Ohdan, Appl. Catal., A. 109 (1994) 135. 12. M. Ai, Proceedings, 7th International Congress on Catalysis, Tokyo, 1980, Kodansha, Tokyo - Elsevier, Amsterdam, 1981, p. 1060. 13. E. Muneyama, A. Kunishige, K. Ohdan, and M. Ai, J. Catal., 158 (1996) 378. 14. M. Ai, E. Muneyama, A. Kunishige, and K. Ohdan, J. Catal., 144 (1993) 632. 15. M. Ai, E. Muneyama, A. Kunishige, and K. Ohdan, Catal. Lett., 24 (1994) 355.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

535

S h a p e Selective E p o x i d a t i o n of Crotyl Alcohol with H20 2 in the Presence of TS-1 Graham J. Hutchings', Paul G. Firth', Darren F. Lee', Paul McMorn a, Donald Bethell', Philip C. Bulman Page b, Frank King c and Frederick Hancock c 'Leverhulme Centre for Innovative Catalysis, Department of Chemistry, University of Liverpool, Liverpool, L69 3BX, United Kingdom. bDepartment of Chemistry, Loughborough University, Loughborough, Leicestershire, LE 11 3TU, United Kingdom. clCI Katalco, PO Box 1, Billingham, TS23 1LB, United Kingdom.

The epoxidation of C 3 and C4 allylic alcohols with hydrogen peroxide has been investigated using the microporous titanium silicalite TS-1. The reactivity of these substrates is found to decrease in the order cis-crotyl alcohol > trans-crotyl alcohol > allyl alcohol >> 2methylallyl alcohol, the latter showing very little reaction under the conditions investigated. The use of alcohols as solvents was investigated for allyl and cis- and trans-crotyl alcohols; methanol was found to be the best solvent for allyl alcohol and ethanol for crotyl alcohol. However, the reaction was possible using water as solvent without loss of selectivity of the oxirane due to triol formation at short reaction times (1-3 h). A non-catalysed homogeneous oxidation, using meta-chloroperbenzoic acid as oxidant, showed that the relative reactivity of cis-crotyl alcohol compared to the trans-crotyl alcohol was significantly decreased, and hence the higher reactivity of cis-crotyl alcohol using TS-1 indicates that the oxidation reaction is shape selective within the microporous environment. The results are discussed with respect to the nature of the reaction mechanism for the formation of the oxirane product.

1. I N T R O D U C T I O N The heterogeneous epoxidation of compounds which contain carbon-carbon double bonds is an important industrial process in both the manufacture of fine chemicals and in the synthesis of natural products. A number of studies have demonstrated that alkenes can be readily epoxidised by hydrogen peroxide using the titanium silicalite TS-1 [ 1-31. However, it has been found that substitution of the alkene by electron withdrawing groups significantly decreases the reactivity of the carbon-carbon double bond since the decrease of the electron

536 density renders it less susceptible to electrophilic attack by the oxidant. This was recently illustrated by Clerici and lngallina [4] when it was shown that the rate of epoxidation of allyl chloride and allyl alcohol were ca. 14 and 30 times slower respectively than the rate of 1-butene epoxidation. It is therefore of interest to determine how the epoxidation of allylic alcohols can be achieved. To date, TS-1 is the only heterogeneous epoxidation catalyst that utilises hydrogen peroxide as oxidant, an important environmental factor as the only waste products are water and molecular oxygen. In most studies the reaction is carried out in an alcohol solvent, e.g. methanol, and it has been suggested that the alcohol plays an important part in the formation of the active site [41. Although TS-I has been investigated for the epoxidation of a range of molecules, e.g. butene, pentene, hexene, allyl chloride and allyl alcohol, little attention has been given to the effect of shape selectivity in the MFI zeotype framework in these reactions. In this paper we address this aspect and exemplify the shape selective epoxidation using a range of allylic alcohols. In particular, the shape selective epoxidation of crotyl alcohol is compared and contrasted with the reaction of allyl alcohol in a range of solvents.

2. EXPERIMENTAL 2.1 Preparation of TS-I Samples of TS- 1, with SiFFi ratio of 32, were prepared by a variation [5] of the original method of Taramasso e t al. [6] using tetraethyl orthotitanate (TEOTi), tetraethyl ortho silicate (TEOSi) and water as reagents and tetrapropyl ammonium hydroxide (TPAOH) as the template. TEOTi (0.75g) was added with stirring to TEOSi (22.5g) and the resulting yellow solution was aged at ambient temperature for 2 h. The aged solution was added slowly to TPAOH (50 ml, Sigma, 20% aqueous solution) with vigorous stirring at 60 ~ the addition being carried out in 1 h to form a clear gel which was then aged at ambient temperature for 24 h. The aged gel was placed in a teflon lined stainless steel autoclave and maintained at 175 ~ for 24 h to enable the product to crystallise. The product was recovered by filtration, washed with water, dried and calcined (6 h, 550 ~ The TS-1 prepared was highly crystalline as determined by X-ray diffraction and electron microscopy and comprised orthorhombic crystallites c a 0.3~m in size. 2.2 Catalyst Testing The epoxidation reaction was c~u'ried out in a flask fitted with a stirrer, thermometer, a reflux condenser and a septum to enable samples to be withdrawn for analysis. In a typical experiment TS-1 (0.0375g) was stirred at a constant temperature for 24h with the allylic alcohol (0.01mol), hydrogen peroxide (0.011 mol, 30%) in the solvent (4 ml). The course of the reaction was monitored by gc and the final reaction products were analysed by gcms and I3C nmr spectroscopy. The allyl alcohols were obtained commercially (Sigma Chemicals) and

537 distilled prior to use, the crotyl alcohol was obtained as a mixture of the cis- and trans- isomers (trans:cis = 19.2:1,4.45:1 and 1:1).

3. R E S U L T S

AND

DISCUSSION

3.1 Comparison of the reactivity of ailylic alcohols The reaction of allyl alcohol, 2-methyl allyl alcohol and crotyl alcohol (trans:cis= 4.45:1) were reacted with hydrogen peroxide at 30 ~ using methanol assolvent, and the results are shown in Figures 1-3. It is apparent that the order of reactivity is cis-crotyl alcohol > transcrotyl alcohol > allyl alcohol > 2 methylallyl alcohol. The latter alcohol was not particularly reactive and this indicates that substition at the 2-position inhibits the reaction significantly. 100

80

u 60

trans-crotyl conv.

9 cis-crotyl conv. a

allyl alc. conv.

A methylallyl alc. conv. 40

20

t

0 ~ 0

I

9

10

I

20

,

30

Reaction Time/h Figure 1 Comparison of the reactivity of allylic alcohols with hydrogen peroxide as oxidant at 30 ~ in methanol solvent using TS-1 as catalyst

538 120 100 80 [] 9 [] ,~

60

conversion glycidol selectivity glycerol selectivity ether diol selectivity

40 20

0

10

20 Reaction Time/h

30

Figure 2 Product selectivity as a function of time for the reaction of allyl alcohol at 30 ~ in methanol solvent. 100 "r-"--'~. 80

~' tr-crotyl conversion 9 cis-crotyl conversion

60

[]

~' cis-oxirane selectivity

40

ether-diol selectivity o triol selectivity

20 0-~/ 0

tr-oxirane selectivity

10

20

30

Reaction Time/h Figure 3 Product selectivity as a function of time for the reaction of crotyl alcohol (cis/trans ratio = 4.45) in methanol solvent at 30 ~

All the allylic alcohols fom~ the oxirane as the initial product, but as the conversion increases the ring opened products, the ether diols (formed by reaction with the alcohol solvent) and the triol (formed by reaction with water) are also observed. It is apparent that cis-crotyl alcohol reacts approximately three times more rapidly than the trans-crotyl even though it is present as the minor component. This was investigated further using crotyl alcohols with a range of trans/cis ratios and the results are shown in Table 1.

Table 1 Reaction of cis- and trans-crotyl alcohol in methanol solvent at 30 "Cusing hydrogen peroxide as oxidant.

Reactant translcis

reaction time 1h

conversion / %

selectivity 1 Ti

trans-crotyl alcohol

cis-crotyl alcohol

trans-oxirane

cis-oxirane

ether diols

trio1

3 24

32.8 61.9 89.0

83.3 95.8 97.9

86.0 89.0 74.4

14.0 9.4 3.4

0 1.6 11.4

0 0 10.8

4.45

1 4 24

28.3 47.1 68.8

76.3 95.4 99.3

58.4 68.2 69.6

40.9 29.6 18.6

0.5 1.9 6.9

0.2 0.3 4.9

1

1 5 24

18.7 42.4 61.3

54.4 86.0 979

25.7 21.9 8.7

74.3 40.9 3.0

0 37.2 66.1

0 0 22.2

1

21.6 48.8

34.2 68.1

75.9 21.8

24.1 11.2

0 0

0 66.9

19.2

4.45l

1

3

Meta-chloroperbenzoic acid as oxidant under the same reaction contions i n the absence of TS- 1.

VI W

a

540 The enhanced reactivity of the cis-crotyl alcohol is observed for all three reactant mixtures. It is also apparent that the cis-disubstituted oxirane is more reactive for the sequential acid catalysed ring opening reactions to form either ether diols or the triol. The relative reactivity of the cis-oxirane when compared to the trans-oxirane is not unexpected since this results from the eclipsing strain of the ring substituents in the cis-oxirane. However, the enhanced reactivity of the cis-crotyl alcohol is unexpected. To confirm this effect was a result of the microporous nature of TS-1, a non-catalysed homogeneous oxidation of crotyl alcohol was investigated using meta-chloro-perbenzoic acid as oxidant under the same reaction conditions (Table 1). In this case the relative rate of oxidation of the cis-crotyl alcohol is not so pronounced as observed for TS-1, although the rate of cis-crotyl alcohol oxidation is higher than that of trans-crotyl alcohol. We therefore conclude that the enhanced rate of oxidation of the cis-crotyl alcohol, observed when TS-1 is used as catalyst, is a result of shape selectivity within the micropores. When TS-1 was used as the catalyst the selectivity to the triol product was very low and only became significant at extended reaction times. When the homogeneous stoichiometric oxidant meta-chloro-perbenzoic acid was used it is apparent that the selectivity to the triol was significant even at short reaction times. The low triol selectivity observed with TS-1 is considered to be due the hydrophobic nature of the titanium silicate framework. 120

100

80

:~

60

40

f

20

l

j

j

j f l r9 . . . . . . . - - "

0

9

,

,

,

10

20

30

Reaction Time/h Figure 4. Oxidation of allyl alcohol in water at 65 ~ glycerol selectivity. ~

GI conversion; 4, glycidol selectivity; II

541 3.2 The effect of solvents on the oxidation of allylic alcohols The oxidation of allyl alcohol and crotyl alcohol (trans-/cis- = 19.2) using hydrogen peroxide as oxidant and TS-1 as catalyst was investigated further using water, methanol and ethanol as solvents and the results are shown in Figures 4 - 9. The observation in the

preceeding experiments that the triol selectivity was not significant when TS-1 was used as catalyst prompted us to investigate the use of water as solvent. It is apparent that the use of water does not lead to a significant loss of selectivity since with both allyl alcohol (Figure 4) and crotyl alcohol (Figure 7) selectivity to the oxirane of >90% could be achieved at c a . 30% conversion, although the selectivity to the triol was more marked with crotyl alcohol as substrate due to the enhanced reactivity of the derived oxirane.

12~t 100 !

80

0

-J

f

v

40

,,.~ . , . . . / "

j 20

0 "~ 0

I,,.,/, -'j'''

10

20

30

Reaction Time/h Figure 5. Oxidation of allyl alcohol in methanol at 65 "C: ~l conversion; 9 glycidol selectivity; A ether diols; II glycerol selectivity.

542

12~t 100

80

,,,.,""~"""~"""~

60 40 20 0 0

10

20

30

Reaction Time/h Figure 6. Oxidation of allyl alcohol in ethanol at 65 ~ A ether diols; II glycerol selectivity.

O conversion; 9 glycidol selectivity;

120 100 80 ~:

60 40 20

~~......~

0 0

10

20

30

Reaction Time/h Figure 7. Oxidation of crotyl alcohol in water at 30 ~ (3 conversion; 9 oxirane selectivity; II triol selectivity.

543

120 100 80 60 40 20

0

10

20

30

Reaction Time/h Figure 8. Oxidation of crotyl alcohol in methanol at 30 ~ Q conversion; t oxirane selectivity; A ether diols; II triol selectivity. 120 100

v v

v v

9

80 60 40 20

.......................

0

i ~ " ............... " ' ' "

. . . . . . ~. . . .

10

i

20

30

Reaction Time/h Figure 9. Oxidation of crotyl alcohol in ethanol at 30 ~ A ether diols; II triol selectivity.

D conversion; 9 oxirane selectivity;

544 3.3 Comments on the reaction mechanism

There are two features that require comment. First, the enhanced reactivity of cis-crotyl alcohol relative to trans-crotyl alcohol indicates that this isomer binds preferentially to the active site in TS-1. The very low reactivity observed with 2-methyl allyl alcohol confirms that the configuration around the carbon-carbon double bond is of prime importance with respect to reactivity. Secondly, the low selectivity to the triol at short reaction times, even when water was used as solvent, is unusual since water (either as solvent or added with H202) could be expected to be reactive towards the oxirane ring opening reaction, and indeed this is observed when meta-chloro-perbenzoic acid was used as oxidant. The observation that water can be used successfully as solvent is important in two respects. It is significant with respect to environmental factors since the use of organic solvents can be avoided. In addition, it is of significance with respect to the reaction mechanism since it indicates that the alcohol solvent does not form an integral part of the active site in the TS-1 catalyst. These observations therefore indicate that the mechanism proposed by Clerici and Ingallina [4], in which an alcohol molecule forms a complex with I-t202 at the active site of TS- 1, cannot be a valid represention. We consider that the role of the alcohol is simply to decrease the water concentration within TS1 and thereby suppress the formation of the non-desired triol at the expense of ether diol formation.

ACKNOWLEDG EMENTS

We thank the DTI, EPSRC, ICI Katalco and Robinson Brothers for financial support.

REFERENCES

M.G. Clerici, G. Bellussi and U. Romano, J. Catal., 129(1991) 159. G. Bellussi, A. Carati, M.G. Clerici and A. Esposito, Stud. Surf. Sci. Catal.,

63

(1991) 421. 3

U. Romano, A. Esposito, F. Maspero, C. Neri and M.G. Clerici, Stud. Surf. Ca tal., 55 (1990) 33. M.G. Clerici and P. Ingallina, J. Catal.,140 (1993) 71. G.J. Hutchings, D.F. Lee and A.R. Minihan, Catal. Lett. 33 (1995) 369. M. Taramasso, G. Perego and B. Notari, US Patent 4410501 (1983).

Sci.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

545

Epoxidation of tertiary allylic alcohols and subsequent isomerization of tertiary epoxy-alcohols : a comparison of some catalytic systems for demanding ketonization processes J.-M. Br6geaulff'*, C. Lepetit b, F. Ziani-Derdar a, O. Mohammedi c, L. Salles~ and A. Deloffre ~ a Laboratoire de Catalyse Homogene et de Chimie des Surfaces, URA 1428 CNRS, Universit6 Pierre et Marie Curie, Tour 54-55, case 196, 4, Place Jussieu, F-75252 Paris cedex 05, France e-mail : [email protected], fr b Laboratoire de Rractivit6 de Surface, URA 1106 CNRS, Universit6 Pierre et Marie Curie, Tour 54-55, case 178, 4, Place Jussieu, F-75252 Paris cedex 05, France. c Institut de Chimie Industrielle, Universite des Sciences et Techniques de Blida, Boulevard E1Aich, B.P. 270-Blida, Algeria *corresponding author

ABSTRACT The single-step ketonization of multi-functional olefinic substrates by 02, H202, R-OOH, quinones, catalysed by group 8 (Rh, Ir) or group 9 (Pd, Pt) transition metal complexes is often considered as a demanding reaction compared to that of simple alkenes. A two-step reaction has therefore been studied: first, alkene epoxidation and, secondly, an efficient way of isomerizing the epoxy-alcohol to the keto-alcohol. Alkene epoxidation by anhydrous t-BuOOH is catalysed by vanadium oxoalkoxide, [OV(OC3H7)3], in benzene or toluene; on a large scale, this system is safer and more advantageous than peracetic acid. Three other catalytic systems can compete in terms of yields and turnover numbers: i) a biphase system with CH3ReO3(MTO)/H202-H20/CH2CI2; ii) a homolytic system with dioxygen in the presence of transition metal-based precursors or even with finely divided inorganic solids (silica, TS-1, KTiOPO4, etc.) and of a coreducer (Me2CHCHO). Phase transfer catalysis systems involving with anionic tungsten species onium salts, e.g. "H2WO4"/H202-H20/H3PO4/Arquad| are also effective, but optimization is difficult. The isomerization of epoxides to carbonyl compounds, in good to excellent yields, can be catalysed by LiBr/HMPA (1:1), cobalt carbonyl complexes or nickel complexes. EPR studies of the nickel-based system and silica-supported Ni(I) complexes are consistent with Ni(I)-based active species. The two-step reaction path gives the best results (yields > 95 %) with several substrates: 2-methyl-3-buten-2-01, 3-methyl1-penten-3-01, 1-vinyl-cyclohexanol. 1. INTRODUCTION Oxidation of olefinic substrates and, more specifically, of ot-olefins to the corresponding ketones is more difficult to perform on functionalized alkenes (e.g. 2-methyl-3-buten-2-ol, la) but is industrially important. These processes can be used to synthesize fine chemicals or

546 organic intermediates such as 2a (2-methyl-butane-2-ol-3-one) [ 1] or other ct-hydroxyketones which are useful intermediates for the synthesis of various natural products including biologically active compounds, sugars and 13-hydroxy-ot-amino acids. As either the yields [2,3], the kinetics [4] or the procedures [5] involved in the earlier syntheses left much to be desired, the development of improved procedures to prepare 2a and parent compounds has been undertaken. Catalytic ketonization of la with dioxygen (or quinone) and a palladium precursor (PdCI2 or [PdCI(NO2)(CH3CN)2]) under mild conditions (room temperature or 3 5°C) may give good conversiom (ca. 98%) and good selectivities(ca. 90°,4),but the turnover numbers are low [6]. Two-step transformations involving an intermediate epoxide show promise. For example, epoxidation may be performed with peracetic acid alone and more efficiently with "vanadium complex/tert-butyl hydroperoxide" combinations. The results obtained with these two systems are compal:ed with those for other catalytic systems. An efficient way of isomerizing the epoxy-alcohol has therefore to be found in order to achieve a two-step catalytic synthesis of the keto-alcohol (Scheme 1). Ring-opening reactions to yield aldehydes or ketones were reviewed [7-10]. Isomerization may be thermally induced [7] or may occur in the presence of basic or acidic agents [11]. Lithium bromide associated with tributylphosphine oxide or hexamethylphosphoric triamide (HMPA) has been used [ 12,13], but transition metal complexes may be more attractive [ 1421]. In this work the performances of catalytic systems, e.g. "LiBr/HMPA/toluene", "Co2(CO)dMeOH", "NiBr2(PPh3)J Zn/PPh#THF", etc. are compared for the isomerization of 3a and of analogues. Supported catalysts have also been studied. We have shown that the second reaction path, 1 -~ 3 -~ 2, gives the best results (yields > 90% and good turnover numbers) with several substrates. O

~l

a

OH

directketonization lip [M"] + "0"

..

""". epoxidation

~~/~ 2a

.¢,

0 "-

[M] + "0"

OH

. ""

",1~

. . ' " isomerization [M'I + "o"

Scheme 1 ([M], [M'] and [M"] : metal salts or transition metal complexes). 2. RESULTS AND DISCUSSION

2.1. Epoxilatim of ~ 2 ~ la, and M" otimr' tmiary alyiic ~ wilh d° metal alloxilatoC4m~ hydropetmiie combimfim~ A comparkm with peraceficatiL It was claimed that la can be epoxidized with cumene hydroperoxide in the presence of catalytic amounts of [VO(acac)2] to give 98% 1,2-epoxy-3-methyl-3-butanol, 3a (at 70% conversion). This epoxide could be separated in 99% purity from the reaction mixture by distillation [22]. Due to the chelating effect of the acetylacetonate ligand, it may be better to use oxovanadium alkoxide complexes. The vVO(OR) moiety has been known for many years [23-25], but only recently has its chemistry attracted significant attention. For example, it was found useful in catalytic systems for the oxidative cleavage of C-C bonds of ketones [26] and

547 of a-diols [27]; these results were recently extended [28]. The observation that the -OR groups can be easily replaced by-OOR' and/or-OCR1R2R 3 led to vanadium oxo-alkoxides being used as precursors and/or catalysts for the epoxidation of tertiary allylic alcohols by alkyl hydroperoxide. Three substrates have been considered; the corresponding epoxides are 3a, 3b and 3r 0

0

OH

OH

OH

OH

la

3a

lb

3b

0

lc

3c

The reactions were carried out homogeneously in anhydrous benzene or toluene. The selectivity depends to some extent on the reaction conditions and must be determined for each case. Typical data are presented in Table 1 for catalytic systems involving t-BuOOH. Table 1 Epoxidation of 2-methyl-3-buten-2-ol, la, catalysed by "titanium alkoxide or tungsten or vanadium oxoalkoxide/benzene or toluene/t-BuOOH" systems. Run

T (~

Time (h)

1

25

24

2

40

24

3

25

24

4

25

5

Precursor [M] (mol.1q) [OW(OC2Hs)4]

0.07 [Ti(OC3HT)4] 0.02 [OV(OC3HT)3] 0.02 [OV(OC3HT)3] 0.01

3a ~elds 3a/[M]

la (mol.l-*)

SolvenP (cm3)

t-BuOOHb (mol.1q)

Conversionr (%)

1.4

6

1.6

5

4 and 5

1.4

6

1.6

10

major 99

1.4

6

1.6

98

67

98

2

7.5

2.2

99

192

98

Selectivityd (%)

a benzene or toluene; b anhydrous t-BuOOH (17% benzene or toluene solution); c % of substrate, la, consumed; d 100 (mol 3aJmol la converted); c,d products analysed by GC (OV1701) and GC-MS coupling; internal standard: 2a added at the end of the reaction. The formation of 3-methyl-2-buten-l-ol, 4, and of 3-methyl-2-butenal, 5, in some experiments can be related to the isomerization of la and subsequent oxidation of product 4, thus formed, to 5. We observed that the vanadium oxo-alkoxide, [OV(OC3H7)3], is a good precursor for the catalytic epoxidation of substrates la, lb and lc with systems involving tertbutyl hydroperoxide and tertiary allylic alcohols. These systems proved to be more active than those involving [VO(acac)2] usually associated with anhydrous t-BuOOH in benzene or toluene (Figure 1 and Table 2). Moreover, even after 24 h, with the acetylacetonato complex, the reaction is incomplete with lb and le (Table 2). Epoxidation of these substrates give satisfactory yields and good turnover numbers (ca. 200, Table 2); there is no evidence for a deactivation process, so even higher TON (ca. 1500) can be obtained under appropriate conditions. Not detected were 4 and 5 and the ethers (7 and 8) which are normally formed under conditions of acid catalysis.

548 91}

2

0

2 2' l'

8 t~

I

I

2

3

0 1

I

..........I~ .............. ii I ....

4 time (h)

5

6

'., 1

7

Figure 1 Table 2 Epoxidation of lb, lc and 4 catalysed by [OV(acac)2] or [OV(OCaHT)3] / benzene or toluene / t-BuOOH systems. Run

T (~

Time (h)

5

25

24

6

25

5

7

25

24

8

25

5

9

25

12

10

0

2

Precursor Substrate [M] (mol.l"l) (mol.lq) [OV(acac)2] lb 0.02 (1.5) [OV(OC3HT)3] lb 0.01 (2.0) [OV(acac)2] lc 0.02 (1.5) IOV(OC3HT)3]

lC

O.O1 [OV(acac)2] 0.01 [OV(OC3HT)~] 0.008

(2.0) 4 (1.0) 4 (2.0)

Solvent" (cm3)

Epoxide yields t-BuOOHb . Conversion~ Epoxide/ Selectivitya (mol.l") (%) [M] (%)

6

1.6

92

70

3b (97)

7.5

2.2

99

197

3b (98)

6

1.6

90

68

3C (98)

7.5

2.2

99

196

3C (97)

6

1.6

100

98

6 (99)

7.5

2.2

99

222

6 (99)

a,b see Table 1; c % of substrate consumed; d 100 (mol epoxide/mol substrate converted); GC and GC-MS determinations.

\

/--o. 4

\

/=o 5

c,d

OH \ 6

7

/

8

The reaction of 4 with the "[VO(C3HT)3]/t-BuOOH/toluene or benzene" systems was also examined (runs 9 - 10). This substrate gives the highest epoxide yield (ca. 99%) in a short time, even at 0~ these experimental conditions favour a highly selective and active system. Peracetic acid (aqueous solution or 33% in dichloromethane) can also give high conversions

549 and selectivities for epoxidation at room temperature, up to 94% and 97%, respectively, but on a large scale the "[VO(OC3H7)3]/t-BuOOH/benzene or toluene" systems are safer and more advantageous. Three other catalytic systems are now considered employing the environmentally acceptable hydrogen peroxide or dioxygen.

2.2. A biphase system : "CH3ReO3 o r M T O / H 2 0 2 - H 2 0 / C H 2 C I 2 " for epoxidation of acidsensitive epoxides Methylrhenium trioxide, CHsReO3 or MTO, has been proposed for epoxidation in homogeneous media with aqueous acidic hydrogen peroxide and tert-butanol mixtures [29]; the transformation may proceed via an rl2-peroxorhenium intermediate [30]. The Br~nsted acidity leads to subsequent reactions, mainly related to the acid-catalysed hydrolysis of the epoxide(s). PTC systems and a biphase medium were considered in order to avoid these undesired processes [31]. Basically, a biphase system with an aqueous and an organic phase would be simpler than phase transfer catalysis systems involving onium salts (vide infra), for which there are dramatic effects of the nature of the cation Q+, of the assembling ligands (PO43-, AsO43-, SO42-, O2-, R2SiO22-, etc.) and of the ratio [Q+]/[M] where M - Mo, W, etc. Optimization has to take into account the existence of several anionic species and of several equilibria [32]. In many cases, when the olefin and the epoxide are dissolved in the organic phase (dichloromethane or toluene, for example), while the aqueous phase is a reservoir of peroxidic oxygen which can regenerate the activated peroxo species, another method for the preparation of acid-sensitive epoxides can be developed and the use of a biphase medium may make it possible to suppress subsequent reactions. For example, epoxidation of relatively weakly nucleophilic terminal alkenes, such as oct-1-ene, is possible even at room temperature. For a 30 h reaction time, overall conversion (95%) and 1,2-epoxy-octane selectivity (99%) with the biphase "MTO/H2Ojoct-l-ene (in CH2C12)" system (molar ratios 1:150:100) are significantly different from those obtained with the homogeneous "MTO/H2Ojt-BuOH/oct-1ene" system (diol yield 70% for a 3 h reaction at 20~ Previous work has shown that inorganic stabilizers, such as phosphates, incorporated into some commercial aqueous hydrogen peroxide solutions, can reduce the selectivity. Even with phosphate-free 30% H202 solutions and for several substrates, the ideal selectivity cannot be obtained with the homogeneous system. With the "MTOd-I202-H20/CH2CI2" system at room temperature, la conversion (90%) and selectivity for 3a (99%) are again significantly better than with the homogeneous medium, for which these ratios are lower (50 and 20%, respectively; 24 h reaction for the two sets of experiments at 20~ 2.3. Phase Transfer Catalysis (PTC) with tungsten anionic peroxo species for epoxidation of la Heteropolyoxoperoxotungstate complexes containing the rl2-rl1 peroxo linkage are edso effective oxidation catalysts with hydrogen peroxide. The epoxidation of non-functionalized alkenes using these tungsten precursors is well documented [33,34]. In previous work [35], some results with demanding substrates such as (R)-(+)-limonene and oct-1-ene with two different systems were reported: -"H2WOa/H2OJHaPOa/Arquad| (I) -"[HPW20~4](Arquad| (II) We have shown that the active species is basically the same in both cases. These two precursors can epoxidize la, and the results favour system (I). In each case the ratio [W] to

550 [la] was 1:100 (7 h reaction at 60~ With system (I), la conversion is 99% with a high selectivity for 3a (95%); with system (II) these values are 50% and 95%, respectively. The differences are tentatively attributed to the pH of the aqueous phase (4.4 for II and 0.5 for I). Ion pairing between the polyanion and the phase transfer agent might be less effective with (II) than with (I), and we observed that the recycling of the catalytic species from peroxoand oxo- moieties is pH-dependent. 2.4. Catalytic epoxidation of la by dioxygen in the presence of transition metal-based precursors and a coreducer In the presence of dioxygen (P c a . 0.1 MPa), at room temperature (or at 40~ dichloro1,2-ethane solutions of la and Me2CHCHO give 3a and Me2CHCO2H in the presence of a soluble catalyst or of some highly divided inorganic oxides. Many soluble metal complexes can be used: [TiO(acac)2], [Ti(O-iPr)4], [mu(acac)3], [Ru2In'n CI(CH3COO)4], [Ru(H20)6] (MeC6H4-SO3)2. 3H20, "[P~uC13.xH20]", [RuCl2(PPh3)3], [Ru~3)5C1]C12, PdCl2, Pd(OAc)2, Pd~ zeolite, [Rh(OAc)2]2, [Rh(COD)CI]2, "RhC13.3H20", etc. Inorganic oxides (silica, TiO2, Ti-silicalite, KTiOPO4, VOPO4.2H20, etc.) are among the most efficient precursors or catalysts. With these early transition metal precursors the overall reaction is highly selective and yields can be higher than 93%. On the other hand, with the late transition metal precursors, except [Ru(NH3)sCI]C12, side reactions are frequently observed. Yields from other tertiary allylic alcohols such as lb and l c range between 85 and 98% [36]. The simplicity of the experimental conditions coupled with the high yields observed may encourage wider application of this methodology involving dioxygen at room temperature. 2.5. lsomerization of the tertiary epoxy-alcohols LiBr associated with tributylphosphine oxide or hexamethylphosphoric triamide (HMPA) is one of the most efficient systems [13]. It was suggested that the key intermediate species involves an acid-base interaction between Li+ and the epoxide [ 12]. However, the activity and selectivity of this system depend very much on the substrate. Transition metal complexes are attractive [ 14,15]. Palladium complexes [ 16,17] have been shown to catalyse isomerization of epoxides with good activities and regioselectivities, and are particularly suitable for substrates bearing acid-sensitive functionalities. For model Rh I complexes a catalytic cycle [15,37,38] was proposed involving the oxidative addition of the epoxide and the formation of a rhodium hydride and ketone by 13-H-elimination. Dicobaltoctacarbonyl [ 19,20] was reported to catalyse epoxide rearrangement but it cannot be used on a large scale because of its possible dismutation. More recently, nickel-based catalysts were used [21], and an isomerization mechanism involving a low-valent nickel active species and an oxanickelacyclobutane intermediate was proposed. The advantage of this system is the possibility of developing equivalent supported or grafted metal complexes. The performances of three systems, LiBr/HMPA (1:1), Co2(CO)8 and Ni complexes, were therefore compared for the isomerization of three epoxy-alcohols, the results are summarized in Table 3. 2.5.1. Comparison of some catalytic systems The performances of LiBr/H A (1:1) for the isomerization of epoxy-alcohols 3a, 3b and 3c were tested in toluene. A relatively high temperature (85~ is required to obtain good conversion (82-100%). The selectivity for the keto-alcohol is very high (_> 97%). The reaction

551

may be catalytic (Run 12-Table 3) but the conversion is improved under stoichiometric conditions (Run 11-Table 3). Unexpectedly, the catalytic performances are much better than with epoxy-l,2-octane [39]. With transition metal-based systems, the reaction is slower. Using a solution of Co2(CO)8 in methanol, the isomerization of the epoxy-alcohol proceeds under mild conditions (35~ The reaction is catalytic but not complete (conversion c a . 50%). As for LiBr-HMPA (1:1), the keto-alcohol is obtained with good selectivity (Run 13). The conversion is lower than for epoxy- 1,2-octane. Table 3 Li-, Co- or Ni- based systems for catalysed ring opening reactions of 1,2-epoxy-alcohols, 3a, 3b and 3e; a two-step synthesis of keto-alcohols, 2a, 2b and 2c. Keto-alcohol fields Run

T (~

Time (h)

11

85

2

12

85

2

13

35

24

14

50

46

15

50

46

16

50

46

17

50

20

18

85

2

19

35

24

20

35

24

21

85

2

Precursor [M']" (mol.1-1)

Epoxyalcohol (mol.l1)

LiBr/HMPA 0.75 0.75 LiBr/HMPA 0.75 0.75 Co2(CO)8 0.05 NiBre(PPh3)2/Zn/PPh3 0.01 0.1 0.01 NiBr2(PPh3)JZn/PPh3 0.016 0.160.016 NiBr2(PPh3)2/Zn/TBA 0.016 0.160.016 NiBr2(PPh3)JZn/PPh3/ 0.016 0.16 0.016 TBA 0.016 LiBr/HMPA 0.75 0.75 Co2(CO)8 0.05 Co2(CO)8 0.05 LiBr/HMPA 0.75 0.75

3a (0.75) 3a (4.5) 3a (1.52) 3a (0.12) 3a (0.38) 3a (0.16) 3a (0.16)

3b (4.5) 31a (1.54) 3c (1.5) 3c (4.5)

Solvent (cm 3)

Conversion b

EpoxySelectivity c alcohol/[M'] (%)

Toluene 10 Toluene 10 MeOH 1.5 CH2C12 4 THF 3 THF 3 THF 3

100

1

2a (98)

85

5

2a (98)

50

15

2a (99)

37

-

43

10

16

2

24

2

la (97) 2a (3) 2a (96) la(4) 2a (96) la(2) 2a (96) la(3)

Toluene 10 MeOH 1.5 MeOH 1.5 Toluene 10

83

5

2b (99)

44

14

2b (98)

48

14

2c (98)

82

5

2e (97)

a [M]

----lithium, cobalt or nickel precursors; b0~ of substrate consumed; c 100 (mol 2/mol 3 converted); b'cproducts analysed by GC (OV-1701) and GC-MS coupling;TBA---tributylamine.

There is a dramatic solvent effect with the nickel-based system (Runs 14-15). A tetrahydrofuran (THF) solution of the nickel-based catalyst appears to be less active than the

552 dicobaltoctacarbonyl complex. NiBr2(PPh3)2/PPh3/Zn (1:1:10) (referred to as Ni/P/Zn) catalyses isomerization of 3a with a TON similar to those obtained by Miyashita et al. [21 ] for epoxy-propane isomerization. Good selectivity for the keto-alcohol, similar to that observed for the above systems is obtained. The same catalytic performances were observed for epoxy1,2-octane. On the other hand, with a non-coordinating solvent such as CH2C12, la is "regenerated". When PPh3 is replaced by tributylamine (TBA), the selectivity is unchanged but the reaction rate decreases and the catalyst deactivates quickly. NMR studies show that, in contrast to PPh3, TBA is not bound to the nickel centre. TBA is not convenient for building a stable active complex and, when used as adduct, it does not significantly modify the catalytic performances as it does with palladium catalysts in the Heck reaction [40]. Trialkyl phosphine is more promising and is under study in our group as well as novel heterogeneous systems. Epoxides 3b and 3c are selectively isomerized by LiBr/HMPA in a short time (Runs 18, 21).

2.5.2. EPR study of the nickel-based system In order to characterize the active site of the homogeneous nickel-based catalyst (oxidation state, composition of the coordination sphere), the reaction was monitored by EPR. Before addition of the epoxide The catalytic solution is obtained by mixing NiIIBr2(PPh3)2 (0.016 mol.r 1) with a reducing agent (Zn) in the presence of 1-1.6.10 -2 mol.1-] additional PPh3 which may also act as a reducing agent, as already described for palladium-based catalysts [40] or supported nickel complexes [41]. The best catalytic performances are obtained when the solution is prepared before adding the epoxide, suggesting that the active species is generated by the reduction of Ni II to a Ni 0 or Ni I complex. EPR is expected to be able to discriminate between these various oxidation states. Most Ni II complexes, especially low symmetry ones [42], are EPR-silent; only perfect octahedral Ni II complexes or Ni II ions with an isotropic environment exhibit an isotropic signal near g = 2.17 [43,44]. Ni 0 which is diamagnetic, is also EPR-silent except in metallic particles, which exhibit a broad ferromagnetic resonance signal [45]. EPR is however a suitable technique for characterizing Ni I. At room temperature, before addition of the epoxide, the catalytic solution exhibits an isotropic EPR signal A at giso = 2.21 (Figure 2a). However, a broad axial signal B located at g_L = 2.40 > gll = 2.07 > ge (Figure 2b) is observed when the EPR spectrum of the same solution is recorded at 77K. Both signals are due to the same complex because (i.e. the average g value expected from signal B). The initial giso = (2g• NiIIBr2(PPh3)2/THF solution presents a very broad EPR signal, which is ten to twenty times less intense than signals A and B. The g tensor values of signals A and B lie in the range expected for Ni I d 9 ions [46-48]. Our findings suggest that the Ni I complex is monomeric and rule out the formation of a ligandbridged bimetallic species [49] [the half-field region (g = 4) is EPR-silent; no broadening of the EPR signal related to spin-spin interaction was observed]. Many Ni I complex solutions show such an evolution of their EPR spectrum with temperature [50-52]. The relative order of the g tensor components of signal B suggests a dz2 ground state for the single electron, and consequently a distorted tetrahedral [53] or trigonal-bipyramidal symmetry [54,55] for the corresponding Ni I complex. Unfortunately, because of the poorly

553 resolved superhyperfine structure (more apparent on the parallel component), no doubt resulting from the interaction of the unpaired electron with 31p (I = 1/2), 79Br or 81Br ( I - 3/2) nuclei of the Ni I ligands, the numbers of each ligand cannot be deduced. Nevertheless, signal B is very similar to that reported by Gmznyh et al. [56] for NilCI[P(n-C4H9)3]3. We therefore propose to assign signal B to NiIBr[PPh3]3. The observed g value shift of signal B compared to that of NilBr[P(n-C4H9)3]3 agrees with the decrease of ligand field strength expected on replacing P(n-C4H9)3 by PPh3 and/or C1 by Br. In order to measure the amount of Ni II reduced to Ni I in the presence of Zn, the intensities of signals A and B were estimated using a copper-based standard [57]. The number of spins measured at 77K for a sample containing initially 6.6 10-3 mol Ni II is (4.7 + 1.0) 10-3. This means that (70 + 14)%, i.e. most of the initial Ni II, is reduced to NiI.

L ~ ~2H~ _

DPPH J 100G

g|so

~-

IDPPH [g//=.2 07

=2.21

Figure 2a

Figure 2b

These findings suggest that NiIBr(PPh3)3 is the catalytic centre and explain why additional equimolar PPh3 is required: it replaces one Br atom to form the active complex. After addition of the epoxide The conversion never reaches 100% whatever the reaction time, suggesting deactivation of the catalyst. Indeed, after addition of the epoxide, the intensity of signal B decreases with time. Similarly, the initial red-orange colour fades with time and vanishes at the end of substrate conversion. This colour may be attributed to the active complex. The catalytic site proposed above involves an oxophilic Ni I centre. Deactivation might arise from oxidation of this active complex and simultaneous epoxide deoxygenation yielding the small amount of allylic tertiary alcohol la observed (Table 3). Toluene or dichloromethane (Run 14) solutions of NilIBr2(PPh3)2 do not exhibit A or B

type EPR signals in the presence of Zn. We propose that if Ni II is reduced to Ni I, under such conditions, it cannot be stabilized. This could explain the low activity of these samples for epoxide isomerization and show once more the influence of the solvent on catalyst performance. Our findings suggest that NilBr(PPh3)3 is the active site, in contrast to Miyashita et al. [21 ] who proposed a Ni ~ active site.

554 2.5.3. Silica-supported Ni I complexes The catalytic properties of the homogeneous catalyst were compared with those of similar well defined supported Ni I complexes. In order to confirm the nickel oxidation state in the active entity and to study the influence of the composition of its coordination sphere [coordination number, nature of the trialkylphosphine (TAP) ligands] on the catalytic performances, we have used a set of silicasupported unsaturated Ni I complexes [58] with a given number of TAP ligands of different electronic and steric properties [59,60]. NiI(pPh3)2(O)2 (O = silica surface oxygen atom), whose coordination sphere may be similar to those of the homogeneous epoxide isomerization catalysts, led to the conversion of 15% of epoxy-l,2-octane to 2-octanone after a 24 h reaction at 50~ In contrast, NiI(PEt3)(O)2 and NiI(p(c-C6Hll)3)(O)2 showed no ability to isomerize epoxy-alcohol 3a. The immediate colour change observed after epoxide addition suggests the formation of Ni II by Ni I oxidation in the presence of epoxide. The latter catalysts are more oxophilic than NiI(pPh3)2(O)2 because of their high unsaturation and because PEt3 and P(c-C6Hll)3 are more electron-donating than PPh3. These findings are in agreement with the results of Miyashita et al. [21 ]: they found that the deoxygenation reaction becomes predominant when the c~-basicity of the trialkylphosphines bound to nickel is increased. In any case, these preliminary results obtained with NiI(pPh3)2(O)2 are promising; they are consistent with a Nikbased active species in the homogeneous catalytic system and encourage us to pursue this approach to improving catalyst performance.

3. CONCLUSION The advantage of these systems over traditional catalytic ketonization routes lies in the selectivities and conversions; the two-step reaction path is very convenient from a preparative point of view. Work is continuing to control these systems better, and to investigate their mechanisms and their synthetic utility. Novel and highly efficient ways of isomerizing epoxyalcohols by careful control of the oxophilic character of the active species are sought. ACKNOWLEDGMENTS We thank Dr. J.S. Lomas for constructive discussions and for correcting the manuscript. REFERENCES 1. A.B. Smith, P.A. Levenberg, P. J. Jerris, R. M. Scarborough and P. M. Wovkulich, J. Am. Chem. Soc., 103 (1981) 1501. 2. G. B0chi and H. W0est, J. Am. Chem. Soc., 100 (1978) 294. 3. W. Adam, M. Braun, A. Griesbeck, V. Lucchini, E. Staab and B. Will, J. Am. Chem. Soc., 111 (1989) 203. 4. A.L. Baumstark and P. C. Vasquez, J. Org. Chem., 53 (1988) 3437. 5. M.S. Newman and V. Lee, J. Org. Chem., 40 (1975) 381. 6. F. Derdar, J. Martin, C. Martin, J.-M. Br6geault and J. Mercier, J. Organometal. Chem., 338 (1988) C21. 7. R.E. Parker and N. S. Isaacs, Chem. Rev., 59 (1959) 737.

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10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32.

33. 34. 35. 36. 37. 38. 39.

R. J. Gritter, in Chemistry of Functional Groups, S. Patai Ed., Wiley, New York, 3 (1967) 400. A. S. Rao, S. K. Paknikak and J. G. Kirtane, Tetrahedron, 39 (1983) 2323. Organic Reactions, Wiley, New-York, 29 (1983) 345. V. N. Yandovskii, B. A. Ershov, Russ. Chem. Rev., 41 (1972) 403. B. Rickbom and R.M. Gerkin, J. Am. Chem. Soc., 90 (1968) 4193. ibid, 93 (1971) 1693. G. Magnbusson and S. Thoren, J. Org. Chem., 38 (1973) 1380. H. Alper, D. Des Roches, T. Durst and R. Legault, J. Org. Chem., 41 (1976) 3611. D. Milstein, O. Buchman, J. Blum, J. Org. Chem., 42 (1977) 2299. S. Kulasegaram and R. J. Kulawiec, J. Org. Chem., 59 (1994) 7195 and references therein. J. H. Kim and R. J. Kulawiec, J. Org. Chem., 61 (1996) 7656 and references therein. A. Cabrera, F. Mathe, Y. Castanet, A. Mortreux and F. Petit, J. Mol. Cat., 64 (1991) Lll. J. L. Eisenmann, J. Org. Chem., 27 (1962) 2706. J. Prandi, J. L. Namy, G. Menoret and H. B. Kagan, J. Organometal. Chem., 285 (1985) 449. A. Miyashita, T. Shimada, A. Sugawara and H.Nohira, Chem. Let., (1986) 1323. Japan Kokai Tokkyo Koho, JP 60023376, A2 850205, Showa to Mitsui Toatsu Chemicals, Inc., Japan ; CA 103 : 6207. W. Prandtl and L. Z. Hess, Z. Anorg. Chem., 82 (1913) 103. C. N. Caughlan, H. M. Smith and K. Watenpaugh, Inorg. Chem., 5 (1966) 2131. L. G. Hubert-Pfalzgraf, Chem. Rev., 90 (1990) 969. J.-M. Brrgeault, B. E1 Ali, J. Mercier, J. Martin and C. Martin, C.R. Acad. Sci. Paris, 307, Serie II (1988) 2011. J.-M. Bregeault, B. El Ali, J. Mercier, J. Martin and C. Martin, C.R. Acad. Sci. Paris, 309, Serie II (1989)459. A. Atlamsani, J.-M. Bregeault and M. Ziyad, J. Org. Chem., 58 (1993) 5663. W. A. Herrmann, R. W. Fischer and D. W. Marz, Angew. Chem. Int. Ed. Engl., 30 (1991) 1638. A. M. AI-Ajlouni and J. H. Espenson, J. Org. Chem., 61 (1996) 3969. J.-M. Bregeault et al., unpublished results. J.-M. Bregeault, R. Thouvenot, S. Zoughebi, L. Salles, A. Atlamsani, E. Duprey, C. Aubry, F. Robert and G. Chottard, in New Developments in Selective Oxidation II, V. Cortes Corberb.n and S. Vie Bellrn Eds., Elsevier Sci. Publ., Amsterdam (1994) 571. I. V. Kozhevnikov, Catal. Rev.-Sci. Eng., 37 (1995) 311. C. L. Hill and C. M. Prosser-McCartha, Coord. Chem. Rev., 143 (1995) 407. L. Salles, C. Aubry, R. Thouvenot, F. Robert, C. Doremieux-Morin, G. Chottard, H. Ledon, Y. Jeannin and J.-M. Bregeault, Inorg. Chem., 38 (1994) 871. A. Atlamsani, E. Pedraza, C. Potvin, E. Duprey, O. Mohammedi and J.-M. Bregeault, C.R. Acad. Sci. Paris, 317 (1993) 757. R. Grigg, R. Hayes and A. Sweeney, J. Chem. Soc., Chem. Comm., (1971) 1248. G. Adams, C. Bibby and R. Grigg, J. Chem. Soc., Chem. Comm., (1972) 491. F. Ziani-Derdar, Thesis, Univ. Pierre et Marie Curie, Paris, France (1988).

556 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60.

C. Amatore, E. Carrd, A. Jutand, M. A. M'Barki and G. Meyer, Organometallics, 14 (1995) 5605. C. Lepetit, Thesis, Univ. Pierre et Marie Curie, Paris, France (1987). A. Bencini and D. Gatteschi, in Transition Metal Chemistry, G. A. Melson and B. N. Figgis Eds, Marcel Dekker, New York, 8 (1982) 1. J. E. Bolton and J. R. Wertz, Electron Spin Resonance, Me Graw-Hill, New York, (1972) 291. L. Bonneviot, M. Che, K. Dyrek, R. Sch611ner et G. Wen&, J. Phys. Chem., 90 (1986) 2379. L. Bonneviot, M. Che, D. Olivier, G. A. Martin and E. Freund, J. Phys. Chem., 90 (1986)2112. K. Nag, A. Chakravorty, Coord. Chem. Rev., 33 (1980) 87. C. Lepetit, M. Kermarec and D. Olivier, J. Mol. Catal., 51 (1989) 73. J. G. Van Ommen, J. G.M. Van Rens and P. J. Gellings, J. Mol. Catal., 13 (1981) 313. M. F. Rettig, E. A. Kirk and P. M. Maitlis, J. Organometal. Chem., 111 (1976) 113. M. J. Nilges, E. K. Barefield, R. L. Belford and P. H. Davis, J. Am. Chem. Sot., 99 (1977) 755. E. K. Barefield, D. A. Krost, D. S. Edwards, D. G. Van Derveer, R. L. Trytko and S. P. O' Rear, J. Am. Chem. Sot., 103 (1981) 6219. G. A. Bowmaker, P. D. W. Boyd and G. K. Campbell, Inorg. Chem., 22 (1983) 1208. reference 43, p. 324 M. Kermarec, D. Olivier, M. Richard, M. Che and F. Bozon-Verduraz, J. Phys. Chem., 86 (1982) 2818. R. Barbucci, A. Bertcini and D. Gatteschi, Inorg. Chem., 16 (1977) 2117. B. A. Gruznyh, B. B. Saraev, F. K. Schmidt, G. M. Warin and E. H. Siedyh, Koord. Khim., 9 (1983) 1400. K. Dyrek, A. Rokosz and A. Medej, Appl. Magn. Reson., 6 (1994) 309. F. X. Cai, C. Lepetit, M. Kermarec and D. Olivier, J. Mol. Catal., 43 (1987) 93. M. Kermarec, C. Lepetit, F. X. Cai and D. Olivier, J. Chem. Sot., Faraday Trans. I, 85 (1989) 1991. C. Lepetit, M. Kermarec and D. Olivier, J. Mol. Cat., 51 (1989) 95.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

557

M e t a l - c a t a l y z e d oxidations with alkyl hydroperoxides: a c o m p a r i s o n b e t w e e n

tert-butyl

h y d r o p e r o x i d e and p i n a n e h y d r o p e r o x i d e

H.E.B. Lempers and R.A. Sheldon Laboratory of Organic Chemistry and Catalysis, Delft University of Technology, Julianalaan 136, 2628 BL, Delft, The Netherlands

The oxidizing capacities of the bulky pinane hydroperoxide (PHP) and

tert-butyl

hydroperoxide (TBHP) were compared in oxidations catalyzed by Mo, V, Se, Os and Ru. The general trend is that metals which react via an oxometal pathway show a very similar behaviour using these two hydroperoxides as oxidant, e.g. the selenium catalyzed allylic oxidation of olefins to the corresponding c~, [3-unsaturated alcohols. Reactions which involve a peroxometal pathway, e.g. the molybdenum catalyzed epoxidations of olefins, show a completely different behaviour using these two hydroperoxides, namely virtually no reaction is observed with the bulky PHP. We conclude that PHP is a suitable mechanistic probe for distinguishing between oxometal and peroxometal pathways in catalytic oxidation.

1. INTRODUCTION In homogeneous metal catalyzed oxidations with alkyl hydroperoxides two different reaction mechanisms can be distinguished: the peroxometal- and oxometal mechanism [1] (Scheme 1). Peroxometal mechanisms are favoured when the metal in its highest oxidation state is both a Lewis acid and a weak oxidant, e.g. Mo(VI), W(VI) or Ti(IV) [2,3]. Polar solvents, particularly alcohols greatly retard the reaction by competing with the hydroperoxide for coordination sites on the metal [4]. Substituent effects on the olefin are consistent with the active alkyl peroxometal species being electrophilic in nature [5]. Oxometal mechanisms are favoured with later and/or first row transition elements having a high oxidation potential, e.g. Os(VIII), Ru(VIII) and Cr(VI) [2]. Generally, polar solvents

558 have a much less negative effect than is observed for peroxometal catalyzed reactions. For example, osmium catalyzed dihydroxylations are camed out in tert-butanol as solvent [6].

-HX

M_O2 R

S

r_

MOR + SO

PEROXOMETAL PATHWAY MX + RO2H

~-

ROH

M--O

S

MX + SO

I

X OXOMETAL PATHWAY Scheme 1

A borderline case is V(V) which has an intermediate oxidation potential [3] and can react via both mechanisms depending on the substrate used. It catalyzes epoxidations with TBHP via the peroxometal mechanism and alcohol oxidations via the oxometal pathway [7]. In this present investigation we used pinane hydroperoxide (PHP) as a mechanistic probe to distinguish between peroxometal and oxometal pathways. In the peroxometal pathway the bulky pinane group is present in the active oxidant while in the oxometal pathway it is not. Hence, in the case of the peroxometal mechanism one might expect more steric constraints and consequently a slower reaction using the bulky PHP compared to the much less bulky TBHP [8]. In the case of oxometal mechanisms the difference between PHP and TBHP should be much smaller as the alkyl group is not present in the active oxidant. PHP is also of industrial interest [9]. Catalytic hydrogenation affords pinanol which is pyrolized to linalool. The latter is used in the fragrance industry and in the synthesis of vitamin E. Hence we were interested in utilizing the active oxygen in PHP to perform useful oxidations.

559

OH OH OOH TBHP

PHP 1

2

3

OH

HO

J J7

~ovi 6 ~

OH

2. RESULTS AND DISCUSSION 2.1. Molybdenum catalyzed epoxidation Previous studies indicated that the structure of the alkyl hydroperoxide in molybdenum catalyzed epoxidations has only a minor effect on the rate and selectivity [ 10]. Hence, we were initially surprised to observe that PHP failed to give the expected epoxidation of cyclohexene (1) and limonene (2) in the presence of a molybdenum catalyst (Table1). Epoxidation of limonene with TBHP as oxidant, in contrast, gave the epoxide of the more highly substituted double bond in 84% selectivity, consistent with nucleophilic attack of the olefin on the alkylperoxomolybdenum(VI) [3,5]. We tentatively concluded that this low reactivity of PHP is a result of steric hindrance in the putative alkylperoxomolybdenum(VI) intermediate. This prompted us to carry out a systematic investigation [8] of steric effects of the alkyl substituents in the alkyl hydroperoxide on the rate of molybdenum catalyzed epoxidations. Molybdenum-catalyzed epoxidations of cyclohexene under pseudo first-order reaction conditions showed the highest rate, of the investigated tertiary alkyl hydroperoxides (Figure 1), with TBHP.

560 Table 1. Molybdenum catalyzed a epoxidation with TBHP

Successive substitution

and PHP as oxidant

of the methyl groups in TBHP by higher

substrate

oxidant

conv. (%)

sel. (%)b

1

TBHP

89

96

reaction rate. Substitution

2

TBHP

81

84

of all three methyls by

1

PHP

0

2

PHP

4

alkyl groups resulted in a steadily decreasing

three ethyl groups 0

resulted in a 99% decrease in reaction

a Conditions: 10 mmol substrate, 10 mmol oxidant,

rate compared to TBHP

0.1 mmol Mo(CO)6, 1 g internal standard was heated

and a further increase

for 5 h at 80~ in 10 ml chlorobenzene

in the steric bulk of the

b selectivity to epoxide

alkyl group as in PHP resulted in a complete

loss of epoxidation activity. These results are rationalized on the basis of a peroxometal mechanism involving nucleophilic attack of the olefin on an alkylperoxomolybdenum(VI) intermediate [8]. Bulky substituents at the o~ position in the alkyl hydroperoxide seriously impede the approach of the olefin to the O - - O bond.

lOO O

,~==.(

9

~8 o 9 '--2.*"

'

~.,

0 9

I"1

(C3H7)(CH3)2COOH

. ~

(C2Hs)2(CH3)COOH

A

i

10

I

I

I

10

20

30

time (min)

Figure 1

(CH3)3COOH

(C2H5)3COO H

561 2.2. Vanadium catalyzed oxidations

Similar results were observed with vanadium catalyzed epoxidation of cyclohexene (1) and limonene (2) with TBHP and PHP (Table 2). Epoxidations with TBHP gave reasonable conversions and Table 2. Vanadium catalyzed a epoxidation with TBHP

selectivities while

and PHP as oxidant

PHP gave virtually no reaction, consistent with

substrate

oxidant

conv. (%)

selec. (%)b

a peroxometal mechanism [ 11, 12].

1

TBHP

66

34

2

TBHP

60

70

catalyzed epoxidation of

1

PHP

6

7

cyclohexenol (3) and

2

PHP

10

0

carveol (4) gave high

3

TBHP

89

98

substrate conversions

4

TB HP

100

80 c

and high epoxide

3

PHP

61

100

selectivities both with

4

PHP

100

85 c

PHP and TBHP [13].

5

TBHP

100

> 95

These results can be

6

TBHP

99

> 95

rationalized by assuming

7

TBHP

99

> 95

that efficient coordination

5

PHP

0

6

PHP

57

> 95

of the allylic alcohol to

7

PHP

88

> 95

the alkylperoxovana-

In contrast, the vanadium

of the hydroxyl group

dium(V) forces the a Conditions: 10 mmol substrate, 10 mmol oxidant,

oxidant and substrate

0.1 mmol VO(acac)2 and 1 g internal standard were heated

into close proximity, thus

for 5h at 80~

facilitating intramolecu-

b selectivity to o~,13-epoxy alcohol.

lar oxygen transfer to the

c the rest is o~,13-unsaturated ketone.

double bond. The electron rich allylic

alcohol (7) exhibited high conversions with both hydroperoxides. When the electron density of the substrate (6) was decreased we observed no effect when TBHP was used, while with PHP a decrease in conversion was observed from, 88% to 57%. A further decrease in electron density of the double bond, as in allyl alcohol (5) resulted in a complete loss of activity with PHP while with TBHP no effect was observed. Hence, the observed differences between

562 cyclohexenol epoxidation with TBHP and PHP are a reflection of the intermediate reactivity of this olefin. The double bond of carveol, on the other hand, is more substituted and shows no difference in reactivity with these two alkyl hydroperoxides. 2.3 Selenium catalyzed oxidation SeO2 displays a unique mode of interaction with olefins, involving an initial ene reaction followed by a [2,3] sigmatropic rearrangement to a Se(II) species [14] (Scheme 2). H ene

[2,31

H O • e

+ TBHP ~ - TBA

O

H

+ SeO 2

Scheme 2 The resulting Se(II) is reoxidized by the TBHP. Hence, selenium-catalyzed allylic oxidations with RO2H involve an oxometal pathway and, assuming that this is the rate-limiting step, the rate should not significantly be influenced by the structure of RO2H. This is indeed the case: using 0.05 equivalent of SeO2 and 1.5 equivalent of oxidant at room temperature smooth allylic oxidation was observed with both TBHP and PHP (Table 3). The reactivity order for Table 3. Selenium catalyzed a allylic oxidation

selenium catalyzed oxida-

substrate

tion of allylic C - - H

oxidant

conv. (%)

sel. (%)b

groups is CH2>CH3>CH 8

TBHP

100

96

[ 15, 16]. This order of

9

TBHP

68

92

reactivity was observed

8

PHP

100

99

in the selenium-catalyzed

9

PHP

91

95

allylic oxidation of ~-pinene (8) and

a Conditions: 10 mmol substrate, 15 mmol oxidant, 0.2 mmol SeO2 and 1 g internal standard were stirred

2-carene (9) which both

for 24 h at room temperature.

dation. We studied the

b selectivity to the cx,13-unsaturated alcohol

selenium-catalyzed

showed mainly CH 2 oxi-

allylic oxidation of 13-pinene with TBHP and PHP in more detail, using an olefin/alkyl hydroperoxide ratio of 10. Under these conditions we observed zero order kinetics rather than the expected pseudo-first

563 order reaction. The selenium catalyzed allylic oxidation of 13-pinene gave similar reaction rates with PHP (4.8 910 -4 mol-1-1 .min -1 ) [ 17] and TBHP (5.0.10-4 mol.1-1 .min -1 ) which is consistent with the idea that the reoxidation of Se(II) by RO2H is not the rate-limiting step. Synthetic applications of the SeO2 allylic oxidation have revealed that the reaction gives predominantly E-allylic alcohol products [ 18]. These results parallel the geometric preference shown by [3,3] sigmatropic rearrangements of the Cope and Claisen variety and apparently arise from the most favourable transition state geometry being a pseudo chair cyclohexane with the largest groups in the equatorial position [ 19, 20]. This similar geometric selectivity in the SeO2 oxidation of olefins indicates an analogous steric effect might be operating in the [2,3] pseudo cyclopentane transition state [21]. This preference for the E-allylic alcohol product for selenium catalyzed allylic oxidation was also investigated with PHP as oxidant. Allylic oxidation of a mixture of cis/trans dupical (10) yielded with both TBHP and PHP more than 95% of the E allylic alcohol (Figure 2), further demonstrating that selenium catalyzed oxidations are independent of the oxidant used.

~

O

I0

O

fO SeO 2

TBHP/PHP

OH

~21

10

13 C (ppm) substrate product

1

2

3

115.3 114.6 120.1

147.4 146.7 151.5

35.t3 37.8 76.13

Figure 2 2.4 Ruthenium and osmium catalyzed oxidation

Two other examples of metal-catalyzed oxidations with RO2H which involve an oxometal pathway are the OsO4 catalyzed dihydroxylation of olefins [22] and the RuC13 catalyzed oxidation of alcohols [23]. OsO4 catalyzed dihydroxylation is believed to involve a 2 + 2 cycloaddition of oxoosmium(VIII) to the olefinic double bond followed by

564 rearrangement to a cyclic osmate(VI) ester (Scheme 3). Reaction of the latter with RO2H/H 2 0 affords the diol product with regeneration of OsO4.

t)

+ OsO4

o os (v

O3Os/O~

o 2.

NO

+.

o.o HO

-t- OsO4 +

ROH

Scheme 3

This reoxidation of the relatively substitution inert osmium(VI) ester to a substitution labile osmium (VIII) ester is the rate-limiting step [22]. This step would be expected to be slower with the more sterically demanding PHP than with TBHP. This is indeed what we observed: OsO4-catalyzed dihydroxylation of cyclohexene (1, Table 4) gave a faster reaction with TBHP than with PHP, the final conversion being reached in 6h and 24h respectively. Table 4. Ruthenium a and Osmium b catalyzed oxidations with TBHP and PHP as oxidant substrate

catalyst

oxidant

conv. (%)

sel. (%)

4

RuC13

TBHP

57

72

4

RuC13

PHP

33

75

1

OsO4

TBHP

100

78

1

OsO4

PHP

73

65

a Conditions: 10 mmol substrate, 10 mmol oxidant, 0.05 mmol RuC13 and 1 g internal standard were stirred for 24 h at room temperature. b Conditions: 10 mmol olefin, 16 mmol oxidant, 20 ml tert-butanol, 0.75 m120% aqueous tetraethyl ammonium hydroxide and 0.005 mmol OsO4 were stirred for 24 h at 0~ RuC13-catalyzed alcohol oxidation of carveol (4, Table 4) similarly showed a higher rate with TBHP than with PHP, suggesting that the rate-limiting step in ruthenium catalyzed oxidation of alcohols may involve reaction of a ruthenium alkoxide with RO2H, resulting in formation of the carbonyl compound with simultaneous reoxidation of the ruthenium (Scheme 4 ).

565

/OH

n+

Ru--~- O

N

+/ck H

/ 0-

/c\ H

Rn+OH RO 2H

O.

\

X

~-

H

Ru/n+

n+

Ru~O

+

H20

"OH \ + /C=O

Scheme 4 3. Conclusions

In conclusion, we believe that a comparison of the oxidizing capacities of PHP and TBHP in metal-catalyzed oxidations can provide valuable mechanistic insights. When the reaction involves rate-limiting oxygen transfer from a peroxometal species to the substrate, e.g. in Moand V-catalyzed epoxidations the bulky PHP is not reactive. The steric constraints are less of a problem, however, when coordination of the substrate, e.g. in V-catalyzed epoxidation of allylic alcohols, provides for an intramolecular oxygen transfer. When the reaction involves reaction of an oxometal species with the substrate as the rate-limiting step little or no difference is observed. Intermediate reactivities may be observed when reoxidation of the catalyst by RO2H to the active oxometal species, is the rate-limiting step. Hence, we conclude that PHP is a suitable probe for distinguishing between alternative mechanistic pathways in catalytic oxidations with RO2H. Acknowledgement: We wish to thank the Netherlands Institute for Research on Catalysis (NIOK) for financial support and Quest International for supplying us with pinane hydroperoxide. 4. References

1

R.A. Sheldon and J. Dakka, Catal. Today, 19 (1994) 215.

2

R.A. Sheldon and J.K. Kochi, Metal-catalyzed Oxidations of Organic Compounds, Academic press, New York, 1981.

3

R.A. Sheldon, J.A. van Doom, C.W.A. Schram and A.J. de Jong, J. Catal., 31 (1973) 438.

566 R.A. Sheldon and J.A. van Doom, J. Catal., 31 (1973) 427. M.N. Sheng and J. G. Zajacek, J. Org. Chem., 35 (1970) 1839. K.B. Sharpless and K. Akashi, J. Am. Chem. Soc., 98 (1976) 1986. R.B. Dehnel and G.H. Witham, J. Chem. Soc., Perkin Trans. I, 4 (1979) 935. H.E.B. Lempers, M.J. van Crey and R.A. Sheldon, Recl. Trav. chim. Pays-Bas, 115 (1996) 542. Ullman's Encyclopedia of Industrial Chemistry, 4th ed., VCH Weinheim (Germany), 20 (1981) 199. 10

R.A. Sheldon, Applied Homogeneous Catalysis, Vol. 1 (W.A. Herrmann and B. Cornils, Eds.) VCH Weinheim (Germany) pp. 411-423.

11

H. Mimoun, M. Mignard, P. Brechot and L.Saussine, J. Am. Chem. Soc., 108 (1986) 3711.

12 13

A.O. Chong and K.B. Sharpless, J. Org. Chem., 42 (1977) 1587. H.E.B. Lempers, A. Ripoll~s i Garcia and R.A. Sheldon, to be submitted to J. Am. Chem. Soc.

14

M.A. Warpehoski, B. Chabaud and K.B. Sharpless, J. Org. Chem., 47 (1982) 2897.

15

A. Guillemonat, Ann. Chim., 11 (1939) 143.

16

U.T. Bhalerao and H. Rapoport, J. Am. Chem. Soc., 93 (1971) 4835.

17

Unpublished results, H.E.B. Lempers and R.A. Sheldon.

18

J.J. Plattner, U.T. Bhalerao and H. Rapoport, J. Am. Chem. Soc., 91 (1969) 4933.

19

D.J. Faulkner and M.R. Peterson, Tetrahedron Lett., (1969) 3243.

20

C.L. Perrin and D.J. Faulkner, Tetrahedron Lett., (1969) 2783.

21

D. Arigoni, A. Vasella, K.B. Sharpless and H.P. Jensen, J. Am. Chem. Soc., 95 (1973) 7917.

22

K.B. Sharpless, A.Y. Teranishi and J-E. B~ickvall, J. Am. Chem. Soc., 99 (1977) 3120.

23

S. Zhang and R.E. Sheperd, Inorg. Chim. Acta, 193 (1992) 217.

24

E. Erdik and D.S. Matteson, J. Org. Chem., 54 (1989) 2742.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

567

On the Way to Redox-Molecular Sieves as Multifunctional Solid Catalysts for the One-Step Conversion of Olefins to Aldehydes or Ketones M. van Klaveren and R. A. Sheldon Delft University of Technology, Laboratory of Organic Chemistry and Catalysis, Julianalaan 136, 2628 BL Delft, The Netherlands Ti-A1-Beta was used as a bifunctional catalyst in the one-step conversion of 2,3-dimethyl-2butene to 3,3-dimethyl-2-butanone. The Ti-AI-Beta catalyzed epoxidation of the olefin with aqueous H202 to the corresponding diol 2,3-dimethyl-2,3-butanediol using different solvents was studied. Good yields and high H202 selectivities were obtained when using polar aprotic solvents, such as dioxane or diglyme. Furthermore, the rearrangement of 2,3-dimethyl-2,3epoxybutane to 3,3-dimethyl-2-butanone was studied. Results obtained from these separate reaction steps were used in the one-step conversion.

1. Introduction Since the pioneering work of Enichem workers on titanium silicalite (TS-1)[1] there has been great interest in the use of redox-molecular sieves as oxidation catalysts for the synthesis of fine chemicals. I2'31 TS-1 is a highly efficient catalyst for the oxidation of a large number of substrates such as amines, m alcohols, tSl paraffins t6~ or olefins I71 with H202 (30 wt% aqueous) under mild reaction conditions. The high reactivity of TS-1 as an epoxidation catalyst has been attributed to the greater hydrophobic character of the internal surface compared to other molecular sieves, allowing a more facile diffusion of the olefin into the pore channels. R--CHO

H202 R--CH=CH2

~

O / \ R--CH--CH2

[Ti(IV)]

R'OH ~ H20

OR' OH I I R--CH--CH2 (R ' = H or alkyl)

R--CH2CHO + R--C(O)CH3 Scheme 1

568 However, the use of TS-1 is limited due to its rather small pore size, i.e. 5.6 x 5.3 A, [7'8] SO that only linear olefins are efficiently epoxidized. For this reason, a large pore molecular sieve, i.e. Ti-A1-Beta, was developed by Camblor et al. tg'~~ Ti-AI-Beta was much more active than TS-1 for the epoxidation of branched olefins, due to its larger pore size (7.6 x 6.4 ,~).[11,12] However, owing to the Bronsted acidity (due to the presence of framework AI) of Ti-A1-Beta, the main products were those resulting from acid-catalyzed ring-opening reactions, i.e. diol and diol monoethers, or rearrangement products (Scheme 1). t12'~3'~4~Sato et al. showed that these ring-opening reactions are suppressed by neutralization of the Bronsted acid sites by ionexchange with Li+ or Na+.t~4j In contrast to TS-1, the presence of A1 and a large amount of SiOH groups give the pore channels ofTi-A1-Beta a more hydrophilic character. Interestingly, in the case of TS-1 a large solvent effect was found during the epoxidation of propylene, i.e. protic polar solvents such as MeOH largely enhanced the reactivity. ~15'16jIt was suggested that this is due to formation of a hydroperoxo species I, in which the protic alcohol ROH (R = alkyl) coordinates to the Ti site and stabilizes the complex via additional H bonding (Scheme 2). [17]

SiO

O

SiO

.O_--H

\/

O~O~H -

SiO~Ti~ SiO

SiO/'

/ R

I

--_~'N? - H

II Scheme 2

An analogous solvent effect was observed in the Ti-A1-Beta catalyzed epoxidation of 1hexene with H202, with an even higher reactivity being observed in MeCN than in MeOH. rlsl This enhanced reactivity was explained by the formation of Ti-peroxo species II in which water, instead of an alcohol, coordinates to the Ti site (Scheme 2). Furthermore, MeCN is able to coordinate to the Bronsted acidic sites of the zeolite, thus preventing ring-opening reactions &the epoxide to glycol and glycol derivatives (Scheme 1). tlsl Recently, advantage was taken of the fact that several Ti-molecular sieves contain both Lewis as well as Bronsted acidic properties. For example, Corma et al. showed that Ti-MCM41 and Ti-A1-Beta are able to catalyze the epoxidation of linalool in the presence of TBHP followed by ring-closure to cyclic hydroxy ethers in a one-pot reaction, t191 Neri et al. found that TS-1 catalyzes the epoxidation of styrene, with aqueous H202 in MeOH, with in-situ rearrangement to the corresponding aldehyde in 75% yield within 4h. [2~ Based on the above mentioned literature examples, we were especially interested in the use of bifunctional catalysts based on redox-molecular sieves, such as Ti-A1-Beta, 111'12] TAPSO [2H

569 or Ti-MCM-41, [22'23] in the one-step conversion of an olefin to the corresponding aldehyde or ketone. Such processes have industrial potential based on the simplicity of a one-pot procedure, its salt-flee nature and the facile recovery and recycling of the zeolite. In this way several pharmaceutical, flavor and fragrance intermediates can, in principle, be prepared from cheap raw materials. In this paper we report on the use of Ti-Al-Beta as a bifunctional catalyst in the one-step conversion of an olefin to an aldehyde or ketone. As a model system the Ti-AlBeta catalyzed oxidation of 2,3-dimethyl-2-butene 1 with H202 (30 wt% aqueous) to 2,3dimethyl-2,3-epoxybutane 2 (or 2,3-dimethyl-2,3-butanediol, pinacol, 3), followed by subsequent rearrangement to 3,3-dimethyl-2-butanone 4 (pinacolone) was studied (Scheme 3). 0 ~_~

[Ti(IV)]u202 ~

HO

OH

~--~---~

1

0 -H20~

2

3

[[

4

Scheme 3 Initially, the separate reaction steps were investigated, i.e. the epoxidation of olefin 1 via epoxide 2 to diol 3 (Section 2.1.), as well as the rearrangement of epoxide 2 to ketone 4 (Section 2.2.). Results obtained from these investigations were used in the one-step conversion of the olefin 1 to the ketone 4 (Section 2.3.). Research is currently in progress to broaden the scope to other substrates, such as styrene, ~-methylstyrene or ~-pinene as well as the use of other redox-molecular sieves such as TAPSO t21] or Ti-MCM-4 1.t22'23] 2. R e s u l t s and D i s c u s s i o n 2.1. Epoxidation of olefin to diol

The Ti-AI-Beta catalyzed epoxidation of 2,3-dimethyl-2-butene 1 with H 2 0 2 (30 wt% aqueous; molar ratio 1 : H 2 0 2 - 2.4 : 1), via the epoxide 2 to pinacol 3 (Scheme 3), using different types of solvents, i.e. methyl-tert-butylether (MTBE), acetone, t-BuOH, i-PrOH, diethyleneglycoldimethylether (diglyme) and 1,4-dioxane was studied in detail (Table 1). It is important to note that no epoxide 2 was detected by GC during the reaction, indicating that 2 is converted immediately to pinacol 3 over the strong Bronsted acidic sites of Ti-A1Beta. Furthermore, the results presented in Figure 1 and Table 1 show that the nature of the solvent has a large influence on the reactivity of the epoxidation, resulting in different yields of pinacol 3 aider 24h. These solvent differences can be explained by competitive sorption of both solvent and reagents in the zeolite pore channels. MTBE and acetone gave a very low yield of pinacol 3 after 24h of 8% and 21%, respectively. The low reactivity in the case of MTBE can be explained by mass transfer problems due to a liquid two-phase system. In the case of the alcoholic solvents t-BuOH and i-PrOH an increased reactivity was found, which resulted in a

570 better yield of 3 atter 24h of 48% and 66%, respectively. This higher reactivity when using protic polar solvents, which is in agreement with literature reports, tlsl is explained by preferential formation of complex I as the catalytic active species (Scheme 2). However, we found a rather low H202 selectivity in the case of both solvents. Table 1. Epoxidation of 2,3-dimethyl-2-butene 1 to pinacol 3. a) Entry

Solvent

H__2_O_Q2 conv. (%)b) (at t = 6h)

Yield of 3 (%)c) (at t = 6h)

H202 select, d) (at t = 6h)

TON c) (at t = 6h)

1.

M T B E f'g)

-

6

-

42

2. 3. 4. 5. 6.

acetone g) t-BuOH i-PrOH diglyme dioxane

18 44 40 19 25

10 13 28 19 25

56 30 70 100 100

68 90 197 133 172

a) 2,3-dimethyl-2-butene 1 (3.24 g; 38.5 mmol), Ti-AI-Beta (50 mg; 0.02276 mmol Ti), H202 (30 wt% aqueous; 1.8 g; 15.88 mmol), internal standard 1,3,5-tri-tert-butylbenzene, 65 ~ 20 mL of solvent, one-pot procedure, b) Determined by iodometric titration, c) Determined by capillary GC; based on added amount of H202. d) Yield of 3/H202 -conv. x 100. e) mmol of 3/mmol Ti. f) two-phase system; H202 conversion not determined, g) 50 ~ reaction temperature. lOO o

90

;~

80

+

dioxane +

diglyme

---tr-t_BuOH

~ acetone

= MTBE

"-- i-PrOH

70 60 50 40 30 20 10 I

I

10

20

Figure 1

!

Time (h)

571 Interestingly, diglyme and even better dioxane, resulted in a reasonable yield after 24h of 49% and 62%, respectively, as well as a 100% H202 selectivity, i.e. no other products were identified by GC beside the desired product. No reaction took place under analogous experimental conditions without catalyst. From the results presented above can be concluded that the Ti-A1-Beta catalyzed epoxidation of olefin 1 to diol 3 can best be performed in a polar aprotic solvent, such as dioxane or diglyme, resulting in both a good yield of 3 as well as a high H202 selectivity. 2.2. Rearrangement of epoxide to ketone The direct rearrangement of 2,3-dimethyl-2,3-epoxybutane 2 to pinacolone 4 using several catalysts, i.e. TS-1 (1.13 wt% Ti; Si/Ti = 75.2), t241Al-free Ti-Beta (1.7 wt% Ti; Si/Ti = 59), I251 Beta (Si/Al = 12.5), [261 amorphous silica-alumina and Ti-AI-Beta (2.18 wt% Ti; Si/Ti = 44.3; Si/Al = 72.1) ll~ in benzene (50 mL) at 45 ~ was investigated in detail (Scheme 3). The results are presented in Table 2. Table 2. Rearrangement of epoxide 2 to pinacolone 4. a) Entry 1. 2. 3. 4. 5.

Catalyst

Time (h) b)

Conv of 2 (%)c)

Yield of 4 (%)~

TS-1 Ti-Beta Beta Silica-Alumina Ti-Al-Beta

0-24 0-24 4 4 4

0 0 100 100 100

0 0 49 a) 52 a) 39 a)

a) Epoxide 2 (5 mmol), catalyst (80 mg; 0.0364 mmol Ti; 16 wt%), benzene (50 mL), internal standard 1,3,5-tri-tert-butylbenzene, 45 ~ one-pot procedure, b) Time at which a sample is taken from the reaction mixture, c) Determined by capillary GC. d) Side products (identified by GC-MS) were pinacol 3 and allylic alcohol CH2=C(Me)C(Me)2OH. With TS-1 (entry 1) and N-free Ti-Beta (entry 2) no reaction was observed, even after 24h. However, Beta (entry 3), amorphous silica-alumina (entry 4) as well as Ti-A1-Beta (entry 5) gave a moderate yield (39-52%) of 4. Besides pinacolone 4, pinacol 3 and allylic alcohol CHz=C(Me)C(Me)2OH were also formed. Ongoing studies show that the allylic alcohol can be further dehydrated to the corresponding diene in the presence of the acidic sites of a zeolite. [2vl Because the above mentioned experiments gave pinacol 3 in addition to pinacolone 4, attention was focused on the rearrangement of 3 to 4 over the various catalysts. In order to promote the rearrangement, water was removed from the reaction medium using a Dean-Stark apparatus. The results are presented in Table 3. TS-1 (entry 1) and Al-free Ti-Beta (entry 2) again gave no product formation within 24h. Beta gave after 24h a 90% conversion of 3 and a 61% yield of 4 (entry 3). Amorphous silica-alumina (entry 4) resulted in a rather poor conversion of 3 (26%) and yield of 4 (5%). Interestingly, Ti-AI-Beta gave after 4h a 100%

572 conversion of 3 and a 62% yield of 4, indicating that the Bronsted acidic properties of Ti-AlBeta are sufficient to catalyze the rearrangement of pinacol to pinacolone (entry 5). Table 3. Rearrangement of pinacol 3 to pinacolone 4. a~ Catalyst

Time (h) b)

TS-1 Ti-Beta Beta Silica-Alumina Ti-Al-Beta

0-24 0-24 24 24 4

Entry 1. 2. 3. 4. 5.

Conv. of 3 (%)c~ Yield of 4 (%)c~ 0 0 90 26 100

0 0 61 5 d)

62 d>

a) Pinacol 3 (5 mmol), catalyst (80 mg; 0.0364 mmol Ti; 16 wt%), benzene (50 mL), 80 ~ internal standard 1,3,5-tri-tert-butyl-benzene, Dean-Stark apparatus, b) Time at which a sample is taken from the reaction mixture, c) Determined by capillary GC. d) Side product (identified by GC-MS) was allylic alcohol CH2=C(Me)C(Me)2OH. From Table 2 (entry 5) and Table 3 (entry 5) it is apparent that the Ti-Al-Beta catalyzed pinacol rearrangement results in a higher yield of pinacolone 4 than the direct rearrangement starting from the epoxide. Consequently, in order to promote the rearrangement of epoxide 2 to pinacolone 4 via pinacol 3, 2 equivalents of water (relative to the amount of epoxide) were added at the start of the reaction (Scheme 4; conditions A). Furthermore, after formation of the diol a Dean-Stark apparatus was used to remove the excess water (conditions B). O

Y<

HO

OH

A

2

O

B 3

II 4

Conditions A: epoxide (5 mmol), I-I20(10 mmol), Ti-A1-Beta (16 wt%) internal standard, solvent (50 mL), 45 ~ Conditions B: Dean-Stark; 80 ~ Scheme 4 Interestingly, a large solvent effect was found. Benzene as solvent gave, after 24h, a 100% conversion of 2 and a 75% yield of 4. The use of 1,2-dichloroethane gave after 30 minutes already a complete conversion of 2 and a 90% yield of diol 3 was found (Figure 2, conditions A). Subsequent removal of the formed water after 3h using a Dean-Stark apparatus resulted in a 100% selective rearrangement of 3 to pinacolone 4, as determined by GC (Figure 2, conditions B). The overall reaction resulted after 4h reaction time in a 100% conversion of 2 and a 92% yield of pinacolone 4. The blank reaction performed under analogous experimental conditions without catalyst showed no reaction.

573

100 -'~

90

~

-~

80 70

yield of 3 60

"-- yield of 4

50 40 30 20 10 0

m

i

6

8

n

0

2

4

10

Time (h) Figure 2 The reaction performed in a benzene-MeOH mixture (in a 1 : 1 ratio) resulted in the direct formation of the monomethylether, CH3CH(OMe)CH(Me)OH, which did not react further to pinacolone 4. The possible acetal formation between the diol and the ketone in the presence of the Bronsted acidic sites of Ti-A1-Beta was tested by the reaction of equimolar amounts of 3 and 4 under the optimized experimental conditions (1,2-dichloroethane, Dean-Stark, 80 ~ However, no acetal was formed under these conditions.

2.3. Epoxidation with in-situ rearrangement Finally, the one-pot procedure for the Ti-AI-Beta catalyzed epoxidation of olefin 1 with H202 (30 wt% aqueous) to pinacol 3, followed by in-situ rearrangement to pinacolone 4 was studied (Scheme 3). From the results presented in Section 2.1. and 2.2. the logical choice for a solvent system would be dioxane for the epoxidation step of 1 to 3, followed by addition of 1,2-dichloroethane to enforce the rearrangement of 3 to 4. This resulted for the epoxidation step performed in dioxane in a 62% yield of diol 3 after 24h and more than 90% yield of 3 after 48h. However, the expected in-situ rearrangement of 3 to 4 was very slow, i.e. after 4h 25% pinacolone was obtained beside 63% of diol. As the current solvent combination of dioxane with 1,2-dichloroethane is not completely suitable for the performance requested, we are currently testing other polar aprotic solvent systems. Preliminary experiments showed that promising results are obtained when using sulfolane (tetrahydrothiophen-l,l-dioxide) in combination with 1,2-dichloroethane in the epoxidation followed by in-situ rearrangement.

574

3. Experimental 3.1. Analysis The prepared catalysts were analyzed by X-ray powder diffraction (XRD), Diffuse Reflectance Spectra (DREAS), Inductive coupled plasma-atomic emission spectrometric analysis (ICP-AES) and scanning electron microscopy (SEM). The catalysts were characterized by XRD before and after calcination, using a Philips PW 1877 automated powder diffractometer with CuI~ radiation. DREAS measurements were recorded on a Varian Cary-1 spectrophotometer using BaSO4 as a reference. ICP-AES measurements were recorded on a Perkin-Elmer Plasma 40 (ICP) or Perkin-Elmer 1100 (AES). SEM measurements were recorded on a Philips XL-20 microscop.

3.2. Catalyst preparation Zeolites beta (Si/AI = 12.5) t2~ as well as TS-1 (1.13 wt% Ti; Si/Ti = 75.2) [24]were prepared according to literature procedures. Amorphous silica-alumina (HA-HPV; Si/A1 = 2.67) was purchased from AKZO and aluminium-free Ti-Beta was kindly donated by J. C. van der Waal. t151 Ti-A1-Beta was prepared by a slightly modified procedure, t28~ For this preparation TEAOH (freshly prepared; 28% aqueous solution; 80 g) was diluted with the required amount of H/O (18 g). Then TEOT (Aldrich; 1.0 mL) was slowly added at 25 ~ exactly (controlled with a warm water-bath). The solution (which is clear light yellow) was stirred for 2h at 25 ~ Then SiO/(Aerosil Degussa; 16.4 g; 0.278 mol) was added at 25 ~ and the mixture was stirred for 30 minutes at this temperature. Finally A1NO3.9H20 (Merck; 0.56 g; 0.00139 tool) was added at 25 ~ and the mixture was stirred for another 30 minutes at this temperature. A fluid translucent slurry was obtained. Then the solution was transferred into three 50 mL teflon-insert autoclaves at 135 ~ and stirred (60 rpm) for 12 days at this temperature. After crystallization the solid was washed three times with 100 mL of demi-water until pH = 9 and dried at 80 ~ in a vacuum-oven. Gel-composition: OH/(Si + Ti) = 0.54; (Si + Ti)/A1 = 200; H20/(Si + Ti)= 15; Ti/(Si + Ti) = 0.016. Crystalline material was obtained both before and after calcination as determined by XRD measurements. ICP measurements showed a Si/Ti and Si/A1 ratio of 44.3 and 72.1, respectively, corresponding to a 2.18 wt% Ti incorporation. DREAS measurements showed the presence of an absorption band at 48.000 cm -1, indicating that Ti(IV) was tetrahedrally incorporated into the framework of the zeolite, f291No absorption bands of TiO2 species were observed. SEM measurements showed an average crystal size of about 0.25 ~tm.

3.3. Catalytic Reactions 3.3.1. General The catalysts were pre-activated before use at 450 ~ The olefin 2,3-dimethyl-2-butene (purchased from ACROS, 95%) was distilled before use. H202 (30 wt% aqueous) was purchased from Merck, 1,3,5-tri-tert-butylbenzeen from Fluka (>97%), pinacol (2,3-dimethyl2,3-butanediol) from Aldrich (98%), pinacolone (3,3-dimethyl-2-butanone) from ACROS

575 (95%). The from Baker, mixture and were made standard.

solvents dioxane from Merck, t-BuOH, i-PrOH, benzene and 1,2-dichloroethane diglyme from ACROS. During the reaction samples were taken from the reaction analyzed by capillary GC (CP Wax 52 CB) and GC-MS. GC Calibration curves with mixtures of authentic samples and 1,3,5-tri-tert-butylbenzene as internal

3.3.2. Rearrangement

of 2,3-dimethyl-2,3-epoxybutane

(2) via pinacol (3) to

pinacolone (4) In a three-necked flask equipped with reflux-condenser and magnetic stirrer, 2,3-dimethyl2,3-epoxybutane 2 (5 mmol), solvent (50 mL), demi-water (10 mmol) and an exact amount of internal standard 1,3,5-tri-tert-butylbenzene were mixed and, under stirring, the solution was warmed to 45 ~ The catalyst was added at once (t = 0). The reaction was followed in time by taking samples during the reaction and the products were analyzed by GC and GC-MS. After complete conversion of the starting substrate to pinacol 3, the reflux-condenser was replaced by a Dean-Stark apparatus and the reaction temperature raised to 80 ~ The formation of 4 was followed by taking samples of the reaction mixture followed by GC analysis.

3.3.3. Epoxidation of 2,3-dimethyi-2-butene (1) via epoxide (2) to pinacol (3) In a three-necked flask equipped with reflux-condensor and magnetic stirrer, the olefin 2,3dimethyl-2-butene 1 (3.24 g; 38.5 mmol), solvent (20 mL), H202 (30 wt% aqueous; 1.8 g; 15.88 mmol) and an exact amount of internal standard 1,3,5-tri-tert-butylbenzene were mixed and, under stirring, the solution was warmed to the required reaction temperature (Table 3). The catalyst (50 rag) was added at once (t = 0) and the reaction mixture was stirred vigorously during the whole reaction time. The formation of 3 was followed by taking samples of the reaction mixture followed by GC analysis. 4. A c k n o w l e d g e m e n t We wish to thank the Dutch Innovative Research Programme on Catalysis (IOP) for their financial support. 5. References 1. 2.

3. 4. 5. 6. 7. 8.

M. Taramasso, G. Perego, B. Notari US Pat. 4,410,510 (1983). R. A. Sheldon in: Heterogeneous Catalysis and Fine Chemicals II, M. Guisnet, J. Barrault, C. Bouchoule, D. Duprez, G. P6rot, R. Maurel, C. Montassier (Eds.); Stud. Surf Sci. Catal. 59 (1991) 33. R.A. Sheldon, dr. MoL CataL 107 (1996) 75. S. Tonti, P. Raffia, A. Cesana, M. A. Montegazza, M. Padovan, Eur. Pat. 3141475 (1988). A. Esposito, C. Neri, F. Bunomo, US Pat. 4,480,135 (1984). D.R. C Huybrechts, L. De Bruycker, P. A. Jacobs, Nature (1990) 240. C. Neri, A. Esposito, B. Anfossi, F. Buonomo, Eur. Pat. 100,119 (1984). T. Tatsumi, M. Nakamura, K. Yuasa, H. Tominaga, CataL Lett. 10 (1991) 259.

576 9. M.A. Camblor, A. Corrna, J. P&ez-Pariente, SP Pat. 9,101,798 (1991). 10. M. A. Camblor, A. Corma, A. Martinez, J. P6rez-Padente, J. Chem. Soc., Chem. Commun. 589 (1992). 11. A. Corma, M. A. Camblor, P. Esteve, A. Martinez, J. P6rez-Pariente, J. Catal. 145 (1994) 151. 12. A. Corma, P. Esteve, A. Martinez, S. Valencia, or. Catal. 152 (1995) 18. 13. M. A. Camblor, A. Corma, J. P6rez-Padente, Zeolites 13 (1993), 82. 14. T. Sato, J. Dakka, R. A. Sheldon, Stud. Surf. ScJ. Catal. 84C (1994) 1853. 15. M. G. Clerici, G. Belussi, U. Romano, J. Catal. 129 (1991) 159. 16. M. G. Clerici, P. Ingallina, Jr. Catal. 140 (1993) 71. 17. G. Belussi, A. Carati, M. G. Clerici, G. Maddinelli, R. Millini, Jr. Catal. 133 (1992) 220. 18. A. Corma, P. Esteve, A. Martinez, or. Catal. 161 (1996) 11. 19. A. Corma, M. Iglesias, F. S~inchez, 3'. Chem. Soc., Chem. Commun. (1995) 1635. 20. C. Neri, F. Bunomo, Eur. Pat. 0102097 (1986). 21. A. Tuel, Zeolites 15 (1995), 228. 22. M. A. Corma, M. T. Navarro, J. P6rez-Pariente, Jr. Chem. Soc., Chem. Commun. (1994) 147. 23. P. T. Tanev, M. Chibwe, T. J. Pinnavaia, Nature 386 (1994) 239. 24. A. J. H. P. van der Pol, J. H. C. van Hooff, Appl. Catal. A92 (1992) 93. 25. J. C. van der Waal, P. Lin, M. S. Rigutto, H. van Bekkum, Stud. Surf. Sci. Catal. 105C (1997), 1093. 26. R. L. Wadlinger, G. T. Kerr, E. Rosinski, US Pat. 28,341, Reissued (1975) to Mobil Oil Corporation. 27. J. A. Elings, R. A. Sheldon, to be published. 28. M. A. Camblor, private communication. 29. M. R. Bocutti, K. M. Rao, A. Zecchina, G. Leofanti, G. Petrini, Stud. Surf Sci. Catal. 48 (1989) 133.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

577

Liquid-phase oxidation of cyclohexane to adipic acid catalysed by cobalt containing 13-zeolites I. Belkhir 1., A. Germain 1, F. Fajula 1 and E. Fache 2 1 Laboratoire de Mat~riaux Catalytiques et Catalyse en Chimie Organique, UA/IR-CNRS 5618, ENSCM, 8, Rue de l'Ecole Normale,34296 Montpellier Cedex 5, France. Tel: 33 (0) 4 67.14.43.90 ; Fax: 33(0) 4 67.14.43.49. 2 Rh6ne-Poulenc Industrialisation, CRIT-Carribres, 85, Avenue des Frbres Perret, BP 62, 69192 Saint-Fons Cedex, France.

Abstract Cobalt exchanged 13-zeolites obtained by impregnation and solid state ion exchange and cobalt substituted 13-zeolites obtained by incorporation of cobalt in the synthesis gel were studied towards the oxidation of cyclohexane into adipic acid. The Co-substituted 13-zeolites were found to be effective catalysts for the oxidation of cyclohexane in acetic acid. In contrast, the use of Co-exchanged 13-zeolites always led to inhibition of the oxidation. It was demonstrated that the catalytic activity came as a result of the dissolved cobalt in the reaction medium, while inhibition was ascribed to the accessible uncompensated aluminic sites of the zeolites. 1. I N T R O D U C T I O N Adipic acid is an important intermediate extensively used for the manufacture of nylon 66. It is currently produced from cyclohexane oxidation by a two steps process [1 ]. During the first step, oxidation of cyclohexane by air in the liquid phase forms cyclohexanol and cyclohexanone. Further oxidation of this mixture by nitric acid gives adipic acid. In addition to its cost, the use of nitric acid generates corrosion risks and requires recovery of the nitrogen oxides effluents. Direct aerial oxidation of cyclohexane in a single step implies a partial and selective oxidation of the substrate. Oxidation without catalyst but in the presence of initiator gives adipic acid as a minor product [2]. Homogeneous catalysis by cobalt acetate in acetic acid provides good selectivity for adipic acid (88% at 21% conversion) [3, 4]. Recently, solid CoAPO were found to be effective heterogeneous catalysts [5]. However, the adipic acid selectivities were low [5-7] and the heterogeneous nature of the catalysis was not clear [8]. Moreover, redox properties of the framework cobalt ions are now subject to debate [9-10] and the reactive redox process could be attributed to non-framework cobalt species. Thus we have decided to explore the catalytic activity of cobalt containing zeolites and compare cobalt exchanged (or impregnated) zeolites to cobalt substituted zeolites obtained by incorporation

578 of cobalt into the zeolite synthesis gel as it was already achieved to obtain cobalt silicalite [11-13]. In order to favour adsorption of the organic apolar substrate and subsequent desorption and diffusion out of the catalyst of the polar products, the zeolite must possess both surface hydrophobicity and an open large pored structure. Taking into account these reasons, high silica 13-zeolites were chosen. Since 13-zeolites cannot be obtained without trivalent metal, cobalt substituted 13-zeolites were synthetised in the presence of aluminium or boron. The aim of the present work is to investigate and compare the cyclohexane oxidation activities of cobalt exchanged zeolites prepared by conventional impregnation or solid-state ion-exchange methods and cobalt substituted zeolites, in order to gain insight into the type of catalysis involved. Herein, the results obtained for oxidation of cyclohexane to adipic acid catalysed by cobalt exchanged zeolites (Co/BEA) and cobalt substituted 13-zeolites (Co-BEA) are presented. 2. E X P E R I M E N T A L BEA stands for 13-zeolites and the numbers after the structure type code of zeolites denote Si/AI or Si/B ratio (determined by analysis).

2.1. Materials Cobalt (II) acetate tetrahydrate, Cobalt (II) nitrate hexahydrate, cobalt chloride, sodium aluminate (Na20.A1203.3H20), boric acid (H3BO3) and tetraethylammonium hydroxide (Aldrich), acetic acid purex and cyclohexane for analysis (SDS) were used as received. Ludox HS-40 colloidal silica solution was obtained from Dupont. Zeolites BEA 15 and BEA 27 were synthesized in the presence of tetraethylammonium hydroxide (TEAOH) according to the procedure described by Wadlinger and al. [ 14]. Dealuminated BEA 1100 was obtained by treating BEA 15 with concentrated nitric acid [ 15].

2.2. Catalysts Zeolite impregnation: The zeolites were impregnated with 1 to 2% Co using a cobalt (II) acetate-water solution [ 16]. After evaporation until dryness at 343 K, the solids were calcined at 823 K for 6 hours (Co/BEA 15, Co/BEA 27 and Co/BEA 1100). Solid state exchange: Mechanical mixtures of powders of the zeolite and COC12 were grounded and calcined in air at 823 K for 6 hours [ 17] (Co/BEA 15S). Cobalt substituted zeolites synthesis: A: 0.97 g of Co(NO3)2.6H20 was dissolved in 13 cm 3 distilled water. Next, 0.55 g of sodium aluminate was added to this solution. Then, 48 g of Ludox HS-40 were dissolved in the mixture. A second solution was prepared by dissolving 0.56 g of sodium hydroxide in 36.8 g of a 40% aqueous solution of tetraethylammonium hydroxide. The final gel composition was 10Na20.CoO.AI203.110SIO2.1170H20. After 4 hours of stirring, the gel solutions were transfered into autoclaves and crystallised at 403 K for various periods from 1 to 2 weeks (Co-AI-BEA). B: 0.57 g of H3BO3 was added to 32.5 g of a 35% aqueous solution of tetraethylammonium hydroxide. Then, 0.27 g of Co(NO3)2.6H20 were dissolved in this solution. Next, 25 g of Ludox HS-40 were added to the mixture. The final gel composition was 3Na20.CoO.5B203.190SIO2.2170H20. After 4

579 hours of stirring, the gel solutions were transfered into autoclaves and crystallised at 423 K for various periods from 2 to 3 weeks (Co-B-BEA). Characterization of catalysts: The zeolite structure was checked by X-ray diffraction patterns recorded on a CGR Theta 60 instrument using Cu Kal filtered radiation. The chemical composition of the catalysts was determined by atomic absorption analysis after dissolution of the sample (SCA-CNRS, Solaize, France). Micropore volumes were measured by N2 adsorption at 77 K using a Micromeritics ASAP 2000 apparatus and by adsorption of cyclohexane (at P/Po=0.15) using a microbalance apparatus SETARAM SF 85. Incorporation of tetrahedral cobalt (II) in the framework of Co-A1-BEA and Co-B-BEA was confirmed by electronic spectroscopy [ 18] using a Perkin Elmer Lambda 14 UV-visible diffuse reflectance spectrophotometer. Acidity measurements were performed by Fourier transform infrared spectroscopy (FT-IR, Nicolet FTIR 320) after pyridine adsorption. Self-supported wafer of pure zeolite (20 mg/cm 2) was outgassed at 673 K for 6 hours at a pressure of 10 -1 Pa. After cooling at 423 K, the zeolite was saturated with pyridine vapour (30 kPa) for 5 min, evacuated at this temperature for 30 min and the IR spectrum was recorded.

2.3. Procedure Cyclohexane oxidation was carried out in a 300 cm 3 titanium, semi-batch, mechanically stirred Parr-type reactor. A typical procedure used for the oxidation was described in detail for an experiment at 383 K and 21 bars of total pressure. The reaction feed consisted of cyclohexane (45 cm3; 690 mmol), glacial acetic acid (68 cm3), catalyst (0.5 to 3 g) and acetaldehyde (0.24 g; 5 mmol) used as promoter. The autoclave was brought to the operating temperature and pressure, then held there for 3 hours under a constant flow of 20 dm3.h 1 of oxygen and nitrogen (10/90). Oxygen consumption was followed by the measure of the oxygen concentration and the flow rate in the output. The reactor was cooled and depressurized, and the product mixture was removed. The reaction mixture (2 g) was esterified at reflux with methanol (15 cm 3) in the presence of 2 drops of concentrated H2SO 4 to obtain the diacids in the diesters form. The products were analysed using a Hewlett Packard gas chromatograph equipped with a Carbowax 52 CB polar capillary column and a flame ionization detector assembled with a Shimadzu programmed and computerized Chromatopac CR6A. The reaction products consisted of adipic, glutaric, succinic and 6-hydroxycaproic acids, cyclohexanone, cyclohexanol and butyrolactone. Filtrates of Co-containing zeolites were obtained by treatment of 1 to 2g of zeolites in 75 cm 3 acetic acid at reflux overnight. After centrifugation of the solid, the desired amount of filtrate was fed into the reaction system. 3. R E S U L T S

AND DISCUSSION

3.1. Cyclohexane oxidation catalysed by cobalt exchanged zeolites We observed that the aerial oxidation of cyclohexane without catalyst, but in the presence of initiator (acetaldehyde) and in acetic acid as a solvent, occurred at 110~ The first step of the mechanism was the formation of the cyclohexylhydropero• which was converted to cyclohexanol and cyclohexanone. As cyclohexanone catalysed the decomposition of the hydroperoxide, the oxidation was autocatalytic.

580 Table 1 summarizes the activity of aluminic and boric zeolites in the proton form in the oxidation of cyclohexane. The results show that the reaction rate was reduced in the presence of aluminic zeolites. The addition of aluminic H-zeolites thus inhibited the oxidation of cyclohexane in acetic acid and the inhibition effect was stronger the larger the amount of aluminic sites was. Thus, the inhibition of the uncatalysed oxidation is attributable to the presence of the strong Br6nsted acid sites of the zeolites. It might come from the proton assisted heterolytic decomposition of the cyclohexyl hydroperoxide which is an intermediate in the autoxidation of cyclohexane [19]. Such a decomposition in acetic acid leads to the formation of 6-acetyloxy-hexanal [20] due to an ionic mechanism. Later, this aldehyde is oxidised into the acid analogue, leading to 6-hydroxycaproic acid after methanolysis. This reaction contributes to the termination of the free radical chain mechanism. In contrast to aluminic zeolites, the use of a boric H-zeolite led to an activity level equal to that of the uncatalysed reaction. The oxidation of cyclohexane was thus not inhibited by boric Hzeolites. It has been showed that boron atoms incorporated into the framework of 13-zeolites during synthesis, were removed from the solid after activation [21]. This could explain the very weak acidity presented by the solid. Table 1 Catalytic activity of zeolites in the proton form in the oxidation of cyclohexane. Zeolite

Amount (g)

None

Acid sites * Reaction rate*** (mmol)

(mmol/min)

Cyclohexane

Adipic acid

conversion **

yield **

(%)

(%)

0

0

0.36

6.6

1.1

H-A1-BEA 15

1.96

1.74

-~0

0

0

H-A1-BEA 1100

1.01

0.021

0.30

6.1

0.85

H-B-BEA 15

1.0

0.96

0.38

7.7

1.5

Cyclohexane: 690 mmol; acetaldehyde: 5 mmol; acetic acid: 68 cm3; N2/O2 90/10; 21 bars; flow: 20 dm3.hl; 110~ * Overall aluminium (or boron) content in the reaction medium. ** The reaction lasted for 3 hours. *** Rate of oxygen consumption measured after 2 hours of reaction. The use of Co-exchanged zeolites always led to an activity level below that of the uncatalysed reaction (Table 2). This was true for the impregnated Co/BEA and for the Co/BEA prepared by solid state ion exchange. Taking into account the amount of cobalt cations and considering that each cation compensates two negative charges of the framework, the number of residual (noncompensated) acid sites was calculated and the activity of samples was plotted as a function of the latter in Figure 1. It was observed that all the Co/BEA zeolites exhibited the same behaviour as the aluminic zeolites in the proton form whatever the exchange method used. The oxidation activity was thus a decreasing function of the number of aluminic acid sites. In order to understand the mechanism occuring during the oxidation, the Co-exchanged zeolites were treated in acetic acid at reflux overnight. After centrifugation of the solid, the filtrate

581 was fed into reaction. The results of the activities of Co-exchanged zeolites filtrates are reported in Table 3 and in Figure 1. Table 2 Catalytic activity of Co-exchanged zeolites in the oxidation of cyclohexane. Catalyst

Co content* Non exchanged Reaction rate (mmol)

acid sites

(mmol/min)

Cyclohexane

Adipic acid

conversion

yield

(%)

(%) 0.7

(mmol) Co/BEA 15

0.84

0.52

0.10

5.6

Co/BEA 27

0.26

0.11

0.20

6.0

0.6

Co/BEA 1100

0.60

0

0.34

7.0

0.87

Co/BEA 15S** 0.42 1.03 0.04 0.7 0 Same conditions as Table 1. * Total cobalt content in the reaction medium. ** Obtained by solid state ion exchange.

0.5

~L

LL

r-

E

0.4

o

E E

0.3

tO

o

0.2

0 L_

L_

> 0 0

0.1

0

0.3

0.6

0.9

1.2

1.5

1.8

Non compensated aluminic sites (mmol) FIG. 1 Dependency on the rate of oxygen consumption as a function of the free aluminic sites of zeolites present in the reaction medium (N none, II H/BEA, 5 Co/BEA, A Co/BEA filtrates linked to the corresponding Co/BEA, O Co-BEA)

582 The filtrates exhibited effective catalytic activity higher than the solids. This confirmed that the inhibition was due to the zeolitic aluminic sites. Thus, the catalytic activity of the filtrates must be attributed to the dissolved cobalt. We also observed that the addition of an aluminic 13-zeolite (Si/AI=I 5) in the proton form to an active filtrate inhibited the oxidation reaction. All these results demonstrate that the catalysis is homogeneous and that the zeolitic aluminic sites are responsible for the inhibition. If the Co-exchanged zeolites are not catalysts in the oxidation of cyclohexane, it is due to the presence of uncompensated aluminic sites in the solids. Table 3 Catalytic activity of the Co-exchanged zeolite filtrates in the oxidation of cyclohexane. Filtrate

Zeolite cobalt

Reaction rate

Cyclohexane

Adipic acid

content

(mmol/min)

conversion

yield

(%)

(%)

9.8

2.5

(mmol) Co/BEA 15

0.66

0.62

Co/BEA 15S

0.42

0.60

9.2

1.8

Co/BEA 1100

0.60

0.62

10.1

1.9

Co/BEA 15+H-AI-BEA 15

0.66

0

0

0

Same conditions as Table 1.

3.2. Catalytic activity of cobalt substituted zeolites The Co-exchanged zeolites were not effective catalysts for the oxidation of cyclohexane. The cobalt exchanged ions were not stabilized enough by the zeolite interactions and part of these cations were released in the oxidation medium. Thus, we decided to explore the activity of 13-zeolites in which cobalt ions were incorporated into the framework. We hoped that the incorporation would increase the stability of the cation within the solid. We studied the catalytic activities of cobalt substituted 13-zeolites containing aluminium (Co-A1-BEA) and boron (Co-B-BEA) towards the oxidation of cyclohexane into adipic acid. Table 4 Catalytic activity of the Co-substituted 13-zeolites in the oxidation of cyclohexane. Catalyst

Zeolite cobalt

Reaction rate

Cyclohexane

Adipic acid

content

(mmol/min)

conversion

yield

(%)

(%)

0.62

9.1

18

(mmol) Co-AI-BEA

0.27

calcined Co-AI-BEA

0.27

0.64

9.2

19

Co-A1-BEA filtrate

0.29

0.60

8.3

14

Co-B-BEA

0.24

0.66

9.4

25

calcined Co-B-BEA

0.24

0.64

9.3

2.1

Co-B-BEA filtrate

0.24

0.68

11.5

3.4

Same conditions as Table 1.

583 Three types of catalytic experiments were achieved on each Co-substituted p-zeolites. First, the catalytic activity of the as-made zeolite was evaluated. Then, the activity of the calcined one was investigated. Finally, the as-synthetised Co-substituted zeolite was treated in acetic acid at reflux and the filtrate was fed into reaction. The results are reported in Table 4. The boric (as-made or calcined) Co-substituted p-zeolites presented a catalytic activity higher than the activities of the uncatalysed reaction and the reaction with boric zeolite in the proton form. As the boric Co-substituted zeolites were not acid, they did not decrease the reaction rate of the oxidation. The filtrate of the boric Co-substituted zeolite was as active as the solids. This demonstrated that the catalysis resulted from the cobalt in solution. The reaction rates of the aluminic (as-made or calcined) zeolites were higher than the rate of reaction whithout catalyst and they did not follow the inhibition curve of the aluminic Coexchanged zeolites. As the as-made zeolite still contained the templates inside the pores, the solid certainly prevented access of cyclohexane to the aluminic sites responsible for the inhibition, but allowed the dissolution of cobalt. The activity of the as-made Co-substituted zeolite filtrate was similar to the activity of the zeolite. So, as for the Co-exchanged zeolites, the catalysis is homogeneous and came as a consequence of the dissolved cobalt. In contrast to calcined cobalt exchanged p-zeolites, the accessible acid sites of the calcined Cosubstituted aluminic p-zeolite did not inhibit the oxidation of cyclohexane. This result shows that the aluminic sites of the Co-substituted zeolite did not inhibit the oxidation of cyclohexane. The nature of the acidity was investigated in order to explain the catalytic activity of the calcined Co-substituted p-zeolite and the role played by the aluminic sites of this solid. A pyridine adsorption followed by IR spectroscopy measurements was performed on the calcined catalyst. It has been shown that adsorption of pyridine emphasized two distinct bands at 1548 c m -1 and 1451 cm 1 corresponding respectively to the adsorption on Br6nsted and Lewis sites [22]. In the case of the calcined Co-substituted zeolite, only a weak band at 1548 -1 cm appeared in the IR spectrum. Thus, we deduced that very few Br6nsted sites were present in the catalyst. This could explain that the oxidation of cyclohexane into adipic acid in the presence of calcined Co-substituted aluminic p-zeolite was not inhibited. 4. C O N C L U S I O N In the oxidation of cyclohexane into adipic acid, we have shown that the aluminic sites of proton form zeolites inhibited the reaction. When the aluminic sites of a Co-exchanged [3zeolites were not totally compensated, the solids inhibited also the oxidation. The activity was not influenced by the cobalt exchange method. As the acetic acid filtrates of Co-exchanged zeolites presented catalytic activities, we demonstrated that the catalysis is homogeneous and is due to the dissolved cobalt. In this precise case, the cobalt exchanged ions were not sufficiently stabilized by the zeolite and were dissolved in acetic acid. As we expected a better stabilization of the cobalt ions introduced in the zeolite framework, we studied the activity of Co-substituted p-zeolites. They showed an effective catalytic activity towards the oxidation of cyclohexane also linked to the cobalt in solution. We demonstrated that the Cosubstituted p-zeolites were more active than the Co-exchanged zeolites. In both cases, however, when a catalysis occurs, it came as a result of the dissolved cobalt.

584

REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12 13

14 15 16 17 18 19 20 21 22

A. Castellan, J.C.J. Bart and S. Cavallero, Catal. Today 9, 23 7 (1991). D.G. Hendry, C.W. Gould, D. Schuetzle, M.G. Syz and F.R. Mayo, J. Org. Chem. 41, 1 (1976). K. Tanaka, Hydrocarbon Proc. 53, 114 (1974). J. Kollar, W.O. Pat. 94/07833 (1993). S.S. Lin and H.S. Weng, Appl. Catal. A: General 105, 289 (1993); 118, 21 (1994); jr. Chem. Eng. Jpn 27, 211 (1994). B. Kraushaar-Czarnetzki and W.G.M. Hoogervorst, Eur. Pat., 519,569 (1992). D.L. Vanoppen, D.E. de Vos, M.J. Genet, P.G. Rouxhet and P.A. Jacobs, Angew. Chem., Int. Ed. Engl. 34, 560 (1995). B. Kraushaar-Czarnetzki, W.G.M. Hoogervorst and W.H.J. Stork, Stud. Surf Sci. Catal. 84, 1869 (1994). V. Kurshev, L. Kevan, D.J. Parillo, C. Pereira, G.T. Kokotailo and R.J. Gorte, J. Phys. Chem. 98, 10160 (1994). H. Berndt, A. Martin and Y. Zhang, Micropor. Mater. 6, 1 (1996). J.A. Rossin, C. Saldarriaga and M.E. Davis, Zeolites 7, 295 (1987). R. Mostowicz, A.J. Dabrowski and J.M. Jablonski, Stud. Surf Sci. Catal. 49A, 249 (1989). T. Inui, A. Miyamoto, H. Matsuda, H. Nagata, Y. Makino, K. Fukuda and F. Okazumi, New Developments in Zeolite Science and Technology, Proc. 7th Int. Zeolite Conf., Tokyo, 1986, ed. by Y. Murakami and al., 859. R.L. Wadlinger, G.T. Kerr, E.J. Rosinski U.S. Patent 3,308,069 (1967). E. Bourgeat Lami, F. Fajula, D. Anglerot and T. Des Courieres, Micropor. Mater. 1,237 (1993). J.M. Stencel, V.U.S. Rao, J.R. Diehl, K.H. Rhee, A.G. Dhere and R.J. De Angelis, jr. Catal. 84, 109 (1983). A.V. Kucherov, A.A. Slinkin, J. Mol. Catal. 90, 323 (1994). F.A. Cotton, D.M.L Goodgame and M. Goodgame, Jr. Amer. Chem. Soc. 83, 4690 (1961). R.A. Sheldon, J.K. Kochi, in <<Metal-catalysed oxidations of organic compounds>>, Academic Press, New York. 1981. J.C. Brunie, M. Costantini, N. Crenne and M. Jouffret U.S. Patent 3,689,534 (1972). M. Derewinski, F. Fajula, Appl. Catal. A: General 108, 53 (1994). N-Y. Topsoe, K. Pedersen and E.G. Derouane, J. Catal. 70, 41 (1981).

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

585

Nitrogen oxides catalyzed selective oxidation by oxygen in the liquid phase A. B. Levina, S. S. Chornaja, I. A. Grigorjeva, O. N. Sergejeva, S. R. Trusov Riga Technical University, Faculty of Chemical Technology, Azenes Sr., 14/24, L V - 1048, Riga, Latvia Different oxidative transformations in acidic water solutions with the participation of molecular oxygen may be carried out in the presence of nitrogen oxides. Primary and secondary aliphatic alcohols, as well as benzyl alcohol and dibenzyl ether, allyl alcohol and glucose are oxidised to the corresponding carbonyl compounds ( glucose - to gluconic or glucaric acids). Aromatic compounds are halogenated by means of halogen salts in the presence of transition metals salts. Variable valence metal ions may be easily oxidised to higher valence state under the same conditions. 1. OXIDATION OF ORGANIC COMPOUNDS There are at least three possible ways to carry out alcohol oxidations into carbonyl compounds by oxygen in the liquid phase. The first is oxidation in non-aqueous solutions according to free radical mechanism. But usually aldehydes (ketones) formed are much more easily oxidizable than alcohols and oxidation selectivity is not high enough. The second is oxidation in alkaline water solutions. But in alkaline solutions aldehydes may transform into compounds with higher molecular mass. And the third is oxidation in acidic water solutions. This way seems to be free of complications mentioned above. The only problem is the selection of an adequate catalytic system.

1.1 Benzyl alcohol In [ 1,2] it was shown that in 5-10 M H2SO4 at room temperature benzyl alcohol may be transformed into benzaldehyde with practically 100% selectivity by the action of nitrous acid. It is mentioned also [2], that in the presence of 02 partial regeneration of oxidant is possible. However, it was impossible to organize the catalytic cycle because of instability of nitrous acid. We have elaborated a suitable preparative method for benzyl alcohol oxidation into benzaldehyde by oxygen in the presence of catalytic amounts of nitrous acid. The oxidations have been carried out in a thermostated gas-measurement apparatus which consisted of a vigorously shaken glass vessel and a gas burette., filled with oxygen. Benzyl alcohol oxidation kinetics was followed by absorbed 02 volume measurement. Previously it was proven that the process proceeds in the kinetic area. The following oxidation process parameters were varied during kinetic investigations: [PhCH2OH]0 = (2.0-8.0)x 10 -2 M, [NaNO2]0 = (1.0-5.0)x 10 -3 M ,[HC104]0 = 5.0-8.3 M, T = 298 - 318 K, Po2 -(0.05 - 1.0)xl05 Pa. The GLC analysis

586 of the reaction mixture shows that the only benzyl alcohol oxidation product is benzaldehyde and the stoichiometric equation of the reaction is: PhCI-hOH + 1/2 0 2 - ~ PhCHO + H20

(I)

During investigations it was established that there are two process parameters areas with different kinetic equations for the oxidation process. In area A ( [HC104]=5.0-6.0 M, P 02 = (0.5-1.0)• 105 Pa. ) the following equation is adequate to describe oxidation process kinetics: WA = k I [PhCH:OH],

(1)

here WA is the benzaldehyde formation rate; k I is the effective rate constant. In area B ( [HC104]=7.3-8.3 M, P o2 = (0.05 - 0.3)• 105 Pa) and the process kinetic equation appears as follows WB = k0,

(2)

here WB is the benzaldehyde formation rate ' k 0 is the effective rate constant. On the basis of kinetic investigations the following process mechanism is proposed: HNO2 + H + cv NO + + H20 PhCH2OH + NO+--~ [PhCH2OH] +" + NO 2NO + O2 --> 2NO2 [PhCH2OH] +~ + NO2 --> [PhCHOH] + + HNO2 [PhCHOH] + --~ PhCHO + H +

(II) (III) (IV)

(v) (vi)

Depending on process conditions, the rate limiting stage of reaction (I) may be either reaction (III) or (IV). At high oxygen partial pressures and "low" solution acidities the rate limiting stage is (III), but at low oxygen partial pressures and "high" acidities - stage (IV). In the more detail some of the aspects of the process kinetics are described in [3]. It was shown that instead of HC104 it is possible to use H2SO4. As it was shown in [5], chloride ion acts as an inhibitor ofbenzyl alcohol oxidation to benzaldehyde in 5.0-6.2 M HC104 water solutions. The higner its inhibitor activity, the higher is solution acidity. It is because of replacement of the most active catalyst form ( NO +) into less active (NOCI). HC1 also may be used, but in combination with sulpholane [4]. Preparatively benzaldehyde synthesis in the framework of this approach may realised [4] as follows. Into a thermostated two space reactor ( a low wall in the reactor prevents reagent's mixing before the reactor is shaken ) with the volume 75 ml in one part of reactor, are placed : 9.0 ml 8.6 M H2SO4,5.0 ml hexane or heptane and 1.04 g benzyl alcohol. To another part add 0.4 ml 2 M NaNO2 water solution. Then all of the gas phase is filled with 02 and at room temperature reactor shaking starts. After 45-60 min vigorous shaking process is finished and the organic phase, which contains only benzaldehyde, is separated. The benzaldehyde yield is 90-95%.

1.2. Dibenzyl ether Dibenzyl ether is a side product in some industrial processes and its utilisation is an actual problem. It may be transformed into the mixture of benzaldehyde and benzoic acid by oxida-

587 tion by air at 120-150~ in the presence of metal complexes catalysts [6]. We have shown that the catalytic system mentioned above may be used successfully for dibenzyi ether oxidation into benzaldehyde. Reaction proceeds in 10-15 min at room temperature. Benzaldehyde yield is app. 100% [7]. In contrast to benzyl alcohol, dibenzyl ether is practically insoluble in water That is why in kinetic investigations it is necessary to a use supplementary organic solvent for system homogenization. In table 1 there are the initial rate values of dibenzyl ether oxidation in the presence of some organic solvents. In previous experiments it was established that all solvents mentioned in table 1 do not undergo oxidation themselves in these conditions. Data in table 1 refers to the following oxidation conditions :[Dibenzyl ether]o = 2.63• 10-2M, [NaNO2]o = 3.0x 10-3M, [HCIO4]o =5.36 M, T=298 K, Po2 = 1.0x 105pa. Table 1. Rates of dibenzyl ether oxidation in the presence of different organic solvents ( 3 M ). Solvent -PKBI_pW 0 • 10 4,/147/lxs Sulpholane

12.88

1.30

Dimethylsulphone

12.27

1.10

Dimethylsulphoxide

2.48

0.10

Dimethyl formamide

1.35

0

It is evident that dibenzyl ether oxidation proceeds with remarkable speed only in the presence of the solvents with high pKB/~ . All our research on dibenzyl ether oxidation kinetics was carried out in the presence of sulpholane. Oxidation product analysis by GLC shows that the only dibenzyl ether oxidation product is benzaldehyde. At the same time during the oxidation process small amounts of benzyl alcohol are detected in the reaction media. Benzyl alcohol oxidation rate is five times higher than that for dibenzyl ether. The form of kinetic curves for benzyl alcohol oxidation is typical for intermediate products in consecutive processes. It must be emphasized that the benzyl alcohol formation is not a result of hydrolysis. Furthermore it was established that in early stages of oxidation both products : benzaldehyde and dibenzyl ether are formed simultaneously. All these data obtained proved consecutive dibenzyl ether oxidation followed : PhCH2OCH2Ph --~ PhCHO + PhCH2OH --+ 2PhCHO

(3)

Kinetic investigations of dibenzyl ether oxidation shows that like benzyl alcohol oxidation, there are the same two areas of process parameters with different reaction kinetics: area A with "low" acidity ([HC104]=5.0-5.8 M) and high oxygen partial pressure ( ( 0.5-1.0)• 10SPa), and area B with "high" acidities ([HCIO4] = 5.8-6.6 M) and low oxygen pressures ( ( 0.05-0.5)• 105 Pa). The main kinetic features of oxidation, that is rate dependence on concentrations and temperatures, both for dibenzyl ether and benzaldehyde are one and the same. The mechanism of dibenzyl ether oxidation appears as follows: HNO2 + H + < ~ N O + + H20 PhCH2OCH2Ph + NO + --> [PhCH2OCI-hPh] +" + NO

(II) (VII)

588

(VIII) (IX) (X) (XI)

2NO + 02--~ 2NO2 [PhCI-hOCI-hPh] +" + NO2 r [PhCI-hOCHPh] + + HNO2 [PhCI-hOCHPh] + + HOH ~ [PhCH2-O-CH(OH)Ph] + H + [PhCH2-O-CH(OH)Ph] ~ PhCHO + PhCI-hOH

Benzyl alcohol formed in reaction (XI) oxidises into benzaldehyde according to the mechanism mentioned above ( reactions (II) - (VI)). Kinetic aspects of dibenzyl ether oxidation into benzaldehyde are described in more details in [8]. Preparatively benzaldehyde synthesis according this method may be realised as follows [7]. Into a thermostated two space reactor with the volume 75 ml in one part of reactor, are placed 10.0 ml 10.4 M H2SO4 and 0.792 g dibezyl ether but into another - 0.20 ml 2.00 M NaNO2 solution. Oxidation by oxygen proceeds in 10 min. The reaction mixture is diluted by water three times and extracted by benzene. According to GLC data, the benzene solution contains app. 99% benzaldehyde and traces of benzoic acid. Dibenzyl ether transformed completely.

1.3 Aiiphatic alcohols The oxidation of alcohols has been carried out in the same thermostated gas-measurement apparatus. Concentrated aqueous solutions of HCIO4, I-hSO4 and CF3COOH were used as a reaction medium. The reaction conditions are listed in Table 2. Table 2. Conditions of the aliphatic alcohols oxidation processes. No. Amount of catalyst (NaNO2), mmol 1

0.2

Acid solution HCIO4, 8 M + H20

2

0.1

HCIO4, 8 M + 1-120

3

0.05

HC104, 8 M + 1-120

4

0.1

I-hSO4, 11 M + 1-t20

5

0.1

CF3COOH, 11 M + 1420

For all experiments 10 ml of acidic solution and 1 mmol of alcohol were taken. The reaction temperature was 20 ~C, p o 2 - 1• 105pa. As it was established, in all cases alcohol oxidation proceeds according to the following equation: R1RaCH-OH + 0.502 --> R1R2C=O

+

H20

(4)

In Table 3 there is shown the nature of R1 and l~. T.he experiments showed that allyl alcohol as well as primary and secondary aliphatic alcohols are oxidized by oxygen under conditions mentioned in Table 3. The oxidation proceeds quickly (at 3-5 min). For all studied reactions only the initial alcohol and corresponding carbonyl compound were detected in the reaction mixture by GLC. The yields of carbonyl compounds are in agreement with the expended quantities of oxygen, determined by the gas

589 burette. The data in Tables 2 and 3 show that high conversion levels of primary and secondary aliphatic alcohols are achieved in 8.0 M perchloric acid solution ( column 1 in Table 3). Table 3. The alcohols oxidized and the amounts of oxygen used in reaction (4), percent on stoichiometEr Condition No. from Table 2 Alco R1 R2 hol 1

2

3

4

5

I

CH2=CH -

H-

13.3

14.1

6.6

141

0

II

CH3-(CH2)2-

H-

99

70.6

35.3

22 4

0

III

(CH3)2CH-

H-

79

44.8

20.8

174

0

IV

CH3-

CH3-

92

91.3

72.2

913

68.1

V

CH3-(CH2)2-

CH3-

80.5

80

66.4

63 1

70.6

VI

-(CH2)5-

-

73

56.4

35.7

432

63.1

The conversion level decreases with the amount of catalyst is reduced ( columns 2 and 3). The conversion of iso-butyl alcohol (No. III in Table 3) is always less than for the n-butyl alcohol (IV). For secondary alcohols, the conversion falls with an increase in molecular mass. In 11 M solutions of sulphuric acid (column 4) the primary alcohols (IV and V) are oxidised much less effectively than the secondary alcohols (IV-VI). But the largest selectivity is achieved in 11 M trifluoracetic acid (column 5) : the primary alcohols are not oxidized at all, while secondary alcohols are oxidized fairly well. The satisfactory conversion of allyl alcohol (I) was not achieved in all of the mentioned acidic systems. Thus, the results of the oxidation of various alcohols show that there is a complicated dependence on the structure of the alcohol molecule and conditions of the process. The oxidation mechanism is similar to reactions (II) - (VI). Aspects of the kinetics of aliphatic alcohol oxidation under these conditions are described in more details in [9]. 1.4 G l u c o s e The most desirable products of glucose oxidation are gluconic, glucuronic and glucaric acids. All these products are widely used in the pharmaceutical industry. That is why elaboration of new synthesis methods for these compounds is needed. Furthermore, this approach may be used in other carbohydrates oxidation. The kinetics of glucose oxidation was investigated by the same methods as above. The variable process parameters were : [glucose]0=0.5-1.5 M, [NaNO2] = 0.01-0.25 M, [HC104]=4.0-6.0 M, P02 = (0.2-1.0)• 105pa, T=313-343 K. The mixture of water- sulpholane - perchloric acid was used as a solvent. Both glucose and it's oxidation products amounts in the reaction mixture were measured by GLC. The structure of the products obtained was proven by elements analysis, NMR and GCMS. In the course of the investigations it was established that glucose oxidation proceeds according to following scheme: HOCH2(CHOH)4CHO + 0.5 O2 --> HOCH2(CHOH)4COOH

(5)

590

glucose

HOCI-h(CHOH)4COOH

+ 02 -~

gluconic acid

HOOC(CHOH)4COOH glucaric acid

(6)

The yields of compounds obtained are in agreement with the expended quantities of oxygen, determined by the gas burette. As it was established, glucose oxidation according to (5) does not proceed in the absence of NaNO2. Reaction (5) parameters optimization shows that NaNO2 concentration is the most important factor influencing glucose oxidation rate and selectivity. For example, at [NaNO2]0 = 0.25 M and [glucose]o=l M, [HCIO4]-6.0 M , [sulpholane]=2.6 M, T= 333 K, POE = 1.0x 105pa, in 3 rain glucose is oxidised into gluconic acid entirely. In the next 40-45 rain glucaric acid was formed in 95% yield [ 10]. So, using different NaNO2 concentrations it is possible to obtain either gluconic or glucaric acid, because glucose oxidation is consecutive reaction and the absorbed oxygen volume corresponds to reaction stoichiometry. The results of the kinetic investigations proved that glucose oxidation mechanism is the same as established for other substances oxidized in the presence of this catalytic system. The kinetic equation of gluconic acid formation rate appears as follows: (7)

W0 = k[NaNO2]2 xP 02 Here : k = (0.42_+0.08)x 10 -4 L/MxsxPa at 333 K and [HCIO4]=6.0 M W o is the initial gluconic acid formation rate, M/lx s. The equation (7) shows that the rate limiting step in glucose oxidation in these conditions is gas phase reaction:

(s)

2NO + 02 --~ 2NO2 2. OXIDATIVE HALOGENATION OF AROMATIC

COMPOUNDS

As shown in [ 11,12] the NO2/NO catalytic system may be used in the iodination or bromination of aromatic compounds under the following conditions: Aromatics - HCIO4 NaNO2 - NaI(NaBr) - 02. At the same time attempts to chlorinate aromatics were unsuccessful. We have established that chlorinating of aromatics proceeds successfully if some metal salts are introduced into reaction mixture. The chlorinating reaction kinetics are investigated on the same equipment by the measurement of absorbed oxygen volume. The process in general consists of two stages. The first (very fast oxygen adsorbance) stage is connected with the reaction (8) in the gas phase. No chlorinated aromatics form during this stage. The second stage proceeds with lower rate, and in 60 - 70 rain the yield of the reaction product reaches 60-80%. 3. OXIDATION OF Fe(ID AND Ti(HI)

Oxidation of variable valence metal ions to higher valence state is desirable, for example, in hydrometallurgy, water conditioning, etc. The kinetics of Fe(II) and Ti0II) oxidation by oxygen in acidified water solutions was investigated. An investigation of oxidation process kinetics was carried out in the same equipment as mentioned above. It was established that

591 the initial oxidation rate of both Fe(II) and Ti(III)) Wo consists of catalytic Wo, cat. and noncatalytic W o, ncat. initial oxidation rates. Wo, cat. at constant [H + ] and p 02 is not dependent on temperature and depends on [NaNO2] o according to equation Wo, cat= 28• [NaNO2]o

2

(9)

The reaction mechanism includes reactions (II) - (IV), but Men+ oxidation to Me (n+l)+ proceeds according the following reaction Me n+ + NO2, aq --~ Me (n+l)+ + NO2-,aq

(X/I)

The analysis of the obtained kinetic data leads to the conclusion, that as in the majority of other cases when NO2/NO catalytic system was used, the rate limiting stage is gas phase reaction (8) , which according to literature data follows the third order kinetic equation: W = k • PNO 2 • P02

(1 O)

On the basis of the data obtained, effective technologies were created for the regeneration of Fe(III) containing solutions used in hydrometallurgy for ores leaching , in radioelectronics and in printed board production. It is evident, therefore, that NO2/NO is a cheap, simple, effective and very promising catalytic system. REFERENCES

I.D. Ross, C.-L. Gu, G.Hum, R. Malhotra. Int. J. Chem.Kinet.Vol.18 (1986) 1277. 2. R. Moode, S. Richards. J.Chem.Soc.Perkin Trans.II.No. 11 (1986) 1833. 3. A. Levina, S. Chomaja, S.Trusov, T. Stelmah. Kinetics and Catalysis.Vol.32 No. 6 (1991) 1336. 4.A.Levina, T.Stelmah, S.Trusov. Benzaldehyde synthesis method, USSR Patent No. 15112962 (1989). 5. A.Levina, S.Trusov. Kinetics and Catalysis. Vol.33 No.1 (1992) 92. 6. V.Nehoroshkov, G.Kamalov at al. Dibenzyl ether oxidation method, USSR Patent No. 1154261 (1985). 7. A.Levina, S.Trusov. Benzaldehyde synthesis method, USSR Patent No. 1657488 (1991). 8. A.Levina, S.Trusov. Kinetics and Catalysis.Vol. 32 No.6 (1991) 1343. 9. A.Levina, S.Trusov. Journal of Molecular Catalysis 88(1994) L 121. 10. T.Stelmah, I.Grigorjeva, S.Trusov. Glucaric acid synthesis method Latvian Patent No 10857 (1996). 11 .J.Dorfman at al. Kinetics and Catalysis. Vol.29 No. 1 (1988) 59. 12.J.Dorfman at al. Ibid.Vol. 30 No. 2 (1989) 303.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

593

Oxidative Coupling of Isobutene in a Two Step Process H. Hiltner and G. Emig Lehrstuhl for Technische Chemie I, Universit~t Erlangen-N~rnberg, Egerlandstr. 3, 91058 Erlangen, Germany, Tel: (+49-9131) 857420, Fax: (+49-9131) 857421, Email: [email protected]

ABSTRACT Oxidative coupling of isobutene suffers from severe deep oxidation. As in many other partial oxidation reactions selectivity remains low, despite intensive optimization of catalysts and reaction conditions. Among various new reactor concepts, the separation of catalyst reduction and reoxidation is very promising (two step process). Reaction engineering investigations of the two step process have been done. The influence of reaction conditions and reversibility of reduction/reoxidation cycles have been investigated. Based on the reaction engineering results a first approach to a kinetic model of both reaction steps has been developed.

1. INTRODUCTION A typical example for the oxidative dehydrodimerization of alkenes by heterogeneous catalysis is the conversion of isobutene to 2,5-dimethyl-l,5-hexadiene (DMH) catalyzed by metal oxides. The overall reaction scheme is: CH3 I 2 CH2=C~CH3

+ 1/2 02

Cat. ~

CH3 CH3 I I CH2=C~CH2-CH2-C=CH2

+ H20

DMH is an intermediate in peroxide chemistry and could be used for synthesis in the field of lubricant or polyester. So far research has mainly focused on the development of different catalysts and the comparison of their performance [1]. Despite intensive optimization of catalysts and reaction conditions, selectivity to DMH is still low, because molecular oxygen in the gas-phase causes deep oxidation. Beside the development of catalysts and the optimization of reaction conditions, the mode of gas-solid contact and the reactor configurations are important issues as well. As in any parallel-series network with valuable intermediate products, the design of the reactor has a strong influence on the selectivity towards DMH. In principle, modes of contact which maintain a low oxygen concentration in the reactor favour the desired reaction and thus improve selectivity. Obviously, the requirement of a low oxygen concentration in the gas phase can be met by using a redox-type operation in which a

594 reducible metal oxide catalyst is used as a source of lattice oxygen. This requires organization of the process into reduction-reoxidation cycles. The separation of catalyst reduction and reoxidation is based on the assumption that partial oxidation on metal oxides proceeds via a cyclic reduction-oxidation-mechanism [2]. In a first half-cycle the hydrocarbon reacts with lattice oxygen of the catalyst to the product while the oxide is reduced (reduction step). The catalyst acts as an inorganic carrier of oxygen. In the second half-cycle molecular oxygen reoxidizes the reduced catalyst (reoxidation step). The main advantage of the two step process is a higher selectivity as compared to the conventional process. Burch and Swamakar [3] established this for the oxydehydrogenation of ethane just as Mtiller-Erlwein and Guba [4] for the oxidation of isobutyraldehyde to methacrolein. Emig et al. [5] applied the two step approach to the butane-oxidation to maleic anhydride. DuPont intends to install the two step process for butane oxidation to maleic anhydride on a technical scale [6-9]. Weismantel et al. [10] obtained an improvement of selectivity with a two step process for the oxidation of isobutyric acid.

2. EXPERIMENTAL Vapour-phase catalytic oxidation of isobutene was carried out at atmospheric pressure in a completely automated laboratory setup, including a fixed bed reactor (700 mm length, 10 mm inner diameter) with corundum as wall material. In order to ensure isothermicity, the heated section (200 mm in length) was divided into five independently heated zones and the catalyst bed was diluted with inert pellets (cz-A1203). Inert pellets were placed above and below the catalyst bed to ensure a well-mixed feed stream, and to preheat the gas to the reaction temperature. Bi203 catalyst (Merck) was pressed into thin wafers and broken into small particles. Granules with a diameter of 0.8 - 1.2 mm were used. The gas flows (isobutene, oxygen, helium, methane) were controlled by thermal gas flow controllers. Helium was used as a balance. The exit gases were analysed by on-line gas chromatography with FID and TCD. Methane was used as internal standard to calculate component concentrations from peak areas. The whole setup including the gas chromatograph was controlled by a personal computer using a software which carried out predesigned experimental plans. All mass-flows and temperatures as well as the analytic data were recorded by this software. The exhaust gases were catalytically combusted in a subsequent reactor filled with a novel platinum/palladium coated knitted fiber catalyst packing [ 11 ]. Catalyst reduction/reoxidation experiments were usually carried out between 500~ and 600~ The inlet partial pressure of isobutene or oxygen ranged from 2 to 60 kPa, respectively. The catalyst bed was purged with helium between subsequent half-cycles to remove adsorbed species. Prior to the catalyst reduction experiments, the catalyst was completely reoxidized. Before a reoxidation experiment was started, the catalyst was reduced for 90 min at 550~ with an isobutene inlet partial pressure of 2 kPa leading to a reproducible degree of reduction.

595 3. RESULTS AND DISCUSSION 3.1.

Investigation

of the

reduction

step

Figure 1 shows the conversion of isobutene and the selectivity to DMH during the reduction step (index: 2-step). Additionally, the corresponding curves for the conventional oxidation in the presence of molecular oxygen are presented (index: conv). In the case of the conventional process, conversion and selectivity show no temporal dependencies. Since reduction and reoxidation occur simultaneously, the catalyst has a steady oxidation state. In the two step process, conversion decreases with time of reduction due to the consumption of lattice oxygen. An important aspect is the significantly higher selectivity to DMH in the two step process as compared to the convemional oxidation. The improvement in selectivity originates from a lower amount of deep oxidation due to the absence of molecular oxygen during the catalyst reduction step. However, the formation of carbon dioxide was observed, decreasing with time of reduction. Therefore it has to be considered that lattice oxygen reacts with isobutene to carbon dioxide. The increasing selectivity to DMH in the two step process suggests that the formation of DMH is less sensitive to the loss of lattice oxygen than the formation of carbon dioxide. Following Misono [12-14], who has investigated the catalytic properties of heteropolymolybdates, the abstracted protons and electrons in alkene activation are able to move easily into the catalyst bulk and react there with lattice oxygen to water. Due to the fast diffusion of electrons and protons in the catalyst bulk, lattice oxygen in the whole catalyst participates in the formation of water. Thus, the whole catalyst bulk takes part in dimerization, 35

100

30 -51~ 25

-B

.

. . . . .

~

0

0

0

0

0

,'~

- 90 - 80

-I~

. . . . .

9 . . . . .

-U- . . . . .

~

. . . . .

9 . . . . . .

II- . . . . .

9

- 70

-60

= 20 o ra~ t-q

-

> 15 o

50 "~

-40 '~

v

v

~

v

v

10 ""'13-

. . . . .

D._ . . . . .

Q . . . . . .

-

30

-

20

,-,,,,i

r~

1:3- . . . . .

-10 I

I

I

I

20

40

60

80

time of reaction [rain] Figure 1. Conversion of isobutene (dashed lines) and selectivity to DMH (full lines) in the two step process and in the conventional process. T = 550~ W/F = 5000 g-min-mol-x, Pisobutene= 20 kPa, P o x y g e n = 20 kPa. 9 conversionconv; 9 selectivityconv; n conversion2_ste0; O selectivity2.step

596

0.7 ,-, 0.6 .,..~

_.~ o.s & 0.4 9 ,-.,4

o

0.3

,y o 0.2 0.1 0.0

I

0

20

I

I

40 60 time of reduction [mini

I

80

Figure 2. Influence of the partial pressure of isobutene on the space-time yield of DMH during the reduction step. T = 550~ W/F = 2500 g.min.mo1-1, Pisobutene= 9 10 kPa, O 20 kPa, 9 40 kPa, VI 60 kPa although reaction occurs only at the surface (bulk-type reaction). The formation of carbon dioxide, in contrast is directly connected with the consumption of lattice oxygen at the catalyst surface. The reaction proceeds only at the catalyst surface and only oxidized sites at the catalyst surface are involved in the formation of carbon dioxide (surface-type reaction). Classification of dimerization and deep oxidation in bulk-type and surface-type reactions explains the behavior of selectivity well. However the different characteristics between oxidation and oxydehydrogenation may also be explained by assuming two different types of oxygen atoms in the catalyst, i. e. strongly and weakly bonded oxygen [ 1]. This assumption is in principle identical with the classification in surface-type and bulk-type reactions. An increase of the partial pressure of isobutene yields an enhancement in selectivity. Therefore high isobutene concentrations promote the dimerization reaction to DMH at the catalyst surface. This fact can be explained by a higher degree of adsorption by isobutene molecules. With increasing isobutene partial pressure more active sites are covered by isobutene, so there is a greater probability that two isobutene molecules combine to DMH. A higher hydrocarbon concentration also results in a higher space-time yield (Figure 2). So as a whole high isobutene inlet partial pressures improve the performance of the reaction. 3.2. I n v e s t i g a t i o n of the r e o x i d a t i o n step

The reoxidation step serves to restore the initial oxidation state of the catalyst. Normally air would be an economic reoxidation means, but here considerably lower oxygen partial pressures (1 - 4 kPa) are used for keeping an efficient time resolution with the existing analysis technique.

597

100 -

8O

~

6O

=

40

o

20 -

0

t

t

0

20

40

60 80 100 time ofreoxidation [rain]

120

140

160

Figure 3. Influence of the partial pressure of oxygen on the conversion of oxygen during the reoxidation step. T = 550~ W/F = 5000 g-min.mol -~, Poxygen= 9 1 kPa, O 2 kPa, 9 4 kPa In Figure 3, conversion of oxygen during the reoxidation is shown for different partial pressures of oxygen. Conversion starts at 100%, remains there for some time and then decreases very strongly. The reoxidation lowers the degree of reduction leading to decreasing rates of reoxidation and thus decreasing oxygen conversion. At higher oxygen partial pressures the period of total oxygen conversion is shorter, indicating that with increasing partial pressures of oxygen, catalyst reoxidation proceeds faster and the initial oxidation level is approached earlier. 3.3. Reversibility of the redox cycles

Keeping the activity of the catalyst for a long period of time is a prerequisite for an economic operation of the process, since frequent changes of the catalyst activity result in a decrease of efficiency. In the two step process the catalyst is stressed by repeated change between reducing and oxidizing environments. It has to be studied how far this stress affects a reversible function of the catalyst and the reversibility of the redox cycles at all. Figure 4 shows the conversion and the selectivity over 28 successive redox cycles. Each data point was measured after a reduction time of seven minutes. During these redox cycles no significant decrease in conversion or selectivity occurred. 3.4. Kinetic m o d e l i n g

While the kinetics of the oxidative coupling of isobutene have already been studied for the conventional process [1], the two step process has not been examined yet. Following Patience and Mills [8] the simulation of the two step process by applying a kinetic model which is based upon experiments where the catalyst oxidation state essentially remains constant might

598

100 DMH selectivity 80--

9-

60--

o

r~

40--

o r.~

r

0

o

20-

isobutene conversion 9

m

m

m

m

m

m

m

m

m

m

m

m

m

m

m

m

n

--

n

--

_ m

mm m

m

m

m

m

l

I

i

I

i

5

10

15 redox cycles [-]

20

25

0 0

m

n

m

30

Figure 4. Behaviour of conversion of isobutene and selectivity to DMH during several redoxcycles. T = 550~ W/F = 800 g.min.mol l , Pisobutene= 12 kPa, 9 conversion of isobutene, O selectivity to DMH be subject to error, because the kinetic parameters have not been determined under typical unsteady conditions. To explain the effect of catalyst oxidation state on activity and selectivity a separate kinetic modeling of both catalyst reduction and reoxidation step is necessary. Kinetic modeling of both catalyst reduction and reoxidation step as well as coupling between gas phase and catalyst phase yields a set of differential equations: Gas phase: ax i .

.

at

R-T.Os .

.

. a x i + Ri] n tot " a m cat

. P tot

Catalyst phase: Reduction a 00x at

m cat N Ox,0

- - - - - . R

Reoxidation 0

ox

;

a 0 Red at

m cat -N- Ox,0 " R~ Red

(where i denotes isobutene, DMH, carbon dioxide and oxygen, respectively; R i and R0 denotes the reaction rate per catalyst weight; xi is the mole fraction; 0Ox is the fraction of oxidized sites; 0Red is the fraction of reduced sites: 0Red = 1 - 0Ox; Nox,0/mcat is the molar amount of available lattice oxygen atoms per catalyst weight; l:ltot is the total molar flow; Ptot is the total pressure; Ps is the density of catalyst in the reactor; R is the gas constant)

The consumption of isobutene and oxygen as well as the formation of DMH and carbon dioxide can be described by power law rate equations"

599 Reoxidation:

Reduction: mi

mt

Os 9X o 2

RDMH -" kDMH " XC4H 8 " 0 Ox

ROE = - k o 2

ni nt R c o 2 - k c o 2 9XCaH8 90Ox

R%e d = 2.R02

Ot 90Re

d

1 RC4Hs = - ~ - " R c o 2 -

2.

R DMH

Roox = -3. Rco: - RDMH The temperature dependency was taken into consideration by expressing each of the rate constants as: (where T* is an arbitrary reference temperature, * I_~(~_r, 1)] here: T* = 500 r e K;a Tc den~ t i ~ the k i - k i .exp " temperature; k i is the rate constant at the reference temperature; EA, i is the activation energy) Unlike the standard form of the Arrhenius equation, this form does not exhibit the strong and undesirable correlation between the pre-exponential factor and the activation energy. All values of kinetic parameters and the ratio Nox,0/mcat were estimated with SimuSolv T M [ 15] by nonlinear regression analysis based on numerous experiments. The result of parameter estimation for both steps is given in Table 1. Figure 5 shows some of the calculated and measured values of mole fraction of isobutene, oxygen, DMH and carbon dioxide as a function of time for different initial partial pressures. Table 1 Results of parameter estimation for the catalyst reduction and reoxidation step Parameter

Value

Standard deviation

kDM H

[mo1/(gcat'S)]

1.1.10 -9

0.11" 10-9

kco 2

[mol/(gcat'S)]

1.2.10 "12

0.06.10 -12

k*0 2

[mol/(gcat'S)]

1.1.10 -2

0.21.10 -2

EA,DMH EA,CO2

[kJ/mol] [kJ/mol]

57 130

1.2 1.2

EA,O2

[kJ/mol]

18

3.9

mi mt ni nt Os

[-] [-] [-] [-] [-]

Ot

[-]

0.9 0.7 0.6 1.6 1.5 1.9 2.1.10 -4

0.01 0.04 0.02 0.07 0.04 0.07 0.01.10-4

Nox,0/mcat [mol/gcat]

600

Z

0.08

0.08

0.06

0.06

0.04

._o 0.04

0.02

_~ 0.02

o

'-'

_

_

"-"

"-"

~

_

=

z

"-" "-"

"

e-

o o

O

O

E 0.00

E

2500

0.00 ,.

5000

0

_

1000

time of reduction [s]

{

2000

time ofreoxidation [s]

0.008

,._, 0.004

Z 9 0.006 .2 0.004

~

0.002

o t~

'~ 0.002 0

E 0.000

O I

I

2500

5000

time of reduction [s]

E

0.000 0

2500

5000

time o f reduction [s]

Figure 5. Calculated and measured profiles of mole fraction of isobutene, oxygen, carbon dioxide and DMH as a function of time. T = 550~ W/F = 2000 g'min'mol l , Pisobutene,oxygen= O 2 kPa, A 4 kPa, I'-1 8 k P a , - calculated The comparison of calculated and measured values displays a satisfactory agreement. The values of the parameter estimation shown in Table 1 give a good description of the experimental data with the underlying kinetic model. Figure 6 displays the parity plots of experimental data with the values calculated from the model. The points scatter randomly around the bisector of an angle. With few exceptions the points lie inside the borders of 20 % deviation. The disagreement between calculated and measured values of oxygen mole fraction results mainly from the marked gradient of oxygen mole fraction in the course of reoxidation (Figure 5). Therefore even little shifts on time axis cause large errors. The distinct deviation between the calculated and measured mole fractions in the field of very low values results from an error in the gas analysis by gas chromatography at very low oxygen concentrations. The content of lattice oxygen of the catalyst during reoxidation is not accessible to measurements. Figure 7 shows the calculated fraction of oxidized sites (normalized with the fraction of reduced sites at the beginning of the reoxidation cycle 0Red,0) with dimensionless reactor length for different times of reoxidation. The fraction of oxidized sites 0Ox increases with time of reoxidation due to the uptake of oxygen. At the entrance of the reactor the degree of oxidation increases quickly due to the high oxygen partial pressure. At the exit of the reactor reoxidation is slow because less oxygen is available in this part of the reactor. But with increasing time of reoxidation more and more oxygen reaches the exit, so reoxidation can also proceed there intensively. Altogether there is a delay in the increase of 0Ox at the exit of the reactor. This leads to a pronounced profile of the content of lattice oxygen, which is determined by reactor length and time of reoxidation.

601

0.10

0.08 "-" "~ ~o

1 C4H8

E ~9

] 0.04

o

-~

J ~ " "

E

+2

+2

.=o 0.05 ~

- 20%

L o

0

E

02

0

0.00 0.00

I 0.04 mole fraction experiment

E

0.08

0.00

.,.

I

0.00

0.05

0.10

mole fraction experiment

0.010

0.004 C8H14

+ 20% ~9 0.005

.-~ 0.002 f

0.000

,

0.000

201

-

%

~ 0.000

0.005

0.010

0.000

0.002

0.004

mole fraction experiment

mole fraction experiment

Figure 6. Parity plots of calculated and observed mole fractions of isobutene, oxygen, carbon , 9 -1 dioxide and DMH. T = 530~ 550~ 570~ W/F = 1000 g'min'mol ~ 2000 g.mln.mol , Pisobutene = 2 kPa, 4 kPa, 8 kPa

ilill !~iiii? 84

,ii liliputii,i~i~!i~'~(>

9

o

_L

l

0.5

L! 700

time [s]

0

-~0.25

0.5

0.75

dimensionless reactor length [-]

Figure 7. Calculated fraction of oxidized sites with dimensionless reactor length for different times of reoxidation. T = 550~ W/F = 2000 g-min-mo1-1, Poxygen= 2 kPa

602 4. CONCLUSIONS The oxidative coupling of isobutene can be performed in two separate steps, connected with reduction of catalyst and reoxidation of the reduced catalyst afterwards. The two step process leads to an improvement of DMH selectivity as compared to the conventional process. The formation of carbon dioxide requires surface lattice oxygen from the catalyst, while formation of DMH occurs by abstraction of protons and electrons at the catalyst surface. They are absorbed on the catalyst bulk and, finally, react to water there. Thus, the rate of carbon dioxide formation is more affected by catalyst reduction than the rate of DMH formation. The existence of a kinetic model is a prerequisite for scale-up and simulation calculations. Kinetic modeling of the unsteady process is possible by coupling gas phase and catalyst phase. Calculated and experimental data show good agreement. The operation with a two step process is a possible alternative for oxidation reactions with oxide catalysts. Its application is imaginable for all processes with a high degree of deep oxidation.

ACKNOWLEDGEMENTS

Financial support by Bayerischer Forschungsverbund Katalyse (FORKAT) and by AKZO Nobel Faser AG is gratefully acknowledged.

REFERENCES

[ 1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [ 11] [12] [ 13] [ 14] [ 15]

E.A. Mamedov and V.D. Sokolovskii, Catal. Today, 14(1992)343. P. Mars and D.W. van Krevelen, Spec. Suppl. Chem. Eng. Sci., 3(1954)41. R. Burch and R. Swamakar, Appl. Catal., 70(1991)129. E. Mtiller-Erlwein and J. Guba, Chem. Ing. Tech., 60(1988) 1072. G. Emig, K. Uihlein and C.-J. H~icker, Stud. Surf. Sci. Catal., 82(1994)243. R.M. Contractor, H.E. Bergna, H.S. Horowitz, C.M. Blackstone, B. Malone, C.C. Torardi, B. Griffiths, U. Chowdhry and A.W. Sleight, Catal. Today, 1(1987)49. R.M. Contractor and A.W. Sleight, Catal. Today, 3(1988)175. G.S. Patience and P.L. Mills, Stud. Surf. Sci. Catal., 82(1994)1. R.M. Contractor, D.I. Garnett, H.S. Horowitz, H.E. Bergna, G.S. Patience, J.T. Schwartz and G.M. Sisler, Stud. Surf. Sci. Catal., 82(1994)233. L. Weismantel, J. St6ckel and G. Emig, Appl. Catal. A, 137(1996)129. G. Emig, B. Gmehling, N. Popovska, K. H61emann, A. Mayer and A. Buck, Proc. SAE Conf., Detroit, SAE 960138(1996)123. M. Misono, Catal. Rev. Sci. Eng., 29(1987)269. M. Misono, N. Mizuno and T. Komaya, Proc. 8th Int. Congr. Catal., V5(1984)487. M. Misono, N. Mizuno, H. Mori, K.Y. Lee, J. Jiao and T. Okuhara Stud. Surf. Sci. Catal., 67(1991)87. E.C. Steiner, T.D. Rey, P.S. McCroskey, The Dow Chemical Company, 1990.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

603

Solid Solutions for Cleaning up Chemical Processes using Hydrogen Peroxide. Sharon L. Wilson and Craig W. Jones Solvay Interox R&D Widnes Laboratory, P.O. Box 51, Moorfield Road, Widnes, Cheshire, WA8 0FE, U.K. This review examines the work carried out .either in-house or via external research contracts by Solvay Interox into the area of heterogeneously activated aqueous hydrogen peroxide chemistry. The review is focused on the preparation and use of solid catalysts in the presence of H202 for the manufacture of fine chemical and pharmaceutical intermediates.

1. INTRODUCTION

In bulk chemical manufacture the choice of primary oxidant is largely restricted to molecular oxygen for economic reasons. However, liquid phase oxidations with dioxygen are generally radical chain processes [1] and the intermediate alkylperoxy and alkoxy radicals are largely indiscriminate in their reactivity. Selective oxidations are generally speaking, observed only with relatively small molecules containing one reactive group. Thus, although dioxygen is a relatively cheap oxidant with no waste problem its scope is limited to a small number of simple petrochemicals . . . . . The economics of fine chemical or pharmaceutical production, in contrast, allows for a broader range of primary oxidants to be employed. Indeed, even though H202 is more expensive per kilo than 02, it can be the oxidant of choice for such industries, because of its simplicity of operation i.e. for small scale operations the total cost of equipment and raw materials may be lower for oxidation employing H202 than 02 [2]. Several industrial processes using H202 have been developed. For example, the oxidation of HCN to cyanogen [3], the production of hydroquinone and catechol by the hydroxylation of phenol [4], the oxidation of ammonia to hydrazine [5], and the epoxidation of natural oils [6]. However, the majority of synthetic routes developed to date rely upon the homogeneous activation of H202 [7]. The use of homogeneous catalysts usually has the advantage of a relatively high rate of reaction as compared with other catalytic forms. However, immobilization of metal complexes onto either polymer backbones [8] or high surface area inorganic oxides [9] offers several advantages over their homogeneous counterparts. For example, catalyst recycle and isolation is easier, often requiring simple filtration or continuous reactor technologies. The support may or may not be an inert substrate. It can play a positive role leading to preferred orientations of the substrate at the catalytic site so promoting product selectivity [10]. Organic functional groups co-valently bound to the surface of crystalline solids or polymers are subject to special constraints which can alter their chemical reactivity relative to the unsupported metal complex [11]. Other benefits can also be altering

604

equilibrium positions [12], and stability of catalytically active but unstable structures [ 13]. Supported metal complexes for the activation of H202 to afford selective oxidation of organic compounds to fine chemical and pharmaceutical intermediates is the purpose of this review. The work presented here is based on studies carried out either in-house or via externally sponsored research [14]. The work discussed below will be divided into the following areas: 1. The preparation and use of supported polyoxometal complexes with H202 for the oxidation of sulphur containing compounds to sulphoxides, sulphonic acids, and epoxidation of alkenes to epoxides. 2. The preparation of chromium silicalite via a variety of routes for the oxidation of styrene in the presence of H202. 3. The preparation and use of metal(IV) phosphates with H202 for the oxidation of ketones to esters.

2. EXPERIMENTAL 2.1 The Preparation of Heterogeneous Catalysts for the Activation of Aqueous Hydrogen Peroxide The following catalyst preparations are described; phosphotungstic acid on ~/alumina, ammonium molybdate on a cross-linked polystyrene anionic exchange resin, metal(IV) phosphates, and chromium silicalite. General Procedure for the Preparation of Phosphotungstic Acid on Alumina Phosphotungstic acid (3 g) was dissolved in demineralized water (25 ml). To this solution was added y-alumina (20 g), and the mixture stirred at ambient for 4 h. The slurry, was then filtered at the pump, and the residue dried in an oven at 60 ~ for 18 h. The air dried solid was then calcined at 500 ~ for 4 h. in a muffle furnace. The catalyst produced had a nominal polyoxometalate loading of 10 % w/w. General Procedure for the Preparation of Polymer Supported Ammonium Molybdate (NH4)6Mo~O24was supported on a cross-linked polystyrene strongly basic anion exchange resin in chloride form by stirring 2.3 ml of Amberlyst A26 beads in a solution comprising (NH4)6MoTO24 (2.1 g) dissolved in 20 ml of demineralized water at room temperature for 2 h. The resin was filtered off, washed with demineralized water (2x50 ml) followed by vacuum drying for 2 h. Preparation of Metal(IV) Phosphates There are several forms of metal(IV) phosphates, including amorphous, crystalline, and pellicular. However, here only the preparation of amorphous zirconium phosphate is described. Phosphoric acid (10 % w/w. 162 ml) was added to zirconyl chloride octahydrate (14.4 g in 112 ml of demineralized water), and stirred at room temperature for 4 h. The resulting gel was washed repeatedly with demineralized water (3 x 500 ml) and centrifuged each time.

605

After the final wash, the gel was converted to its H* form by immersion in 1 M nitric acid (250 ml) for 24 h. The suspension was then filtered off and washed with DMW, until a pH of 5 was obtained. The residue was then dried at 100 ~ for 18 h. under vacuum. The yield of zirconium phosphate was 11.0 g ~)m,x 3431 (br. OH), 2361 (sh. OH), 1631 (sh. P-O), 1050 (br. P-O)cm 1.

Preparation of Chromium Silicalite

Chromium silicalite has been prepared via three different routes, NH3, F, H2SO4 (see text). However, here we describe the use of the fluoride anion route. 18.1 g of Aerosil 200, 2.5 g of ammonium fluoride, 10 g of tetrapropylammonium hydroxide were added to DMW (178 ml), and the mixture stirred for 18 h. (solution A). Chromium(Ill) chloride (2.2 g in 14 ml of DMW) was added to solution A, and aged in an homogenizer for 30 minutes. The mixture was then placed in a 300 ml Teflon bottle, and placed in a stainless steel autoclave, and heated at 169 ~ under autogeneous pressure for 25 days. After the reaction was quenched, the green solid was filtered off, and the mother liquor diluted with hot water to facilitate isolation of the material. The product was dried at 70 ~ for 12 h. under vacuum, followed by calcination at 550 ~ for 24 h. The xrd pattern was in agreement with a ZSM-5 type zeolite structure. 2.2. The Oxidation of Organic Compounds using H202 in the Presence of

Heterogeneous Catalysts

The following oxidations will be described; the oxidation of penicillin-G potassium salt to the sulphoxide; the oxidation of cyclooctene; the oxidation of cyclohexanone; and the oxidation of styrene.

The Oxidation of Penicillin-G Potassium Salt

Penicilin-G (4 g, 10.8 mmol), water (40 ml), and the catalyst ((NH4)6MoTO24-A26) (0.5 g) were charged to a three necked flask, fitted with, thermometer, cooling bath, overhead stirrer, and peristaltic pump, was equilibrated to 20 ~ H202 (35 % w/w. 1.04 g, 10.8 mmol) was added via the pump over 45 mins. On completion of the H202 addition, the reaction was maintained under these conditions for a further 45 mins. Analysis of the reaction was via H.P.L.C. and comparison to the authentic compound. Benzenethiol was also oxidized using phosphotungstic acid on alumina. However, the solvent was t-butanol, and the reaction temperature was 80 ~ for 4 h.

The Oxidation of Cyclooctene

Cyclooctene (0.097 mol), H3PW12040-AI203 (0.2 g), and t-butanol (40 ml) were charged, as described above. The mixture was warmed to 70 ~ and H202 (30 % w/w. 9.47 g) added over 1 h. The temperature was maintained for a further 5 h. Analysis was via G.C. and comparison with an authentic sample of cyclooctene oxide.

606

The Oxidation of Cyclohexanone H202 (70 % w/w. 51.4 mmol) was added over 1 h. to a suspension of metal(IV) phosphate (0.16 g), cyclohexanone (0.3 g, 3.02 retool), and glacial acetic acid (4 ml) and reacted for 15 h. at room temperature. Analysis was via G.C. and comparison to authentic ~-caprolactone. The Oxidation of Styrene H202 (35 % w/w. 50 retool) was added over 1 h. to a suspension of styrene (100 retool), CrS-1 (0.1 g), and 1,2-dichloroethane (40 ml) at 70 ~ After the addition of the peroxide, the reaction was carried on for a further 3 h. The oxidation products were analyzed for by G.C. and against authentic compounds. 3. RESULTS AND DISCUSSION 3.1. The Activation of H=O2 in the Presence of Supported Polyoxometalates The early transition metal isopoly or heteropolyoxometalates are a substantial family of anionic inorganic cluster-like compounds [15]. The most common, accessible, and investigated class of polyoxometalates are relatively robust to thermal decomposition. A considerable amount of homogeneous chemistry has been carried out in the presence of aqueous H=O= [16-17]. However, we have been interested for sometime in the oxygen transfer ability of immobilized polyoxometalates in the presence of H202. The two complexes we have paid most attention too were, ammonium molybdate, and phosphotungstic acid supported onto a cross-linked polystyrene anionic resin, and y-alumina respectively [18-19]. The oxidations of interest were; Penicilin-G to the sulphoxide (an important intermediate for the preparation of cephalosporin antibiotics [20]), and the preparation of sulphonic acids from thiols [21], and the epoxidation of alkenes. 0

0

SH

SO3H

0

607

3.1.1. The Oxidation of Sulphur Containing Compounds The oxidation of Penicilin-G potassium salt by H202 was carried out in the presence of a range of immobilized polyoxometalates (Table 1). All reactions were run at room temperature or sub-ambient conditions in water or water/lower alcohol mixtures. The (NH4)6Mo7024-A~ system was found to be the most efficacious catalyst screened. It is interesting to note that the un-supported salt was less selective towards the sulphoxide compound compared with the immobilized analogue. The work suggests that the catalyst has been "tamed". The molybdenum containing heteropolyacids were found to be less effective compared to their tungsten counterparts. This type of activity difference was believed to be due to the greater asymmetry observed for 112 peroxotungstate bonds. The titanium silicalite catalyst (TS-1, ZSM-5 structure) was found to be particularly poor for such a transformation. The reason was probably due to the relatively small pore size (5.5 A) thus denying access of the substrate to the active site. Table 1. The Oxidation of Penicilin-G to the Sulphoxide using H202 in the Presence of Supported Polyoxometalates ~ Catalysi

'

'

Solvent

'Temperature 'Yield C0nversion'Selectivity ~ % % % None 'H20 20 9.6 1::;; ' ' 80 H4PMo1104ob H20 20 34 69 49.3 H3PW1204oc H20 20 72.5 96 75.3 H4PMo1104ob ~PrOH/H20 -10 21.8 31 70.3 H3PW1204oc ~PrOH/H20 -10 48.3 74.4 65 (NH4)6Mo7024d H20 20 75.0 96 78 TS-1 e H20 20 5.0 10 50 (NH4},6MozO24f H20 20 69 100 69 aSee experimental, t)Supported'on alumina, CSupportedon alumina and reaction mixture left refrigerated for 5 days at -2 ~ pdor to analysis, dSupportedon A26 beads, eZSM-5redox zeolite catalyst, fHomogeneousunsupportedsalt. As an extension to our studies, we have looked at the preparation of phenylsulphonic acid from benzenethiol using immobilized heteropolyacids in the presence of H202. The transformation was of interest to us because homogeneous metal catalyzed systems generally give poor yields of suiphonic acids due to preferential formation of disulphide compounds. However, DeShriver and co-workers [22] have achieved excellent yields of isothionic acid from 2-mercaptoethanol using tungstate catalysts in the presence of H202. The problem with the system was that the substrate was added to a hot solution of the H202/catalyst. Such an addition regime is unlikely to be viable on a large scale due to inherent safety problems (i.e. H202 decomposition before oxygen atom transfer). Hence, we have attempted to oxidize the thiol to the acid via the safer addition route (i.e. H202 to substrate/catalyst). The results are presented in Table 2.

608

Table 2. Oxidation of Benzenethiol with H202 in the Presence of Heterogeneous Catalysts" Catalyst None H3PW1204ob H4SiW1204ob aReaction run at 80 ~

,Yield , '% ,Conversion % 42 100 96 100 69 100 in t-butanol over 5 h, bSupported on alumina.

Selectivity,% 42 96 69

,,,

The results show that whilst all three systems convert the substrate fully, the immobilized H3PW1204o on alumina affords the desired product in high selectivity. The difference in selectivities between the two heteropolyacids may be due to their relative stabilities under the conditions employed [23]. In conclusion, we have studied the supported polyoxometalate/H202/solvent system for the oxidation of sulphur containing compounds, and found that significant advantages in terms of product selectivity can be obtained when homogeneous species are immobilized.

3.1.2. The Oxidation of Alkenes using Immobilized Phosphotungstic Acid Epoxidation processes currently account for around 30 % of H202 consumption in the chemical industry. Present processes are mainly peracid based and have limitations in scope and safety, as well as incurring high capital investment. Consequently, we have attempted to develop heterogeneous catalyst systems for such transformations. A system which has been extensively studied has been the epoxidation of cyclic alkenes with H202 in the presence of supported phosphotungstic acid. Table 3 illustrates our results. Table 3. The Epoxidation of Alkenes with H=O2 in the Presence of H3PWtaO40-yAI203" Substrate Product ' Conversion % Select, ivity % Cyclooctene Cyclooctene oxide 98 98 Cyclohexene Cyclohexene oxide 29 63 Cycloheptene Cycloheptene oxide 54 90 Cyclopentene Cyclopentene oxide 30 87 1-Octene 1-Octene oxide 5 100 aReactions run in t-butanol at 80 ~ over 5 h. The H3PW1204o-AI203 was calcined at 500 ~ for 4 h. prior to use. As can be seen from the results the order of reactivity was cyclic alkenes greater than linear alkenes. The order is the expected one, and illustrates that the substrate had no difficulty in reaching the catalytically active sites on the solid surface. The immobilized phosphotungstic acid material complements a titanium silicalite type (TS-1) catalyst, which is much more effective towards the epoxidation of linear alkenes in the presence of aqueous H202, and particularly

609

poor when screened against cyclic alkenes. We have also found that the temperature of calcination has an important role to play (Figure 1). Figure 1. The Effect of Calcination Temperature of Phosphotungstic Acid on Alumina on the Yield of Cyclooctene Oxide 90-,.........

........................................................................................................

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80 i~i~i~i~i~i~i~i~i~i~i~i~i~i~i~!~!i;~i~i!~!~i~!~!i~i~i~i!~!i~;~i!~i~!i~i!ii;i~i!~!i!i~i!ii!i!i~i!;i!!;!~i~i!i~i:~i~i~i!i!i~i~i;? ~~

....

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w

....

5o .;i~i~i~;~i;i~i~i~i~i;~;~;~i~i~i~i~;~ii}iii;~;~;!i!i~ii!~i~i!ii~;~i~!~i~i~iii~i~i;ii;~iiii;i.... ;~;ii~i~i~ii;ii~i!~i}~ii~i!;i~!~i!i~:~

~' i ~.ii~iii:i!i!i~i~iii~iiiiii!i~!!ii:i:~iiiiiiii~iiii!!;i~ii!iiiiiiiiii;iii!iiiiiiiiiiiii~i:i::i~iiiii~iii~iiiii~i!i~i~ii:i~i:iii~iiiii~ii!~:~ .... ~ ~~

....

~~

......

0

100

200

300

400

500

600

Calcmatmn TempJ C

For example, at a calcination temperature of 100 ~ the yield of epoxide was found to be only 10.3 %, whereas at 500 ~ the yield dramatically increased to over 88 %. However, at a calcincination temperature of 600 ~ the yield drops off suddenly. Work is still on-going as to the nature of the active species. However, solid state NMR, IR, and Raman all show that changes take place upon calcination which is not just the formation of layered tungstates. For example, a dimeric species appears to be stabilized by the alumina support when the phosphotungstic acid salt was immobilized. 3.2. The Oxidation of Ketones with H202 to Esters in the Presence of Metal(IV) Phosphates The Baeyer-Villiger reaction (1) is of considerable synthetic use for the shortening of carbon chains, hydroxylating aromatic rings, and converting carbocycles to heterocycles and opening up cyclic arrays to prepare functionalized chains or rings [24] o

o

(I) Solvent, H202

Little work has been reported on the use of metal(IV) phosphates as oxidation catalysts. Metal phosphates, are potentially interesting catalytic species, since certain forms posses regular layer structures which have strongly acidic sites

610

which may undergo ion exchange with a range of metal ions. The reaction of ketones with H202 using metal(IV) phosphates will be discussed. Table 4 summarizes our results.

Table 4. The Oxidation of Ketones with 35 % H=O= in the Presence of Metal(IV) Phosphates

Teml~.~

Products convl % Yield % 6-HCA 15 5 s-Cap. 5 TiPA 75 MeCN CX Polymer 90 90 ZrPA 20 AcOH CX Polymer 100 100 SnPC 48 AcOH CX 6-HCA 63 1 8-Cap. 7 Polymer 55 TiPA 47 MeCN CP DVL 72 20 TiPA 47 AcOH CP DVL 61 32 AcOH PIN t-BA 42 42 ZrPA 53 AcOH AP PhA 50 45 ZrWP 50 PhOH 5 ZrPA = Zirconium phosph'ateamorphous, TiPA = Titanium phosphate amorphous,'SnPC = Tin phosphate crystalline, ZrWP = Zirconium/tungsto. phosphate, MeCN = Acetonitrile, AcOH = Acetic acid, CX = Cyclohexanone, CP = Cyclopentanone, PIN = Pinacolone, AP = Acetophenone, 6-HCA = 6-Hydroxycaproic acid, s-Cap. = ~-Caprolactone, Polymer = Polycaprolactone of unknown molecular weight, DVL = ~-Valemlactone, t-BA = t-Butylacetate, PhA = Phenyl acetate, PhOH = Phenol, catalyst ZrPA

20

Solvent MeCN

,

Ketone CX

It can be concluded that TiPA, and TiPC have acidic properties which probably activate H202 to afford a good yield of polycaprolactone when acetonitrile was used as the solvent. Therefore, metal(IV) phosphates appear to be excellent catalysts for the preparation of esters under relatively mild conditions. The selectivity towards t-butylacetate was exceptionally good. The work suggests that possibly two mechanisms are operating, depending upon the solvent and metal(IV) phosphate employed. For example, in acetonitrile the oxidizing species may be activated H202, whereas, in acetic acid the production of peracetic acid within the inter-lamellae spacing was possibly the oxidizing source.

3.3. The Oxidation of Styrene with H202 in the Presence of Chromium(Ill) Silicalite (CrS-1)

There are two major problems associated with selective oxidation catalysis with soluble oxometal complexes which are; the propensity of certain oxometal species (e.g. Ti'v=O) towards oligomerization to inactive p-oxo complexes, and oxidative destruction of organic ligands. These problems can be circumvented by site isolation of discrete oxometal species in an inorganic matrix, whereby

611

the latter functions as a thermodynamically stable ligand [25]. One approach to isolating metal ions in stable inorganic matrices is to incorporate them, via isomorphous substitution into the framework of a microporous molecular sieve. The most famous titanium containing silicalite was developed by Enichem. [26] for a variety of transformations which employed H202 as the primary oxidant. We have also been active in the area of metal silicalite synthesis and investigated the preparation of chromium(Ill) silicalite (CrS-1) under different conditions for the oxidation of styrene using H202. A problem associated with chromium incorporation is the formation of dimers (2) via hydro bridge linkages at high pH typical during the hydrothermal synthesis of zeolites.

[Cr(H20)e]

OH(~

"= [(H20)4Cr

O

4+ Cr(H20)4]

O

(2)

In order to circumvent this problem, we have used three different approaches to suppress dimerization of chromium species. The methods were; addition of H2SO4, or NH3, or F during synthesis of CrS-I. The catalysts activities were then screened with a model substrate (styrene) in the presence of H202. The scheme below summarizes the reactions taking place during the oxidation. Table 5 compares the CrS-1 preparative routes to product selectivities.

CHO

I

2 eq. H202

iI~" "~----~

+

H2CO

I H2021 1 eq. H202,.~

+

~

+ ]

H

OH II

r

O//L~ III

CHO IV

612

Table 5. The Effect of CrS-1 Preparation on the Oxidation of Styrene in the Presence of H202"

Synthesis Method

Conversion %

F 26 NH3 34 H2SO4 9 asee expenmentalfor details'.

Product Selectivity % I II III IV 57 31' '6 6 52 35 4 9 85 0 1 1 '

The highest conversion was observed for the ammonia route. The catalyst prepared via the addition of H2SO4 gave significantly lower substrate conversion. However, the selectivity to benzaldehyde was relatively high. A cursory examination was made of the effect of solvent on the conversion and product selectivities. The NH3 prepared CrS-1 material was employed for the study. Table 6 summarizes the results. 1,2-Dichloroethane proved to be the best solvent of choice from those screened. It is worth noting that analysis of the mother liquors after reaction indicated that no leaching of the chromium had taken place, thus showing the true heterogenaity of the oxidation system.

Table 6. The Effect of Solvent on the Styrene Oxidation with CrS-1 and H=O=

Solvent Toluene MTBE 1,2-DCE 1,2-DCE 1,2-DCE aSee experimental.

'Temperature ~ 70 70 70 55 40

Conversion % 9 4 34 17 3

Product Selectivity % I II Iii + IV 77 0 8 78 12 10 52 35 13 73 17 14 68 0 24

4. CONCLUSIONS A number of oxygen transfer heterogeneous catalysts have been developed for use with aqueous hydrogen peroxide. The majority of the catalysts discussed are inorganic in nature. The inorganic integrity of the catalysts imparts several advantages over solid organic species. Advantages include an increased thermal, and mechanical robustness. We believe that such immobilized species coupled with H202 are well placed to aid the increasing need for clean technology solutions and implementation of integrated pollution control requirements.

613

ACKNOWLEDGEMENTS We thank the Organic Chemistry Group at Widnes R&D. We also wish to thank our university collaborators, particularly; Professor Bob Johnstone, and Dr. Preciosa Pires (Liverpool University, metal(IV) phosphates), Professor Bill Griffith, and Dr. Melanie Gresley (Imperial College, immobilized polyoxometalates), Professor Roger Sheldon, and Dr. Jihad Dakka (Delft University, chromium silicalite). REFERENCES 1. T. Funabiki, T. Sugimato, and S. Yoshida, Chem. Lett., (1982) 1097. 2. R. A. Sheldon, Bull. Chim. Soc. Belg., 9 (1985) 1450. 3. Degussa, German Patent 2012509 (1975). 4. Chemical Marketing Reporter, 3 (1976). 5. Ugine Kuhlmann, German Patent 2752626 (1977). 6. J. Gordon, Hydr. Process & Petr., (1962) 141. 7. G. Strukel (ed.)"Catalytic Oxidations with H202 as Oxidant", Kluwar Academic Press, The Netherlands, 1992. 8. J. Lieto, D. Milstein, R. L. Albright, and B. C. Gates, Chem. Tech., 13 (1983) 46. 9. T. J. Pinnavaia, and P. K. Welty, J. Am. Chem. Soc., 97 (1975) 3819. 10.W. Heitz, Adv. Polym. Sci., 23 (1977) 1. 11.S. Maur, P. Jayalesky, J. T. Anderson, and T. Matusinovic, Am. Chem. Soc., Symp. Ser., 192 (1982) 43. 12.A.T. Jurewicz, L. D. Roilmann, and D. D. Whitehurst, Adv. Chem. Ser., 132 (1974) 240. 13.R.H. Grubbs, C. Gibbons, L. C. Kroll, W. D. Bonds, and C. H. Bruker, J. Am. Chem. Soc., 95 (1973) 2373. 14.University of Liverpool, Metal Phosphates; Imperial College, Immobilized Heteropolyacids; Delft University, Metal Silicalites. 15.M.T. Pope, "Heteropoly and Isopoly Oxometalates", Springer-Verlag, New York, 1983. 16.C. Venturello, E. Alneri, and M. Ricci, J. Org. Chem., 48 (1983) 3831. 17.Y. Matoba, M. Inoue, J. Akagi, T. Okabayashi, Y. Ishii, and M. Ogawa, Syn. Commun., 14 (1984) 865. 18.S.W. Brown, A. Johnstone, C. W. Jones, A. M. Lee, S. C. Oakes, and S. L. Wilson, Recl. Trav. Pays-Bays., 115 (1996)244. 19.Solvay lnterox, World Patent 9421624 (1994). 20.A. Mangia, Synthesis, 1 (1978) 361. 21.F. DiFuria, and G. Modena, Pure & Appl. Chem., 54 (1982) 1853. 22.Solvay Interox French Patent 2616786 (1988). 23.R.C. Chambers, and C. L. Hill, Inorg. Chem., :)8 (1989) 2509. 24.G.R. Krow, Tetrahedron, 37 (1981) 2697. 25.R.A. Sheldon, "Topics in Current Chemistry", 164, Springer-Verlag, BerlinHeidelberg, 1993. 26.M. Taramasso, and B. Notari, U.S. Patent 4410501 (1983).

This Page Intentionally Left Blank

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

615

C a t a l y t i c w e t air o x i d a t i o n o f w a s t e w a t e r s J.C. B6ziat a, M. Besson a*, p. Gallezot a, S. Juif b and S. Dur6cu b Institut de Recherches sur la Catalyse, 2 Avenue Albert Einstein, 69626 Villeurbanne Cedex, France. b TREDI-Laboratoire de Recherche, Technop61e de Nancy-Brabois, 9 Avenue de la for& de Haye, BP 184, 54605 Vandoeuvre-l~s-Nancy, France. Succinic acid was chosen as a model pollutant to study the catalytic wet air oxidation (CWAO) of aqueous solutions of carboxylic diacids (5g 1-~).Using a 5%Ru/C catalyst, the destruction of succinic acid was rapid and substantially complete within one hour, at 190~ and 5 MPa total pressure. Acrylic and acetic acids were formed as reaction intermediates and were further oxidized to CO 2 and H20. Acetic acid was most difficult compound to oxidize, but at the end of the reaction, only traces could be detected and a total organic carbon (TOC) removal efficiency of more than 99% was observed within 4 hours. No leaching of ruthenium from the catalyst was detected, and the heterogeneous catalyst could be removed by simple filtration. Under the same operating conditions, during the oxidation of succinic acid in solutions containing high salt concentrations (NaC1) or mineral acids, the oxidation of succinic acid was only slightly slowed down, while the oxidation of the acetic acid intermediate was diminished. The rate of oxidation was much slower at neutral or basic pH. Preliminary experiments have demonstrated that CWAO over Ru/C was also feasible for treatment of other carboxylic acids (adipic, glutaric, malonic, propionic and acetic acids). 1.INTRODUCTION Many industrial processes in the chemical and pharmaceutical industries produce waste waters containing organic compounds, which cannot be eliminated by traditional removal technologies, e.g. biological treatment (toxicity of products) or incineration (too low concentrations). Wet air oxidation, the chemical oxidation of organic or inorganic compounds in aqueous phase, at high temperatures (150-315~ and pressures (up to 15MPa), gives biologically nontoxic products or ultimately, carbon dioxide and water. However, energy requirements and installation investments for such processes are excessively high. Different catalytic systems have been developed to operate at lower temperatures and pressures than the thermal process and to improve the rates of aqueous phase oxidation. Homogeneous transition metals (Fe 2+, Cu 2+) may be suitable catalysts, but the dissolved ions need to be separated at the end of the process [ 1, 2]. Attempts have been made to overcome this problem, by using solid oxide catalysts. For example, manganese-cerium and cobalt-bismuth composite oxides were active for the oxidation of many lower carboxylic acids or polyethyleneglycol [3], while copper-zinc oxides were effective catalysts in the oxidation of substituted phenols [4-6] and of p-coumaric acid, a pollutant in olive mill wastewaters [7,8]. These oxides were not very stable in acidic corrosive media and some leaching of the active phases was detected. Precious metals (Pt, Pd, Ru) deposited on supports have been reported to be active for catalytic wet air oxidation (CWAO). Gallezot et al [9] have shown that platinum catalysts supported on carbon could decompose formic, oxalic and maleic acids very easily, at

616

atmospheric pressure and using mild temperatures below 100~ However, maleic acid required 1.5 MPa pressure and Pt/C was almost inactive for the oxidation of acetic acid. The oxidation of an aqueous solution of cyclohexanol performed at 150~ and 5 MPa total pressure on a Pt/C catalyst, resulted in a quite rapid oxidation of cyclohexanol and cyclohexanone to produce saturated diacids (adipic, glutaric and succinic acids), but which did not undergo any further oxidative conversion [10]. Higher temperatures did not result in improvements in the performance of the platinum catalyst. On the other hand, there are indications in the literature, that ruthenium could be a better catalyst. Indeed, Imamura et al [ 11 ] found that ruthenium was the most active catalyst among the precious metals examined in the oxidation of poly(ethyleneglycol), at 200~ under a pressure of oxygen (1MPa) and nitrogen (2MPa). Moreover, a Ru/CeO 2 catalyst had a much higher activity than copper salts for formaldehyde or formic acid oxidation. Duprez et al [ 12] reported that 5% Ru/C was a very efficient catalyst for the WAO of acetic acid, without any leaching of noble metal. Graphite-supported Ru-catalysts were even more active for acetic acid oxidation [13]. In this work succinic acid has been chosen as a representative diacid compound and its reactivity was studied on a 5% Ru/C catalyst.

2. E X P E R I M E N T A L 2.1. Materials and catalyst characterization The catalyst employed in this work was a commercial Ru/C catalyst (Aldrich, ref 20,618-0). Inductive coupled plasma-atomic emission spectroscopy (ICP-AES) was used to measure the ruthenium content in the catalyst after dissolution of the solid in an acidic solution, and for the determination of the concentration of various metal ions in the solution after the oxidation treatment. The sizes of ruthenium particles were measured by high resolution electron microscopy (JEOL JEM 2010).

2.2. Oxidation procedure Oxidation of aqueous solutions of organics (0.5wt%, or 5g 1-~) were performed in a 250 ml Hastelloy C22 autoclave, connected to an air reserve and equipped with a magnetically driven turbine. The reactor was loaded with 150 ml of solution and 1 g of catalyst. After flushing with argon, the temperature of the mixture was raised to the reaction temperature under stirring. Air was then admitted until the preset pressure was attained and the reaction was started by adjusting the stirrer to 1800 rpm, which defined t = 0. Typical operating conditions were 190~ and 5 MPa total pressure. The stirrer speed was maintained at 1800 rpm to minimize mass transfer limitation. The total run time for each experiment was approximately 6 h. Samples were periodically withdrawn from the reactor through stainless-steel 1/16" tubing. All samples were analyzed for pH, TOC (Total Organic Carbon) content and by HPLC for reaction intermediates formed during the reaction. The pH gave qualitative information on the progress of the reaction, since the pH increased upon oxidation of the acids into carbon dioxide. Quantitation of the intermediate products was performed by HPLC with UV and RID detectors mounted in series and an ion-exchange column (Sarasep Car-H), used with dilute U 2 S O 4 solutions as eluent (0.01N, 0.5 ml min-'). The decrease in TOC, which quantifies the overall disappearance of organic compounds, i.e. the degree of decontamination, was measured using a Shimadzu TOC 5050A. The TOC was determined by subtracting the measured IC (inorganic carbon, measure of CO 2 evolved by non-dispersive infrared gas analysis, after acidification of the sample in concentrated phosphoric acid) from measured TC (total carbon, CO 2 evolved after catalytic combustion at 700~ The initial reaction rate ri was calculated both from the disappearance of succinic acid or from the TOC removal with time and was expressed either in molsu c h j mO1R~-1 or in mol c h -1 mO1Ru-i. The overall removal efficiency or TOC abatement can be defined by TOC (%) = 100 x [TOC] / [TOC]in~ti,,~.

617

3. RESULTS AND DISCUSSION 3.1. Oxidation of succinic acid

Preliminary measurements Some preliminary tests were performed on aqueous solutions of succinic acid to evaluate the best operating parameters to carry out this reaction, using the 5% Ru/C catalyst. The study of the effect of temperature up to 200~ indicated a strong temperature dependence of the oxidation rates. Thus, the TOC abatement was more than 99 % after 4 hours at 190~ while it was only 77.5% at 180~ The apparent activation energy, deduced from the Arrhenius plot between 180 and 200~ was ca. 100 kJ mol -~. The influence of the oxygen partial pressure was also examined by varying the total pressure between 3 and 5 MPa. The partial oxygen pressure was calculated, after subtracting the autogeneous pressure of water vapor (1.5 MPa). A TOC abatement of more than 99% was obtained after 6 hours of reaction at 190~ above a total pressure of 5 MPa (i.e. at a partial oxygen pressure of ca.0.7 MPa). From the plot of In (rate) as a function of In (Po2) a positive order of ca. 0.2 with respect to oxygen pressure was calculated. A 0.65 rate order with respect to Po2 has been reported for CWAO of acetic acid on a graphite supported ruthenium catalyst [13]. At a constant pressure of 5 MPa, the catalytic system functioned in a kinetic regime, since a linear variation of the initial oxidation rate with the mass of catalyst up to 1.8 g was observed. In the standard experiments, 1g of catalyst (50 mg Ru) was used. The rates were measured as a function of succinic acid concentration, all other parameters being kept constant. Three different concentrations (0.25, 0.5 and 0.75 wt%) were used. The oxidation rate was found zero order with respect to the succinic acid concentration in the range studied, which means that succinic acid is strongly adsorbed on the metal surface. To eliminate the possible catalytic activity of the reactor walls and internal parts, we verified that under the standard reaction conditions, uncatalysed experiments gave negligible conversion of succinic acid (only 1 1 % conversion after 4 h at 190~ compared to complete conversion within 1 h in the catalyzed experiment). Also, negligible adsorption of the products on the support was verified by measuring the same concentration of succinic acid in solution at 190~ in the presence and absence of support.

Kinetic study under standard conditions Figure l a shows the product distribution vs. time during the oxidation of succinic acid (43 mmol 1-]) at 190~ and a total pressure of 5 MPa (partial oxygen pressure 0.7-0.8 MPa) on lg of 5% Ru/C catalyst. Figure lb gives the TOC removal and the corresponding pH profile for the same experiment. A rapid and linear conversion of succinic acid was observed with initial reaction rates of 15 molsu c h -~ molRu-] and 61 molc h -~ mO1Ru-], resulting in complete conversion within one hour. The intermediate products detected were acrylic and acetic acids (maximum yields 10.5 and 2 mmol l-~, respectively), which were then converted into carbon dioxide and water. Acrylic acid disappeared rapidly and completely during the first hour, but acetic acid, known as a refractory molecule towards oxidation, was decomposed at a lower rate. There was a continuous TOC reduction throughout the course of oxidation with the rate of TOC removal progressively decreasing at the end of the reaction. Nevertheless, more than 99% of TOC removal was measured after 6 h of reaction - only traces of acetic acid were then detected (TOC < 9 mg 1-j, i.e. 0.4 btmol l-J). Malonic acid, oxalic acid or formic acid were not detected by HPLC, probably due to their rapid oxidation. Indeed, separate experiments on the malonic acid (vide infra) and previous results [9] have shown that these acids were oxidized to CO~ and H)O at a very high rate at the present reaction conditions. As expected, the acidity of the solution

618 decreased progressively as the acids were decomposed: from an initial pH of 2.75, the pH increased to a neutral pH (figure 1b).

.Aacrylic acid 9acetic acid 9succinic acid II

14~:

45 4O 35 30E~

12 ~'1 0

i8

2s 20N 15~ 10 5 ,0

6 (D ....4

4

0

2

Time (h)

4

i

6

2500 2000

~

1500

4 ~Z

IATOC l p H [

3

9 1000 500

-~ 0

~ 2

time (h)

A

J, 0

4

6

Figure 1. Oxidation of an aqueous solution of succinic acid (43 mmol.11) over 5 wt% Ru/C: a) yield of succinic acid and intermediate products vs. time and b) TOC removal and pH profile vs. time. Reaction conditions: 190~ air, 5 MPa total pressure.

Effect of initial pH and addition of salts Since real industrial waste waters are liable to contain acids, bases or mineral salts, the effect of initial pH value on the oxidation rate of succinic acid was evaluated. The initial value was adjusted at pH 2, by adding a few drops of chlorhydric or sulfuric acids, and at pH 5 or 12 by adding pellets of sodium hydroxide. Table 2 gives initial reaction rates as well as the TOC abatement and the concentration of acetic acid after 6 h. The acrylic acid formed reached a maximum concentration of 2 mmol 11, whatever the initial pH conditions.

619 Table 2 Influence of the initial pH and addition of NaC1 salt on initial reaction rates, composition and TOC removal (initial TOC ca. 2000 ppm, 190~ 1.5 MPa total pressure). run

Conditions

Initial pH

Final pH

1

reference

2.75

6.72

15

61

100

99.6

traces

2

H2SO 4

2.01

2.12

14

51

100

94.1

4.5

3

HC1

2.07

2.33

13

44

100

95.2

5.5

4

NaOH

5.16

7.65

12

41

86.5

81.6

20.5

5

NaOH

12.02

7.80

3

12

47.15

39.5

21.1

6

NaC1 (6.7g.1 -l) NaC1 (26.7g.1-1)

2.75

4.85

15

57

100

97.1

2.8

2.75

4.75

16

49

100

93.6

5.7

7

r i (SUC) a r~ (TOC) b conv.C % TOC d % ACE

e

a) molsvc h -j mO1Ru-1, b) molc h -1 mO1Ru-1, C) conversion of succinic acid after 6 h, d) % TOC abatement after 6 h, e) concentration of acetic acid (mmol.1-1) after 6 h. In the acidified solutions (runs 2, 3), the initial activity of succinic acid was little affected by the addition of mineral acid, compared to the reference experiment. However, the consecutive oxidation of acetic acid occurred at a lower rate. Consequently, residual TOC increased after 6 h of reaction, but still more than 94 % TOC abatement was achieved. In neutral medium ( run 4) a slight decrease in the initial rate was observed. After 6 h of reaction, total conversion of succinic acid was not achieved, compared to less than 2 hours for the reference experiment. The TOC abatement was 81.6 % at 6 h, because acetic acid was formed in larger amounts and was also barely oxidized. In the basic solution (run 5), the reaction rate was greatly reduced and the amount of acetic acid still increased. These results indicate that the oxidation may occur preferentially on the undissociated forms of the acids (pK~ of succinic acid - 4.16, pK 2 = 5.61), rather than on the carboxylate ions, in agreement with previous results on selective oxidation of aqueous solutions of alcohols over noble metal catalysts. Slightly basic conditions favor the desorption of the acid salt from the surface and prevents C-C bond rupture and over-oxidation, whereas acidic pH favor the adsorption of the carboxylic acid and its further oxidation [ 14-15]. Similar results were observed by Imamura, et al. [ 11 ] in the oxidation of formic acid or acetic acid over 5 % Ru/CeO 2. In order to evaluate the effect of high concentrations of inorganic salts, NaC1 was added to the reaction mixture to obtain solutions containing 6.7 and 26.7 g 1-1 of NaC1 ( NaC1/succinic acid = 2.9 and 11.6, respectively), (Table 2). Comparison of the initial rates clearly shows that the addition of NaC1 had little effect on the disappearance of succinic acid, nor on the initial TOC removal rate. However, the TOC and HPLC measurements after 6 hours, indicate that the presence of the salt decreases the oxidation rate of the intermediate acetic acid.

Analysis of the catalyst and of the solutions after oxidation. TEM images of the catalyst before oxidation showed the presence of particles smaller than 2.5 nm. The oxidation treatment under standard conditions did not change the morphology of the catalyst. The final solutions were analyzed by ICP-AES for the possible leaching of the metal in the reaction mixture. In none of the runs was ruthenium detected within the limits of detection of the method, i.e. 0.05 mg 1-1. This result indicates that the reaction proceeds in a heterogeneous

620 manner. However, the amount of ruthenium on the catalysts increased in the catalysts recovered after filtration of the final solutions, particularly in HC1 solution, where the Ru content increased from 5.12 % to 26.5 %. These results imply that there is a partial oxidation of the support. Such problems have also been reported by Duprez, et al. during oxidation of acetic acid. They found that the amount of CO 2 formed (quantified by GC analysis), was higher than the theoretical amount of CO 2 produced by acetic acid oxidation [ 16]. The analysis of other metallic ions in the final oxidation solutions reveal a corrosion of the stainless steel sampling tube by the acidic waste water. Indeed, the ions detected were Ni, Fe and Cr, the main components of stainless steel. 3.2. Oxidation of other carboxylic acids

Further measurements were performed at the same experimental conditions to compare the relative reactivity of different carboxylic acids. Figure 2 shows the concentration-time profiles for all intermediates detected and quantified during the oxidation of adipic acid or glutaric acid. a

14

45 40

,GLU!

~1 2

i lo

f -t,-i /

~8

mSUC &ACE • ,ADI

xlx

~ 9 6

35~" 30 25,~

20.~ 15 +.a E 10 v 5 0 ~

= 4 L)

o 2

t

0

0

2

Time(h)

4

6

14

45 ~ 4/\~

~,12

"

0 ~..

~

wSUC AACE

4O 35~"

•

30~

:

8 6

.~

4 2 0 0

2

Time (h)

4

6

Figure 2. Product distribution during oxidation of a) adipic acid and b) glutaric acid.

621 At the standard conditions (150 ml of an aqueous solution containing 0.5wt % of organic compound, lg Ru/C catalyst, 190~ 1.5 MPa total pressure), the intermediate products were the same: glutaric acid (in the case of adipic acid), succinic, acrylic and acetic acids. They were formed in the same proportions. The initial rates of TOC abatement were in the order: succinic acid > glutaric acid > adipic acid (figure 3). All reaction products were completely oxidized, resulting in a TOC abatement of more than 99.5% after 6 h. The limiting reaction was again the oxidation of acetic acid formed.

100

~

8O

0 [...-,

60

r,.) 0

40

ISUC 1 oGLU oADI i

20

0

1

2

3 Time (h)

4

5

6

Figure 3. TOC abatement during oxidation of succinic, glutaric and adipic acid under standard conditions. Formic, oxalic and malonic acids were never detected during these oxidation experiments. It was verified that formic and oxalic acids were oxidized so rapidly, that they could not accumulate in the reaction mixture and be detected by HPLC. Malonic acid was decarboxylated very rapidly to yield acetic acid. On the other hand, as expected, acetic acid and propionic acid were much less reactive. The initial rates of TOC removal were 13 and 19 molc h -~ mO1Ru-l, respectively, compared to 61 for succinic acid. After 6 h, TOC abatement was 65.9 and 68.5 %, respectively. CONCLUSIONS Wet air oxidation in the presence of carbon-supported ruthenium provides an efficient method for total destruction by air of organic acid pollutants in aqueous solutions. In the presence of high concentrations of NaCI salts or of mineral acids, the oxidation of succinic acid was not modified, whereas the rate of oxidation of acetic acid formed transiently, was slightly lowered. In neutral and basic media, the oxidation of the carboxylate ions was greatly decreased. No leaching of ruthenium was observed, which means that the reaction was catalyzed by a heterogeneous catalytic system. However, the carbon support was partially oxidized, which limits the application of this catalytic system for the CWAO of acetic acid, which requires temperatures close to 200~

ACKNOWLEDGEMENTS The R6gion Rh6ne-Alpes and the TREDI Company are gratefully acknowledged for the financial support of this project.

622 REFERENCES [1] C.J. Chang, J.-C. Lin and C.-K. Chen, J. Chem. Tech. Biotechnol., 57 (1993) 355. [2] F. Luck, Catalysis Today, 27 (1996) 195. [3] S. Imamura, H. Nishimura and S. Ishida, Sekyu Gakkaishi, 30 (1987) 199. [4] J. Levec, Appl. Catal., 63 (1990) L1. [5] A. Pintar and J. Levec, Chem. Eng. Sci., 47 (1992) 2395. [6] A. Pintar and J. Levec, J. Catal., 135 (1992) 345. [7] D. Mantzavinos, R. Hellenbrand, A. Livingstone and I. Metcalfe, Appl. Catal. B: Environmental 7 (1996) 379. [8] R. Hellenbrand, D. Mantzavinos, A. Livingston and I. Metcalfe, Environmental Catalysis, G. Centi et al (eds), Societh Chimica Italiana, Roma (Italy), 1995, pp. 487. [9] P. Gallezot, N. Laurain and P. Isnard, Appl. Catal. B: Environ., 9 (1996) 11. [10] J.-C. Beziat, M. Besson and P. Gallezot, Appl. Catal. A: Gen., 135 (1996) L7. [11] S. Imamura, I. Fukuda and S. Ishida, Ind. Eng. Chem. Res., 27 (1988) 721. [12] D. Duprez, F. Delanoe, J. Barbier Jr, P. Isnard and G. Blanchard, Catal. Today, 29 (1996) 317. [ 13] P. Gallezot, S. Chaumet, A. Perrard and P. Isnard, J. Catal., in press. [14] P. Fordham, M. Besson and P. Gallezot, Appl. Catal. A, 127 (1995) 165. [ 15] P. Fordham, R. Garcia and M. Besson and P. Gallezot, 1l th International Congress on Catalysis, Studies in Surface Science and Catalysis, J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (eds), Elsevier Science B.V., Amsterdam, 101A, 1996, pp. 161. [16] D. Duprez, F. Delanoe, J. Barbier Jr, P. Isnard and G. Blanchard, Environmental Catalysis, G. Centi, C. Cristiani, P. Forzatti and S. Perathoner (eds), Societh Chimica Italiana, Roma (Italy), 1995, pp. 495.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

623

Catalytic partial oxidation of methanol" H2 p r o d u c t i o n for fuel cells L. Alejo, R. Lago, M.A. Pefia and J.L.G. Fierro Instituto de Cat~lisis y Petroleoquimica, CSIC, Campus UAM, Cantoblanco, 28049 Madrid, Spain; [FAX: +34 1 585 4760; E-mail [email protected]]

In this work Cu/ZnO and Cu/ZnO/AI203 catalysts have been studied for the partial oxidation of methanol with Oz to produce H2. These Cu-Zn based catalysts showed high activity for the partial oxidation of methanol and with activity directly related to the Cu metal area. In the series Cu-Zn with copper relative content of 20-70 wt%, the catalyst Cu40Zn60 (Cu 40 wt% and Zn 60 wt%), which showed the highest Cu area, gave the best results for the partial oxidation of methanol. The activation energies and TOF (turnover frequencies) varied with the Cu-Zn catalyst composition. For catalysts with low Cu loading very high Ea and TOF were obtained (for Cu30Zn70 Ea=482 kJ/mol and TOF c a . 200 min -1 at 497-499 K) whereas for higher Cu contents the E a and TOF decreased tending to constant values (for Cu70Zn30 E a : 71 kJ/mol and TOF= 160 min -1 at 497-499 K). These results are discussed in terms of a possible effect of the Cu-ZnO interaction which depends on the catalyst composition. Catalytic experiments with Cu40Zn55A15 showed that the presence of A1 has an inhibiting effect producing slightly lower methanol conversion. On the other hand, higher selectivities for H2 and CO2 were obtained with only traces of the undesirable CO. Moreover, the A1 is very important for catalyst stability and life-time experiments showed that Cu40Zn55AI5 is stable during the partial oxidation of methanol with no significant change in activity and selectivity even after 110 h of operation at 503 K. Overall, the results seem to indicate that Cu~ is active for partial oxidation of methanol to H2 and CO2 whereas Cu + favors the formation of H20 and CO. Keywords: Methanol oxidation, hydrogen production, copper-zinc catalysts, fuel cells, activity

1. INTRODUCTION The use of H2 gas for mobile fuel cell applications is hindered by problems of storage, safety, refueling, etc. These problems have led to a search for alternative hydrogen storage systems or alternative fuels from which H2 can be produced on-board. From the different liquid fuels considered methanol is an interesting alternative since it is the third largest volume commodity chemical after ethylene and ammonia with a production capacity in excess of 25 million tons, much superior to the actual overall demand [1-3]. The production of H2 from methanol by steam reforming has been extensively studied over supported metal catalysts such as Cu [4-9], Ni [10] and Pd [11]. A limitation of steam reforming route is the endothermicity of the reaction:

624 CH3OH + H20 ~- 3H 2

+ CO 2

AH~

= +49.4 kJ/mol

(1)

Moreover, the need to produce steam makes the process even more energy demanding. The partial oxidation of methanol to produce H2 offers some advantages over steam reforming since it is an exothermic reaction, and hence more favorable thermodynamically and uses 02 (air) instead steam as oxidant, which make it more energy efficient. Moreover, it has been reported that the reaction rate of partial oxidation over copper catalyst is higher than that of steam reforming [ 12,13]. CH3OH + 1/202 ~

2H2 +

CO 2

AH~

= -192.2 kJ/mol

(2)

Preliminary work reported by Kumar et al. from Argonne National Laboratory [3] showed that among several catalysts with different metals and supports the system Cu/ZnO showed the most promising results. In this work, a series of Cu-Zn and Cu-Zn-AI catalysts with different compositions have been prepared and tested under differential reactor conditions for the partial oxidation of methanol to H2.

2. EXPERIMENTAL The catalysts were prepared by inverse precipitation by adding 1.25 M copper, zinc and aluminum nitrate solutions into a 0.25 M Na2CO3 solution under vigorous stirring. The notation used in this work show the wt% of each metal in the catalyst, i.e. the catalyst Cu40Zn55A15 contains 40 wt% of Cu, 55 wt% of Zn and 5 wt% of AI omitting the rest of the elements. The powder XRD patterns were obtained with a Seifert 3000P diffractometer. The calcined precursors were reduced in situ in a stream of U 2 (10 %) in N2 heating the catalyst at 10 K/rain to 503 K where it was kept for 1.5 h. The catalytic tests were carried out with 50 mg catalyst diluted with silicon carbide (500 mg) at atmospheric pressure between 473-503 K with methanol (liquid flow of 4 ml/h) and synthetic air (12 ml/min) to produce a O2/CH3OH molar ratio of 0.06. The reaction products were analyzed on-line with a Varian 3400CX gas chromatograph equipped with a TCD and the columns Porapak Q and molecular sieve 13X. A Micromeritics 2900 instrument was used for TPR/TPO analyses and also for N20 pulses experiments for the determination of Cu metal area, assuming a Cus:O=2:l stoichiometry [14-16]. The BET surface area was obtained in an ASAP-2000 instrument. XPS and Auger analyses were carried out in an ESCALAB 200R instrument with MgKa X-ray source (hv = 1253.6 eV). Scans of the Cu 2p3n, O ls, Zn 2p3a and A1 2p were taken. For the analyses by XRD and XPS after reaction or reduction the catalysts were all quenched to room temperature and immediately drenched in isooctane under nitrogen to avoid oxidation of surface metal copper by exposure to the atmosphere.

3. RESULTS

3.1. Catalyst characterization Table 1 shows the BET and Cu metal areas for the Cu-Zn and Cu-Zn-A1 catalysts after reduction. The BET and metal Cu area for the reduced Cu-Zn catalysts increased with Cu

625 Table 1 BET and Cu ~ metal area for the reduced catalysts Cu-ZnO and Cu-ZnO(AlzO3) Catalyst Cu20Zn80 Cu30Zn70 Cu40Zn60 Cu50Zn50 Cu60Zn40 Cu70Zn30 Cu40Zn55A15 Cu40Zn50All0 Cu40Zn45Al15

BET area (m2/g) 38.2 48.2 51.5 45.3 42.2 26.8 57.0 40.0 33.0

Cu ~ metal area (m2/g) 3.4 4.5 6.8 6.6 6.3 4.7 8.5 5.0

Cu ~ area was determined by N20 chemisorption at 353 K on H2-reduced catalysts at 503 K for 1.5 h in a stream of 10% H2 in N2 content reaching a maximum for Cu40Zn60 and decreased for higher Cu concentration. The presence of AI in the catalyst Cu40zn55A15 increased the Cu ~ area whereas for higher A1 content (Cu40Zn50All0 and Cu40Zn45AllS), a decrease of the metallic area was observed. The TPR profile (not shown here) of the catalyst Cu40Zn60 shows a peak between 450-485 K which corresponds to the reduction of Cu 2§ to Cu ~ This profile is very similar for all the other Cu-Zn catalysts. The presence of A1 in the catalyst causes significant differences in the reducibility of copper. The reduction peak is shifted to higher temperatures, first to 513 K for Cu40Zn55A15 and then a second broad peak appears at 553 K for Cu40Zn50AI10. For higher AI content, the catalyst Cu40Zn45Al15, only the broad peak at 553 K is observed. The XPS results showed the presence of A1 and Zn on the catalyst surfaces with peaks for A1 2p at 74.6-74.7 eV and Zn 2p3/2 at 1022.1-1022.2 eV. From the outgassed fresh catalysts two components were observed at c a . 933.7 eV assigned to CuO [18] and another Cu 2+ species at 935 eV which may be related either to Cu2+-OH or a CuAI204 spinel for the Al-containing catalysts [17-20]. Also a Cu 2+ satellite peak is observed at c a . 943 eV. Upon reduction, the Cu 2+ satellite peak disappears and the Cu2p3/z signal shifts down to binding energies near 932.8 eV which may be assigned to either Cu ~ or Cu +. It is interesting to note that the catalysts Cu40Zn50AI10 and Cu40Zn45Al15 even after reduction showed a small component at 935 eV related to Cu 2+ To differentiate between Cu ~ and Cu + the modified Auger parameter, c~A (Auger parameter plus photon energy) [21-23], was used for the reduced catalysts. The c~A values of 1851 eV obtained for Cu40Zn(60-x) (x - 0,5,10) after reduction indicate the presence of only Cu ~ on the catalyst surface after reduction. On the other hand, for the reduced catalyst Cu40Zn45Al15 two small Auger peaks were obtained of 1851.4 and 1849.6 eV suggesting the presence of Cu ~ and also Cu+. 3.2. Catalytic activity of Cu-Zn systems The Cu-Zn catalysts are very active for the partial oxidation of methanol to produce hydrogen and a typical reaction profile is shown in Fig. 1. At 488 K the reaction takes off and the rates of methanol and 02 conversion increase strongly with temperature to selectively

626 0.4 " 0.3

--

0.2

-

~

~

H2

C02

cn~on

.

o.1

I

H20 CO

0.0 470

480

490 Temperature

500 (K)

Figure 1. Partial oxidation of methanol over the catalyst Cu40Zn60. produce H 2 and C O 2. The 02 conversion reaches 100% at 495 K. The CO production rate is very low throughout the temperature range studied and the H20 formation rate decreases for temperatures higher than 488 K. No other products such as formaldehyde, formic acid, methyl formate and dimethyl ether which are often formed in reactions of methanol in the presence of Cu-Zn-A1 catalysts were detected under the reaction conditions employed. The other Cu-ZnO catalysts with different compositions showed similar reaction profiles. The

C3 e,.

4

0.4

0.2 -

20

30

40

50

60

70

%Cu Figure 2. Rates of methanol conversion and H 2 and CO2 formation at 497 K and Cu metal area versus the Cu content in the Cu-Zn catalysts.

627 effect of Cu content on methanol conversion and H 2 and C O 2 formation can be seen in Fig. 2. Methanol conversion to H2 and CO2 increases with Cu content reaching a maximum with the catalyst Cu40Zn60 and decreases for higher Cu loadings. A positive relationship between the methanol partial oxidation and the Cu metal area was observed. The catalyst Cu40Zn60 which had the highest Cu metal area was the most active and selective for the partial oxidation of methanol. If the catalysts were not reduced prior to reaction very low activity for methanol conversion ( < 0.03 mol/g.h) resulted producing CO2, H20 and only traces of H2. Also in blank experiments over pure ZnO under the same reaction conditions the methanol conversion was very low (ca. 0.025 mol/g.h) between 473-503 K. The apparent activation energies (Ea) for the partial oxidation of methanol depended on the Cu-Zn catalyst composition. The catalyst Cu30Zn70 showed a high value of activation energy (482 kJ/mol). As the copper content increased to 70%, E, decreased tending to a value of ca. 71 kJ/mol. From the Cu area the turnover frequency (TOF) was calculated for the different Cu-Zn catalysts. The TOF values at two temperatures, 497 and 499 K, showed that, like the E,, the TOF was also higher for low Cu loading (ca. 250 min -1 for Cu20Zn80) and decreased, apparently tending to constant values near 160 rain -t as the Cu content increased. 3.3. Catalytic activity and stability of Cu-Zn-AI systems It was found that the presence of small amounts of A1 (Cu40Zn55A15) slightly decreased the catalytic activity for the partial oxidation of methanol. Higher A1 contents (Cu40Zn50A110 and Cu40Zn45AI15) result in strong reduction in activity. On the other hand, the presence of A1 greatly enhanced the catalyst stability. The results of life-time experiments with two catalysts, Cu40Zn60 and Cu40Zn55A15, for a period of 110 h at 503 K showed that the binary catalyst Cu40Zn60 deactivated rapidly with methanol conversion decreasind from 0.21 mol/gh to approximately 0.12 mol/gh, selectivity for H2 and CO2 decreasing less. On the other hand, no significant deactivation or loss of selectivity in the ternary catalyst Cu40Zn55A15 was observed even after 110 h on-stream at 503 K.

3.2. Effect of residence time Residence time studies were carried out with 50 mg catalyst at 495 K by increasing the total flow rate while keeping constant the partial pressure of all the components, i.e., CH3OH , O: and N: (Figs. 3a and 3b). Figure 3 shows that the selectivities for H 2 and CO2 increased with residence time. This indicates that part of the H 2 and CO2 are not primary products but are being formed a posteriori by secondary reactions. The only other products observed under the reaction conditions were HzO and CO. By extrapolating to zero residence time it can be estimated that approximately 65 % of CO2 and 35 % of H 2 are formed primarily. It suggests that possible primary reactions taking place are the combustion (ca. 65 %) producing CO2 and H20 (Eq. 3) and the decomposition of methanol (ca. 35 %) to CO and H 2 (Eq. 4): CH3OH + 2 0 2 ~ - - C O 2 -}- 2 H 2 0

(3)

CH3OH ~ CO + 2 H 2

(4)

A possible secondary reaction involved in the formation of H 2 and C O 2 at the expense of CO and HzO is the water gas-shift reaction (WGS):

628

~"~ 0"16f

9

~87

i0-14

f

"~ ~ 0.12[

0.90

|

1.00 W/F (mg/min mi)

F

1.10

Figure 3. Methanol and Oz conversions (a) and catalyst at 495 K.

CO + H20 ~-- CO 2 + H2

H2

I

0.90

,

I

,

I

~

I

1.00 W/F (mg/min ml)

,

1.10

and CO2 selectivities (b) over a Cu/ZnO

(5)

Oxygen conversion increased slightly with residence time while methanol conversion increased much more strongly. This indicates that as the residence time increases methanol is being converted by reactions where oxygen does not participate significantly. These reactions can be, for example, methanol decomposition [6] and steam reforming [24].

3.4. Pulse experiments For the pulse experiments the reduced Cu-Zn catalyst was kept under He flow al 503 K and subjected to different sequences of pulses of methanol(2 %)/He and 02(2 %)/He with the reaction products monitored by mass spectrometry. Figure 4 shows the results for a sequence of 5 pulses of methanol followed by a sequence of CH3OH and 02 pulses. It can be observed that, in the absence of oxygen, methanol decomposes to produce mainly H2 and CO. Small amounts of CO2 are also formed probably due to the presence of strongly adsorbed H20 and OH groups generated during the reduction of the catalyst. Also, despite the very careful procedures to dry methanol a very small amount of H20 was still present in the CH3OH pulse. After 10 pulses of methanol an alternating sequence of CH3OH and 02 pulses was applied to the catalyst and the results are displayed in Fig. 4b. In the first CH3OH pulse the decomposition to H 2 and CO is dominant and only small signals for CO2 and H20 were observed. However in the second pulse of methanol, after the catalyst was exposed to a 02 pulse, the partial oxidation to H2 and CO2 was the main reaction. The intensity of the CO2 signal strongly increased whereas the H20 signal was only slightly more intense and no change was observed for H2. This suggests that O2 can be chemisorbed on the catalyst surface and generate species which are active and stable for the partial oxidation of methanol.

629

v CH3OH ,

~

~

P ~

*

eOH ~O 2 pulse I

,

[

m/e=31

,

m/e=32

r'-__

r

H2 (x0.2) CO C02 ~ 0

15

30 45 Time (rain)

60

!

0

~

I

15

~

I

,~

I

30 45 Time (min)

,

---

60

Figure 4. (a) Pulses of methanol (2%) in He on the Hz-reduced catalyst, followed by (b) alternating sequence of methanol and Oz pulses on the catalyst at 503 K.

3.3. Effect of oxygen partial pressure The effect of the partial pressure of oxygen on the partial oxidation of methanol was studied at 488 K and the results are shown in Fig. 5a-c. It can be observed that as the 02 partial pressure increases from 0.02 to 0.05 bar the methanol conversion and the rate of H2 and CO2 formation increased. It was found that although the rate of H2 formation increased the selectivity for H2 decreased. This indicates that in this Po2 range (0.02-0.05 bar) the chemisorbed O2 on the catalyst surface promotes methanol conversion forming H2 but also the non-selective oxidation leading to H20 which becomes more pronounced. If Po2 is further increased (> 0.05 bar) a sharp decrease in CH3OH conversion and in the production of H2 and CO2 is observed. In the Po2 range between 0.1 and 0.2 bar no significant change was observed. Interestingly, if the 02 partial pressure is then decreased it was found that the catalyst does not recover the initial conversion and H2 and CO2 production. This irreversible decrease in activity suggests that at 488 K and Po2 higher than 0.05 bar the catalyst oxidizes, and this might have been responsible for the large decrease in methanol conversion and H2 and CO2 production. The unreduced catalyst, i.e., copper oxide (CuO/ZnO), showed similar results with very low methanol conversion (<0.03 mol/g.h) producing CO2, H20 and only traces of H 2. Temperature-programmed oxidation (TPO) experiments on the Cu/ZnO catalyst carried out with an O2/He mixture with Po2 of 0.03 bar showed that at 488 K more than 70% of the copper in the sample was already oxidized. Although under these reaction conditions the catalyst seemed to be oxidized at Po2 higher than 0.05 bar it was demonstrated that in the integral regime of operation a Po2 of 0.21 bar, can be used without any apparent catalyst deactivation [25]. It seems that as long as the oxygen conversion during the partial oxidation of methanol is kept high (near 100%) no oxidation of the copper surface takes place.

630 4. DISCUSSION In the Cu-Zn series the catalyst with composition Cu40Zn60 showed the highest copper metal area (6.8 m2/g) and also the highest activity for the partial oxidation of methanol to H2 and CO2. Figure 2 shows that catalyst activity increases with copper surface area suggesting that the exposed copper is essential for an active catalyst. Preliminary studies with an used Cu-ZnA1 catalyst by XPS and Auger spectroscopies quenched and carefully isolated after reaction [25] showed that the catalyst surface was composed solely of Cu ~ and in some cases a very small amount of Cu2§ was also detected. It is interesting to note that in the series Cu-Zn catalysts the activation energy and the TOF for the partial oxidation of methanol depended on the copper content. These results suggest that the interaction Cu-ZnO is very important for a catalyst active for the partial oxidation of methanol. In the Cu-Zn-A1 catalyst series it was observed that the presence of AI has a strong effect on the reducibility of copper. In the presence of AI the TPR copper reduction peaks were shifted to higher temperatures and under the reduction conditions employed part of the copper

0.15

HaOH

0.10 0.05 I

!

0.20 n 2

0.,0

_

0.15

C

C02 0.10 0.05 0

I 0

i 0.10

I

I 0.20

02 partial pressure (atm)

Figure 5. Effect of 0 2 partial pressure on the partial oxidation of methanol at 488 K: (a), methanol conversion; (b), H2 production and (c), CO2 production.

631 in the catalyst was not reduced. This behavior may be related to the formation of Cu-AI compounds such as the spinel CuA1204. XPS analyses showed a Cu 2+ signal which might be related to the spinel CuA1204. It was noted for the calcined catalysts that as the AI content increased the color change from black for Cu40Zn55A15 to black-greenish for Cu40Zn45Al15 (the CuAI204 is green in color). The presence of AI in the Cu-Zn catalyst has an inhibiting in the partial oxidation of methanol. It can be observed that Cu40Zn55AI5 shows much lower TOF compared to Cu40Zn60. This inhibiting effect of AI has also been observed for the methanol synthesis [26]. In the case of Cu40Zn50AI10 and Cu40Zn45Al15 the low activity for the partial oxidation of methanol may be related to the presence of part of the copper in oxized form in the catalyst. Although the presence of Al decreases the catalytic activity, it has an important effect in improving its stability. The catalyst Cu40Zn55A15 was stable for the partial oxidation of methanol at 503 K with no significant change in activity even after 110 h onstream. On the other hand, the catalyst Cu40Zn60 deactivates rapidly, and after 20 h reaction showing a decrease in the selectivities for He and CO2. Moreover, the Cu40Zn55A15 catalyst is more selective for H2 and CO2 formation and only traces of CO were detected during the reaction. The Cu40Zn60 catalyst on the other side forms significant amounts of CO. The mechanism of the partial oxidation of methanol to H2 and CO2 can be very complex due to the possibility of many different reactions occurring simultaneously. Some of these processes are the methanol decomposition, combustion, water gas shift, etc. The residence time studies indicated that possible reactions taking place primarily are the combustion forming CO2 and H20 and the decomposition to CO and H2. There is no evidence to dispute that the partial oxidation of methanol can occur as a primary reaction. In this respect, the partial oxidation of methanol is similar to the partial oxidation of methane to synthesis gas (CH4 + 1/2 02 ~ CO + 2 H2) where a complex network, involving secondary reactions of CO2 and H20 with CH4, enhance the yield of H 2 and CO so that thermodynamic equilibrium can be reached easily [27]. It can be inferred that, like the oxidation of methane to syngas, the partial oxidation of methanol with a CH3OH/O2 ratio near to the stoichiometric value (ca. 2/1) operating with an optimum residence time and temperature can reach very high yields also approaching thermodynamic equilibrium. The pulse experiments showed that the catalyst can operate in a cyclic mode. Molecular oxygen seems to chemisorb to form active and stable species on the catalyst surface which can oxidize methanol to selectively produce H2 and CO2. The concentration of these oxygen species on the catalyst surface seems to be crucial for determining the reaction products. If the reaction is carried out in the absence of 02, methanol decomposition to CO and H2 is the main process as observed by the pulse experiments and also in previous work [25]. If oxygen is present the main reaction products are H2, CO2 and also H20. As the oxygen concentration increases the methanol conversion via the partial oxidation also increases but the formation of H20 becomes more favorable compared to H2. If the catalyst is subjected to high 02 partial pressure and high temperature it deactivates probably due to the oxidation of the copper active phase.

Acknowledgements LAS is grateful to University of San Simon, Cochabamba (Bolivia), and to ICI for a scholarship grant. RML acknowledges the grant from the Ministerio de Educaci6n y Ciencia, Spain. The partial support by the CICYT (Grant No. MAT95-0894) is gratefully acknowledged.

632 REFERENCES .

2. 3. .

5. 6. .

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10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27.

R. Kumar, S. Ahmed and M. Yu, Am. Chem. Soc. Div. Fuel Chem., 38 (1993) 1741. Hydrocarbon Processing, Sept. 31 (1995). R. Kumar, S. Ahmed, M. Krumplet and K.M. Myles, Argonne National Laboratory Report, ANL-92/31, Argon, IL, 1992. E. Santacesaria and S. Carra, Appl. Catal., 5 (1983) 345. H. Kobayashi, N. Takezawa and C. Minochi, J. Catal., 69 (1981)487. C.J. Jiang, D.L. Trimm, M.S. Wainwright and N.W. Cant, Appl. Catal. A: General, 93 (1993) 244. C.J. Jiang, D.L. Trimm, M.S. Wainwright and N.W. Cant, Appl. Catal. A: General, 97 (1993) 145. H. Kobayashi, A. Hirose, M. Shimokawabe and K. Takezawa, Appl. Catal., 4 (1982) 127. J. Barton and V. Pour, Coll. Czech. Chem. Comm., 45 (1979) 3402. H. Kobayashi, N. Takezawa and C. Minochi, in Preparation of Catalyst III, Elsevier, Amsterdam, 1983, p. 697. N. Iwasa, S. Kudo, H- Takahashi, S. Masuda and N. Takezawa, Catal. Lett., 19 (1993) 211. T.J. Huang and S.W. Wang, Appl. Catal., 24 (1986) 287. T.J. Huang and S.L. Chren, Appl. Catal., 40 (1988) 43. K.J. Soerensen and N.W. Cant, Catal. Lett., 33 (1995) 117. J.W. Evans, M.S. Bridgewater and D.S. Yong, Appl. Catal., 31 (1983) 75. T.A.J. Osinga, B.G. Linsen and W.P. van Beek, J. Catal., 7 (1967) 277. A. Sepulveda, C. Marquez, I.R. Ramos, A.G. Ruiz and J.L.G. Fierro, Surf. Interf. Anal., 20 (1993) 1067. A. Wolberg, J.L. Ogilvie and J.F. Roth, J. Catal., 19 (1970) 86. R.B. Friedman, J.J. Freeman and F.W. Lyttle, J. Catal., 55 (1978) 10. G. Ertl, R. Hierl, R. Kn6zinger, N. Thiele and H.P. Urbach, Appl. Surf. Sci., 5 (1980) 49. C.D. Wagner, L.H. Gale and R.H. Raymond, Anal. Chem., 51 (1979) 466. Y. Okamoto, K. Fukino, T. Imanaka and T. Teranishi, J. Chem. Soc., Chem. Comm., (1982) 1405. T.H. Fleish and R.L. Mieville, J. Catal., 90 (1984) 165. K. Takahashi, N. Takesawa and H. Kobayashi, Appl. Catal., 97 (1982) 363. L.A. Espinoza, M.A. Pena and J.L.G. Fierro, in Proceedings of the 15th Iberoamerican Symposium on Catalysis, Cordoba, Argentina, September 1996. S. Gusi, F. Trifiro, A. Vaccari and G. Piero, J. Catal.,94 (1985) 120. S. Fujita, S. Matsumoto and N. Takezawa, Catal. Lett., 25 (1994) 29.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

633

Catalytic liquid-phase phenol oxidation over metal oxides and molecular sieves. Reaction kinetics and mechanism Albin Pintar*, Gorazd Ber6i6, Jurka Batista and Janez Levec Laboratory for Catalysis and Chemical Reaction Engineering, National Institute of Chemistry, Hajdrihova 19, P.O. Box 3430, SI- 1001 Ljubljana, Slovenia

Oxidation of aqueous phenol solutions was studied over various catalysts in a semibatch slurry and continuous-flow fixed-bed reactors at temperatures up to 463 K and pressures slightly above atmospheric. The results show that due to a complex consecutive-parallel reaction pathway and a heterogeneous-homogeneous free-radical mechanism both kinetics and reaction selectivity are strongly dependent on the type of reactor used. Although the catalysts employed were found to be active in converting aqueous phenol solutions to nontoxic compounds, neither metal oxides nor zeolites were stable at the reaction conditions.

1. INTRODUCTION Liquid-phase oxidation of organics by oxygen using solid catalysts is a promising technique for the removal of dissolved toxic pollutants from industrial effluents [1-3]. In this process, organics are oxidized to CO2, H20 and associated inorganic species (e.g., halogenides, sulfate, nitrite and/or nitrate) in a three-phase reactor which operates at milder reaction conditions (i.e., temperatures below 423 K and pressures up to 20 bar) than in the uncatalyzed thermal process. For optimal effect, a catalytic liquid-phase oxidation process should be installed at the source of waste water production in an industrial plant. Since this purification technique is at various stages of development [3], it is obvious that more kinetic and mechanistic studies with different systems in aqueous solutions are needed. The key issue in effective catalytic oxidation of organics is finding a suitable catalyst. Oxidation of aqueous phenol solutions by using different transition metal oxides as heterogeneous catalysts is already known [4-6]. On the other hand, the potential of molecular sieves to catalyze oxidative phenol destruction has not been examined yet. The objective of this contribution is to provide kinetic and mechanistic data on the catalytic liquid-phase oxidation of aqueous phenol solutions obtained in the presence of various transition metal oxides and molecular sieves. The reaction was studied in a semibatch slurry as well as twoand three-phase continuous-flow reactors. Another matter of concern was the chemical stability of catalysts under the reaction conditions. *Corresponding author. Fax: (+386-61) 12 59 244; e-mail: [email protected].

634 2. EXPERIMENTAL

The CuZSM-5 catalysts (Si/AI: 63 [7] and 143) were prepared from the corresponding (Na,TPA)-ZSM-5 precursors by calcination (air stream, 4 h at 823 K) and ion exchange at room temperature in an aqueous copper nitrate solution (1 M) for 24 h under continuous stirring. After exchange the CuZSM-5 zeolites were filtered, washed, dried and calcined again (air stream, 1 h at 823 K). The CuY catalyst was produced from the NaY, FAU zeolite (VEB, Chemiekombinat Bitterfeld, Si/AI: 2.43) by a similar ion exchange procedure. The exchanged sample was filtered, washed and dried. The CoAPO-5 catalyst was prepared by hydrothermal crystallization at 443 K from a reaction mixture with the following composition: 1.5 Et3N 0.10 CoO - 0.9 AI203 - 1.2 P205 - 0.20 HC1 - 50 H20. The (Et3N)-CoAPO-5 precursor was calcined (air stream, 4 h at 623 K and 6 h at 823 K), treated with ammonium nitrate solution (1 M) for 5 h at 353 K, filtered, washed, dried and calcined again (oxygen stream, 6 h at 823 K). The structural type, degree of crystallinity and phase purity of the samples were determined by X-ray powder diffraction analysis, the particle size and potential presence of impurities by scanning electron microscopy and the composition of the catalysts by chemical analysis. The chemical composition and the BET surface area of the catalysts used in this study are given in Table 1. The G-66A catalyst (Siid-Chemie AG, Munich) and the Phillips catalyst (Phillips Chemical Co., Oklahoma) were pretreated for 2 h at 1133 K in an oxygen stream and then cooled to room temperature. XRD examination of the calcined G-66A catalyst revealed the formation of Cu20, CuO and two spinels: ZnAI204 and CuA1204. The latter was present only in small quantities. The catalyst particles were then crushed, sieved and those with sizes below 60 ktm were used in the reaction.

Table 1 Chemical composition and BET surface area of catalysts used in this study Catalyst Composition, wt. % CuZSM-5 (Si/Al: 63) 96.9 % SiO2, 1.3 % A1203, 1.8 % CuO CuZSM-5 (Si/Al: 143) 98.5 % SiO2, 0.6 % A1203, 0.9 % CuO CuY 62.4 % SiO2, 21.8 % A1203, 3.8 % Na20 and 12.0 % CuO CoAPO-5 40.2 % A1203, 57.5 % P205, 2.3 % CoO G-66A 42.4 % CuO, 47.5 % ZnO, 10.1% AI203 Phillips ZnA1204promoted by 4 % Cu, 2 % Mn, 1% La EX- 1144.3 9.3 % CuO, 6.9 % ZnO, 1.4 % CoO on support EX- 1144.8 7.4 % CuO, 9.0 % ZnO, 3.2 % CoO on support

SBET,m2/g N.D. 362 417 208 11 89 12.5 11.5

The activity tests of liquid-phase oxidation of aqueous phenol solution were conducted in a semibatch slurry reactor at operating conditions given in the caption of Figure 1. The experimental apparatus, the procedure of these measurements and the analysis of the reaction samples are described in detail in a preceding paper [6]. Additional kinetic and mechanistic investigations were carded out in an isothermal, differentially operated "liquid-saturated" fixed-bed reactor [8, 9] which was packed with a pretreated EX-1144.3 catalyst (Siid-Chemie

635 AG, Munich). Flow rates of the liquid-phase in this experimental set-up were in a range of 310 mL/min. Catalytic liquid-phase phenol oxidation was also performed at integral conditions in a pilot trickle-bed reactor described elsewhere [10]. In the latter case, an EX-1144.8 catalyst (dp: 0.6 mm) calcined at 1133 K was employed.

3. RESULTS AND DISCUSSION The activity tests of catalysts for liquid-phase phenol oxidation were carried out at operating conditions for which both heat and mass transfer resistances were negligible. Residual phenol concentration as a function of time for some typical oxidation runs is illustrated in Figure 1.

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Figure 1. Relative activities of some experimental and commercial catalysts for oxidation of aqueous phenol solutions in a semibatch slurry reactor. Operating conditions: T: 403 K; p(O2): a, feed (02): 1 L,/min. 5.6 bar; N: 1000 rpm; meat.: 6.5 g; dp: <60 ~tm; CphoH,initial: 0.053 mol/L; "*'vol.

The activity exhibited by the pretreated G-66A catalyst could be attributed to the presence of a Cu+/Cu2+ redox couple in a spinel matrix. This is also in agreement with previous studies where catalytic activity of mixed oxide catalysts with a spinel structure in the phenol oxidation was found to be higher than that of the pure oxides [1 l, 12]. On the other hand, the specific catalytic activity of molecular sieves is limited by the concentration of isolated redox centres located in the microporous voids or in the framework. Although molecular sieves have a much lower concentration of metals (Table 1), they show fairly good activity for liquid-

636

phase phenol oxidation in comparison to the representative metal oxide catalysts. In the light of these observations, it can be concluded that the important factors determining the catalytic properties of these solids are primarily surface structure and composition, and to a minor extent, their bulk composition. Furthermore, the higher Si/A1 ratio of CuZSM-5 sieves makes the character of the catalyst particle surface more hydrophobic, thus increasing the amount of phenol adsorbed and subsequently its disappearance rate. Experimental evidence common to all the catalysts employed demonstrates that the reaction obeys a consecutive-parallel reaction pathway; 1,2 and 1,4 hydroxylated benzenes, benzoquinones, and carboxylic acids are formed during the reaction course towards carbon dioxide production. In addition, the measured reaction rate was affected by the pH value of aqueous solution, and radical as well as inhibitor additions. Based on these facts, a heterogeneous-homogeneous free radical mechanism is proposed. The catalysts probably played a role in the activation of both reactants, i.e., phenol and oxygen. The temporal course of CO2 formation during the catalytic phenol oxidation carried out in a slurry reactor is shown in Figure 2. It can be seen that only about one half of the initial organic content is converted to CO2. At the end of the oxidation run, the rest of the carbon is found in the form of polyaromatic condensates strongly adsorbed on the catalyst surface, as was confirmed by

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Time, rain Figure 2. Percentage of CO 2 formed during the course of catalytic liquid-phase phenol oxidation carded out in a semibatch slurry reactor.

CHN analysis and 13C CPMAS NMR examinations. The total amount of CO 2 produced in the presence of catalyst is in good agreement with the values obtained in the uncatalyzed liquid-

637 phase phenol oxidation [13]. These results suggest that in a slurry reactor the degenerate branched chain mechanism prevails which favors oligomerization and oxidative coupling reactions. The selectivity of the investigated reaction (i.e., carbon dioxide formation) was markedly influenced by the liquid-to-solid volumetric ratio. In fixed-bed reactors (e.g., "liquid-saturated", trickle-bed) this ratio is rather low; correspondingly, propagation reactions taking place in the liquid-phase are suppressed, which leads to the almost quantitative transformation of phenol to carbon dioxide. A thorough kinetic investigation carried out in an isothermal, differentially operated "liquid-saturated" fixed-bed reactor [8, 9] shows that the phenol disappearance rate can be well described by a Langmuir-Hinshelwood (L-H) kinetic model. The proposed intrinsic rate expression

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accounts for model pollutant (poll.) as well as dissociative oxygen adsorption processes on different types of active sites. The adsorption-desorption steps of reactants were determined to be fast and thus at equilibrium. Additional investigations of structure-reactivity relationships have confirmed that the activation of reactants on the catalyst surface occurs via a stepwise oxidation mechanism. On the basis of experimental results, it is concluded that the disappearance rate of phenol and substituted derivatives is dependent on: (1) precursor surface complex formation, and (2) subsequent homolytic electron transfer rate, the step in which phenoxy radicals are produced. The hydroxyl hydrogen radical abstraction from surface phenol complexes, which occurs directly by valence band holes, or indirectly by trapped holes at the particle surface, was found to be the rate controlling step. In the reaction system investigated, oxygen acts as a primary scavenger of electrons in the conduction band. The observed reaction order of 1/2 with respect to oxygen is consistent with oxygen being activated to form surface bound O- and hydroxyl radicals. These species react further with phenoxy species adsorbed on adjacent surface sites, thus enabling total oxidation of the parent molecule. The validity of the proposed L-H kinetic model is shown in Figure 3, where good agreement between the measured and the calculated disappearance rates of model pollutants is achieved. The oxidation rates decrease as the Hammer cr constants of the para substituents increase, indicating that phenols with electron-withdrawing substituents are more resistant to oxidation than phenols with electron-donating substituents. Similar trends as discussed above have been observed for photocatalyzed oxidations of many organic pollutants; the dual-site LH type of rate equations have been successfully employed in simulations of observed oxidation kinetics [14]. In a kinetic investigation of the catalytic liquid-phase phenol oxidation carried out in a semibatch slurry reactor [6], it has been found that homogeneous stepwise polymerization reactions are enhanced in the bulk liquid-phase due to the high liquid-to-solid volumetric ratio. The rate of phenol disappearance has been expressed on the basis of power-law kinetics as a sum of heterogeneous and homogeneous (polymerization) contributions, thus 1/4 + khom." CphOH .ZC(Pn ) (-rphOH) = khet." CphOH "~O2

(2)

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(-r~l.)exp.-lO 6, mol/(gca t h) Figure 3. Comparison between calculated and measured model pollutant disappearance rates evaluated in a differentially operated "liquid-saturated" fixed-bed reactor at different temperatures and oxygen feed concentrations.

where khet. is the apparent rate constant for the heterogeneous oxidation steps, and khom. is the lump polymerization constant including initiation and propagation steps [6]. Considering both the above mentioned experimental facts obtained in a differential fixed-bed reactor, and the role of a solid catalyst in the liquid-phase phenol oxidation, it is evident that Eq. (2) requires some modifications. To describe precisely the temporal course of the catalytic phenol oxidation in a slurry reactor, a new kinetic model has been derived, which is represented by the following rate equation K1/2 .cl/2 k K K [/2~ ~_PhO___H c 9cl/2 _'r'app___z"9?O___U_"~__O__ O~ + "ZC(Pn) khom. 9 02 ) CphOH 1+ K~o22 9c 1/202 (1 + KphoH"CphOH) "(1+ K~O22 C1/2 02

02

(3)

Eq. (3) is presented in Figures 4 and 5 by solid curves. Good agreement between measured and calculated phenol concentrations is obtained. It is interesting to point out that the second right-hand term of Eq. (3) exhibits a Langmuir-type concentration dependence with respect to oxygen concentration (or partial pressure), which indicates that the activated oxygen species (such as O~, HOO ~ H202 and OH ~ possibly formed during the reduction of adsorbed

639

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Time, min Figure 4. Phenol concentration as a function of time obtained in a semibatch slurry reactor at various reaction temperatures.

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100

Time, min Figure 5. Phenol concentration as a function of time obtained in a semibatch slurry reactor at different oxygen partial pressures but constant reaction temperature.

640 oxygen by conduction band electrons play a significant role in oxidative coupling reactions. Since no oligomers and polymers were observed during the phenol oxidation runs conducted in a differentially operated fixed-bed reactor, it can be tentatively concluded that due to the immobilization of a catalyst, diffusion of radicals into the liquid-phase is markedly reduced or even negligible [14]. Long-term experiments of the catalytic liquid-phase phenol oxidation carried out in an isothermal trickle-bed reactor packed with pretreated EX-1144.8 catalyst demonstrate that in a temperature range of 403-423 K and pressures up to 12 bar about 95 % of phenol feed content could be selectively converted to carbon dioxide. Furthermore, Figure 6 shows typical phenol and TOC conversions as a function of catalyst bed length, obtained in the trickling flow regime. It can be seen that due to small differences between the phenol and TOC conversions, small amounts of aromatic and aliphatic hydrocarbons are accumulated as intermediates in the liquid-phase. These observations and the fact that the pH value of the aqueous solution does not change significantly along the catalyst bed length, confirm that the catalyst used is efficient also for oxidation of intermediate (C-2, C-3 and C-4) carboxylic acids. In the off-gas, no carbon monoxide was detected at any operating conditions. 60

60

po2:7.0bar

rr~: 769 g EX-1144.8;dp: 0.6 rrrn Cpto.kfeed:0.032n't31/L ~ 4 0 _ Uquidflow rate: 1.5 L/h Oxygenflow rate: 1 Ln/min

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Figure 6. Phenol and TOC conversion as a function of catalyst bed depth in a trickle-bed reactor (L: 88 cm; D: 3.4 cm).

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TBR model

3O 60 Cat~yst bed lereU~, cm

9O

Figure 7. Measured and predicted phenol conversion profiles along the catalyst bed in a trickle-bed reactor.

Liquid-phase phenol oxidation runs conducted in a trickle-bed reactor were simulated by means of Eq. (1) and heterogeneous models assuming plug flow in the gas-phase, and plug flow or dispersive flow in the liquid-phase. In both cases, due to sufficiently high D/dv and L/dv ratios predicted conversion vs. bed length profiles were practically the same. Good agreement between the measured and predicted values for different liquid flow rates (Figure 7) was achieved only when higher catalytic activity than obtained by intrinsic kinetics

641 measurements in a differential fixed-bed reactor was accounted for. During the operation of a trickle-bed reactor in the trickling flow regime, the catalyst surface is only partially wetted, which means that particles are directly exposed to the gas-phase. Consequently, it is likely that the concentration of bulk and surface trapped valence band holes (i.e., active sites for phenol oxidation) in comparison to those of slurry and differential "liquid-saturated" reactors increases, which in turn accelerates the kinetics of the rate-determining step of hydroxyl hydrogen radical abstraction. The influence of gas-phase composition on the concentration and nature of surface defects, and thus on the reaction rate has already been reported for gassolid oxidations [15]. Based on the parametric analysis performed, it is evident that residual phenol concentration is slightly dependent on the oxygen transfer from the gas-phase into the bulk liquid-phase; the liquid-to-particle mass transfer coefficients exhibit no impact on phenol disappearance rate. The performed calculations also demonstrate that intermediates, even aromatics, and final products have no retarding effect on the phenol disappearance rate. Finally, the chemical stability of the catalysts employed in this study was tested by means of XRD and EDXS analyses. The examination of fresh and used catalysts shows that during the reaction course metal ions are slowly leached into the aqueous solution, which can be attributed either to the temperature of operation or the presence of complexing carboxylic acids and benzoquinones in the liquid-phase. Contrary to the results obtained in continuousflow fixed-bed reactors [8, 9], the extent of catalyst dissolution in the slurry reactor was considerable. This is probably due to the higher accumulation of benzoquinones which are known to form stable complexes with metal ions. Examination of the X-ray powder diffraction patterns of the molecular sieves before and after the liquid-phase phenol oxidation

......... fresh used

i

c~

ti

r .. ..

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•

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20 30 20, degrees

i

40

50

Figure 8. X-ray powder diffraction patterns of molecular sieves before and after liquid-phase phenol oxidation performed in a semibatch slurry reactor at 403 K.

642 in the slurry reactor indicated that the reaction conditions employed here markedly affect the structure of hydrophilic CuY zeolite (Figure 8); it is obvious that an amorphous phase is formed. This effect was less pronounced for CoAPO-5 and hydrophobic CuZSM-5 molecular sieves. The intensities of some peaks were lowered also due to coke deposition on the catalyst surface, which was confirmed by carbon content and BET analyses of spent molecular sieves. Similar observations as above have also been given recently for liquid-phase phenol oxidation with H202 over FeZSM-5 zeolites [16]. To conclude, it should be emphasized that the presence of dissolved copper, zinc and cobalt ions (acting as Lewis acids) in the aqueous solution of a slurry reactor might influence the phenol disappearance rate [6, 17], and consequently the phenol concentration vs. time dependencies depicted in Fig. 1.

4. CONCLUSIONS The catalytic liquid-phase oxidation of aqueous phenol solution, carried out in a variety of reactor systems, demonstrates that phenol can be transformed to nontoxic compounds at milder reaction conditions than used in the thermal processes. The present study indicates that it is advantageous to conduct the reaction in a trickle-bed reactor with partial wetting of catalyst particles, perhaps with cyclic operation, since a direct contact between the catalyst surface and gas-phase increases the concentration of active sites for phenol oxidation. Furthermore, the reaction selectivity in a trickle-bed reactor is higher than that in a slurry reactor. The main drawback of the investigated process is dissolution of metal ions into the liquid-phase, which calls for more stable catalysts.

REFERENCES

1. S. Goto, J. Levee and J.M. Smith, Catal. Rev. - Sci. Eng., 15 (1977) 187. 2. J.R. Katzer, H.H. Ficke and A. Sadana, J. Water Poll. Control Fed., 48 (1976) 920. 3. F. Luck, Catal. Today, 27 (1996) 195. 4. A. Sadana and J.R. Katzer, J. Catal., 35 (1974) 140. 5. H. Ohta, S. Goto and H. Teshima, Ind. Eng. Chem., Fundam., 19 (1980) 180. 6. A. Pintar and J. Levee, J. Catal., 135 (1992) 345. 7. S. Ho~evar and A. Radelj, NIC Intemal Report #1336, Ljubljana, 1992. 8. A. Pintar and J. Levee, Ind. Eng. Chem. Res., 33 (1994) 3070. 9. A. Pintar and J. Levee, Chem. Eng. Sci., 49 (1994) 4391. 10. A. Pintar, Ph.D. Thesis, University of Ljubljana, Ljubljana, 1996. 11. J. Levee, Appl. Catal., 63 (1990) L 1. 12. A. Pintar and J. Levee, Chem. Eng. Sci., 47 (1992) 2395. 13. H.R. Devlin and I.J. Harris, Ind. Eng. Chem., Fundam., 23 (1984) 387. 14. C.S. Turchi and D.F. Ollis, J. Catal., 122 (1990) 178. 15. J. Haber, in J.R. Anderson and M. Boudart (Eds.), Catalysis - Science and Technology, Vol. 2, Springer-Verlag, Berlin, p. 13, 1981. 16. K. Fajerwerg and H. Debellefontaine, Appl. Catal. B: Environmental, 10 (1996) L229. 17. A.I. Njiribeako, P.L. Silveston and R.R. Hudgins, Can. J. Chem. Eng., 56 (1978) 234.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

AMMONIA OXIDATION OVER AND MECHANISTIC STUDY

CuO/TiO2

643

CATALYST:

SELECTIVITY

G. Bagnasco,G. Peluso, G. Russo, M. Turco, Dipartimento di Ingegneria Chimica, Universit~ "Federico II", P.le Tecchio, Naples, Italy G. Busca, G. Ramis, Istituto di Chimica, Facoltfi di Ingegneria, Universit~ di Genova, P.le J. F. Kennedy, Genoa, Italy

A CuO/TiO2 catalysts was characterized by NH3 temperature programmed desorption (TPD) and Fourier transform infrared (FT-IR) spectroscopy and tested for NH3 oxidation. TPD measurements showed two forms of adsorbed NH3, one of which could be removed by treatment with water vapour. FT-IR spectra showed NH3 coordinated to Lewis acid sites, which gave rise, after treatment at 150~ to adsorbed hydrazine and nitrxyl species. In NH3 oxidation tests conversions up to 90% were observed. N2 was the main product, N20 and NO being formed to lower extents. The addition of water vapour in the feed influenced the product distribution. A reaction mechanism was proposed, involving adsorbed hydrazine, nitroxyl and amido species as intermediates for N2, N20 and NO production, respectively. 1. INTRODUCTION The selective catalytic oxidation of ammonia to nitrogen (SCO process) has been proposed for the abatement of ammonia from waste gases. Among different formulations, catalysts based on CuO-TiO2 offer interesting properties [1]. However, selectivity to nitrogen can be decreased by the parallel formation of N20 and NO. The behaviour of metal oxide catalysts upon these reactions can be influenced by water vapour that is always present in the reacting mixture [2 ]. In this work, a CuO/TiO2 catalyst (10wt% CuO, specific surface area = 120 m2/g) was studied in NH3 oxidation. The catalyst was also characterized by NH3 TPD technique and FTIR spectroscopy in order to obtain information about the nature of active sites for SCO and the occurrenece of competitve adsorption of water on such sites. A mechanism of ammonia oxidation is proposed. 2. E X P E R I M E N T A L

NH3 oxidation tests were carried out in a continuous flow fixed bed reactor [3], operating at atmospheric pressure and a temperature ranging from 200 to 350~ with NH3 concentration in the feed of 200, 350 or 700 ppm; a constant concentration of 02 of 2.7%, a water vapour level, when present, of 400, 700 or 1000 ppm and a contact time of 4.5x10 -3 s (total flow rate=80N1/h, catalysts weight-0.1 g).

644 FT-IR spectra were obtained with a Nicolet Magna 750 apparatus. NH3 TPD measurements were carried out in a flow apparatus with a TCD detector as described before [4]. A KOHwater trap was placed before the detector in order to avoid water interference on TCD signal. Some TPD measurements were carried out by treating the sample with a 0.6% water vapour/He mixture after NH3 adsorption. 3. RESULTS In Fig. 1 NH3 TPD spectra obtained with and without water vapour treatment are reported. The spectrum obtained without water vapour treatment shows a composite peak, with maxima at 172 and 258~ As the CuO content of this catalyst is sufficient for the complete coverage of TiO2 surface, the signal is likely to be due to surface copper oxide species [3]. The spectrum indicates at least two forms of NH3 adsorbed at weak to medium strength on copper oxide species. MS analysis of the effluent showed that NH3 is desorbed intact from the weaker sites, while it is oxidized to some extent to N2 on the stronger ones. Water vapour completely romoves the low temperature species, but does not affect the high temperature signal. This results shows that weak interacting forms of adsorbed NH3 are completely displaced by water.

r

2

x

1

Y,=

-1Z

o

I

'

200 Temperature, ~

I 400

Fig. 1. NH3 TPD spectra of the catalyst a) without water vapour treatment; b) with water vapour treatment.

645 Fig. 2,a) reports the FT-IR spectrum of the adsorbed species arising from ammonia adsorption at room temperature on the catalyst outgassed at 350 ~ It presents two strong bands, both quite broad, with maxima at 1598 and 1220 cm l respectively. The higher frequency component is typical of the asymmetric deformation of coordinated ammonia, 8asNH3, and presents a tail towards the higher frequency side. The lower frequency feature is typical for the corresponding symmetric deformation mode, 8symNH3, and presents a tail towards lower frequencies. The position of the 8symNH3 mode is characteristic of NH3 adsorption on cationic sites of quite strong Lewis acid strength [5,6]. The broadness of the two absorption peaks are quite typical of coordinatively adsorbed ammonia, while the tailing is evidence of heterogeneity of the sites. Comparison to ammonia adsorption on the TiO2 support [7] indicates that the adsorbing sites are predominantly copper ions. Additionally, a weak and very sharp band at 1440 cm -1 is also observed. This band, although it falls in the region typical for the asymmetric deformation of ammonium cations,

d)

A b

c)

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1900

1800

1700

1600

1500

1400

1300

1200

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Wavenumber (cm 1)

Fig. 2. FT-IR spectra of the adsorbed species arising from adsorption on 10 % CuO/TiO2 of a) ammonia at r.t.; b) ammonia at 150 ~ c) hydrazine at 150 ~ d) hydroxylamine at 150 ~ 8asNH4+, cannot be assigned to this mode because of its sharpness, which fully contrasts the typical broadness of the 8~NH4 + mode. This sharp band has also been observed after ammonia adsorption over other surfaces such as FezO3/TiO2, CrOx/TiO2 and MnO• [8],which are able to perform more or less selectively ammonia oxidation, and has been assigned to an oxidation or transformation product of ammonia.

646 By heating in vacuum at 150 ~ several absorption bands appear at 1611 c m "1 (strong and sharp), 1560 c m -1 ( w e a k and sharp), 1480 c m "1 (strong and sharp), 1440 c m "1 (strong and sharp), 1380 cm -1 (weak), 1346 cm -1 lweak), 1325 cm -1 (medium), 1295 cm-1 (very weak), 1210 c m (medium), and 1187 c m (medium-strong). It is evident that heating causes ammonia to undergo substantial changes. To understand what happened to ammonia, the adsorption of potential oxidation products and/or intermediates has been investigated. The adsorption of hydrazine (NHE-NH2), hydroxylamine (NH2OH), nitrogen (N2), nitrous oxide (N20), nitric oxide (NO), and nitrogen dioxide (NO2) has been investigated. The spectra of the adsorbed species arising from N2, N20, NO and NO2 have nothing in common with the spectrum reported in Fig. 2 b). On the contrary, the spectrum of the adsorbed species arising from hydrazine (NH2-NH2), shown in Fig. 2 c), is identical, after outgassing at 150 ~ as that of the species arising from ammonia. In particular, the quite strong and sharp bands at 1611 and 1560 c m -1 (symmetric and asymmetric 8NH2 "scissoring" modes), and at 1210 and 1187 cm -1 (N-N stretching and NH2 rocking, respectively) almost certainly belong to coordinated hydrazine species [9]. The strong bands at 1480 and 1440 c m 1 , instead, cannot be assigned to hydrazine, while the bands in the 1400-1250 cm -1 region are also unlikely to be due to hydrazine. The detection of the band at 1440 cm -1 alone after adsorption of ammonia at r.t. suggests, in agreement with previous data [8], that the bands at 1480 and 1440 c m "1 a r e not due to the same species. The spectra of the adsorbed species arising from adsorption of hydroxylamine NHEOH allow the likely identification of the species arising from ammonia. Hydroxylamine adsorption gives rise to a strong band centered at 1565 cm -~ and a weaker band at 1040 cm -~ that can be assigned to the NH2 scissoring and wagging modes of NH2-O- species. These bands nearly correspond to those observed for gaseous hydroxylamine and its liquid O-alkyl derivatives [10,11 ]. These bands, however, are not present in the spectrum of the species arising from the ammonia transformation. After heating at 150 ~ (Fig. 2 d)), however, the spectrum of the adsorbed species arising from adsorption of NH2OH shows a fairly strong band near 1448 c m 1, and additional components at 1415 (shoulder), 1389 and 1350 c m "1. The strong band at 1448 c m -1 certainly does not belong to an adsorbed hydroxylamine species, but could correspond to the component observed near 1440 cm ~, from both ammonia and hydrazine. Also the bands at 1390 and 1350 cm -1 could correspond, due to their relative intensity and shape, to those observed at 1380 and 1346 c m "1 after both ammonia and hydrazine reactive adsorption. To identify these species, the known chemistry of hydroxylamine can be taken into account. It is well known that hydroxylamine can be oxidized by high-valent cations (including Cu 2+) t o N2, N20 o r N O 3 [ 12]. In any case, the reaction carried out in solution is postulated to go through the radical-like NH20" and through its oxidation ~roduct, nitroxyl HNO. The IR spectrum of HNO in matrices shows two bands near 1560 cm- (NO stretching) and near 1125 cm "1 (NH deformation) [13]. However, it is fully reasonable to suppose that when it interacts with a surface cationic center the N=O bond order should decrease, while the NH deformation mode should increase, rising to a value typical for N-H deformations of secondary amines and/or imines, i.e. near 1400 cm "1. We can consequently propose an assignment, that appears to be reasonable both on chemical and spectroscopic grounds, of the pair of bands at 1448 and 1346 c m "1 and at 1415 and 1380 cm l to two different HNO nitroxyl species adsorbed on different sites or with different configurations. The band at 1480 cm -l, produced from ammonia and hydrazine but not from hydroxylamine, can be assigned to a species that does not contain oxygen atoms. According to the chemistry discussed above, it seems reasonable to assign this band to a surface imido

647 . . . . lmldo-metal complex M-NH-M, where M is a surface c a t i o n (Cu 2+ or T1.4+). Bridging complexes typically show the NH deformation mode in the range 1500-1400 cm -] [ 14]. The experiments relative to hydroxylamine adsorption allow the interchange of the previous assignment of the bands at 1480 and 1440 cm l [3], but substantially support the picture proposed previously. The adsorption of ammonia on a water covered surface has also been investigated. When the CuO/TiO2 catalyst is pretreated with water vapour, coordinatively held water is observed, characterized by a broad band at 1620 cm -1, due to the scissoring deformation mode of water (Fig. 3 b)). The subsequent adsorption of ammonia (Fig. 3 c)) gives rise again to the bands of coordinatively held ammonia, and to the sharp band at 1440 cm l, as obtained on the water-free catalyst. A broad weak absorption is also found nearly superimposed to the band at 1440 cm l and can be due to traces of ammonium ions (6asNH4+), due to the Bronsted acidity of adsorbed undissociated water. However, heating easily completely destroys this species. However, the comparison of the spectra of adsorbed ammonia with and without previous water adsorption (Fig. 4) shows that water perturbs the spectrum of coordinated ammonia, in particular, the higher frequency component of the 6symNH3mode is clearly reduced

1 A b S 0

r b

a

n

C e

2000

1900 1800

1700

lsoo

1 5 0 0 1 4 0 0 1:300 " 1 ,00 1 : , 0 0 1 0 0 0

900 " " '

Wavenumber (cm 1) Fig. 3. FT-IR spectra of a 10 % CuO/TiO2 catalyst disk (a) after activation by outgassing at 350 ~ (b) after adsorption of water 5 Torr and outgassing at r.t; (c) after successive adsorption of ammonia (5 Torr) and outgassing at r.t. due to the presence of water, while the lower frequency component is enhanced.This definitely shows that some kind of competition occurs between ammonia and water on the sites responsible for coordinative adsorption.

648 In Figs. 5-8 results of catalytic activity tests performed under anhydrous and water cofeeding conditions are reported. NH3 total conversion was increased rapidly with temperature for any feed concentration, reaching values as high as 90% at 300~ (Figs. 5,6). By increasing the feed NH3 concentration from 200 to 700 ppm, NH3 conversion was lowered, suggesting a reaction order lower than 1. Oxidation products were N2, N20 and NO, coming from the reactions: 4NH3 + 302 --~ 2N2 + 6H20 (1) 2NH3 + 202 ~ N20 + 3H20 (2) 4NH3 + 502 ~ 4NO + 6H20 (3) N2 was the main product at all conditions, the highest conversion to N2 being 60% at about 300~ The conversion to NO markedly increased with temperature, reaching a value of about 30% at 350~ and 200 ppm NH3 feed concentration (Figs. 5,6). Conversion to N20 became constant at 10-15 % at about 300~ (Figs. 7,8). These results could be explained by competition among the different NH3-O2 reactions, reaction (3) probably having the largest activation energy. Moreover it can be supposed that some N20 decomposition occurs at high

1500 r

i400'

i 300

'i200'

'i 100'

Wavenumber (cm "1) Fig. 4. FT-IR spectra of ammonia adsorbed on 10 % CuO/TiO2 on an activated sample (lower spectrum) and on a sample with preadsorbed water (upper spectrum). temperature, since it is known that CuO catalysts are active for this reaction [15]. Conversions to different products appear slightly influenced by the NH3 feed concentration: conversions to N2 and to NO decreased with increasing NH3 concentration while conversion to N20 increased with NH3 concentration. This suggests NH3 reaction orders lower than 1 for the reactions leading to N2 and to NO and higher than 1 for the reaction producing N20. The influence of water vapour on the catalytic activity was investigated by performing some tests with water vapour in the feed at concentrations of 700 or 1000 ppm (Figs. 5-8). The addition of water vapour in the feed caused a decrease in NH3 total conversion at all temperatures and NH3 feed concentrations. This effect is large at low temperatures and high H20/NH3 ratios. Moreover the product distribution was also affected by water vapour.

649

100

100 I

80

=O ,m

t-

oO

9

.A.

9

9

60-

..

,

.

9 9

. .

~i." 9

40-

.< s

.

9 9 9

O

..' .'? .''?

60-

.....

.

A- .... i-II

9

20-

.... : i : - : =

....-

. ..

40-

. .

t-

uo

2o

..

-

i

~,

a . ..... 111:" T

i

i

i

i

250 275 300 325 350

200

Temperature, *C

7O

60

60-

g 4o ~

..-

,:

"~ "

"''A

I .

(JO 2 0 - J

in

> tO

cJ

10

0-

~

...p.. - |

*< 5 0 =o 40im

30

350

Fig. 6. NH3 conversion (empty symbols) and conversion to NO (full symbols) as a function of temperature. Feeding concentrations: NH3 700 ppm, H20 0 (O), 700 ([--I), 1000 ppm (A)

70

*< 50

300

Temperature, *C

Fig. 5. NH3 conversion (empty symbols) and conversion to NO (full symbols) as a function of temperature. Feeding concentrations: NH3 200 ppm: H20 0 (O), 700 (O), 1000 ppm (A)

. "O~"",AlIFl .' ' ' O ' "

250

9 ~ 9 #~1; ~176

30-

e G

20-

9

0

'. o..

o

9 1 4 9

O9

9 . i, ,.

10-

' ' ' ~

~,..

,~'"~~ .~.-'"

~

250 275 300 325 350 Temperature, *C

Fig. 7.Conversion to N2 (full symbols) and to N20 (empty symbols) as a function of temperature. Feeding concentrations: NH3 200 ppm: H20 0 (O), 700 (1--]), 1000 ppm (z~)

200

250

300

350

Temperature, oC

Fig. 8.Conversion to N2 (full symbols) and to N20 (empty symbols) as a function of temperature. Feeding concentrations: NH3 700 ppm, H20 0 (O), 700 (I--1), 1000 ppm (zl)

650 Conversions to N2 and N20 were depressed, mainly at high H20/NH3 ratios. The influence of water vapour on NO formation is appreciable only at high H20/NH3, where a slight enhancement in NO production was observed. This resulted in a decreased selectivity to N2 and to N20. 4. DISCUSSION

The data presented here support the reaction mechanism reported below:

(4) (5) (6) (7) (8) (9) (10) (11) (12) (13) (14)

NH3(g) = NH3(ads) NH3(ads) = NH2(ads) + H + + e 2NH2(ads) = N2H4(ads) NaH4(ads) = N2 + 4 H + + 4 e NH2(ads) = NH(ads) + H + + e NH(ads) + 02.= HNO(ads) + 2 e 2HNO(ads) = N20 + H20 NH2(ads) + 02 = NH2OO(ads) = NO + H20 NH2(ads) + NO = NH2NO(ads) = N2+ H20 2 02 + 4 e = 2022H § + 02.= H20

This scheme also includes reaction (12) between NH2(ads) and NO formed by NH3 oxidation as a likely reaction because CuO based catalysts are active towards the SCR reaction: [ 3 ] 4NO + 4 NH3

+

02 ~ 4N2 + 6 H20

(15)

It was observed that the reaction order with respect to NH3 for N2 and NO formation was lower than 1: this can be explained assuming that the slowest step is the reaction of an adsorbed NH3 species, namely that step (5) is rate determining for N2 formation, while step (5) or (11) is rate determining for NO formation. Moreover the reaction order higher than 1 for oxidation to N20 can indicate that the slowest step for this product is the coupling step (10). The observed behavior upon ammonia adsorption on the water covered catalyst suggests that Bronsted acidity is not involved in the catalytic activity of CuO/TiO2 towards ammonia oxidation. The data suggest that adsorption of ammonia in the absence of water can give rise to different adsorbed species (I and II), as shown below.

HH H N

H

H

Cu

cu

I

II

H

HH,N

H

Cu III

651 The low and medium temperature signals in the TPD spectrum are due to the desorption of NH3 coordinated to species II and I, respectively, as expected from the strength of interaction of these NH3 species with Cu 2+ sites. In TPD measurements species II desorbs one NH3 molecule at low temperature, forming species I that, in turn, desorbs NH3 at higher temperature giving rise to two TPD peaks. Water certainly competes with NH3 for the formation of both species I and II, but allows the formation of new species like species III. The competition of water with ammonia causes a decrease in the FT-IR spectra of the band characteristic of species I (SsymNH3mode near 1220 cm -1 ) while it results in an increase of a band at lower frequency (near 1160 cm-1) assigned to species III. In the TPD experiments, treatment with water vapour at r.t. after NH3 adsorption, can lead to the displacement of ammonia by water, thus transforming species II into species III. When the sample is heated, since water desorbs at a temperature lower than ammonia, species III transforms into species I, while species II cannot be formed. This mechanism can explain why TPD experiments following water treatment show ammonia desorbing only at high temperature, while the low temperature signal is completely absent. It is possible to suppose that species like II are the most favourable to give N2 and N20 by ammonia oxidation, possibly through the following scheme: H H,.. N

H N~H

,,,,

N 2

~' H

Cu

H

t

L

&N "A.Cu ~ N,i

~.~H

H

I

o=N

I

~'~

v'--

~N~ Cu

N20

O

So that, water, that selectively inhibits the formation of species II, selectively inhibits the "coupling reactions" to give nitrogen and nitrous oxide. On the contrary, the more strongly bonded species I could be involved in the formation of NO, as well as in the SCR reaction between ammonia and NO through the following mechanisms:

HH , ~O-O. 9

9

N

O - O J~r

HH H N

~ ++ Cu

HH

/

V+

~

NO

H20

Cu

N"

*+

Cu

~

N-O

u

H~ N - O N ~ +

Cu

v

N2

H20

652 As previously said, water does not inhibit the formation of species I at high temperatures (where species I can be formed by decomposition of species III), and this agrees with the absence of any detectable effects of water in the formation of NO by ammonia oxidation. According to the above mechanism, it is also expected that the reaction NO + NH3 will not be affected by the presence of H20. In previous work it was found that water inhibits the SCR reaction on metal oxide based catalysts [2,16], but to a lower extent in comparison with ammonia oxidation. However, the SCR reaction occurs at lower temperatures with respect to ammonia oxidation to NO, so that at these conditions species III and species I are still in competition. 5. CONCLUSIONS The work allowed to identify the species formed from NH3 adsorption on CuO/TiO2 catalyst that are probably involved in NH3 oxidation. Adsorbed hydrazine leads to formation of N2, while N20 and NO are likely formed through nitroxyl and amido intermediates, respectively. Water vapour reduces conversions of NH3 to N2 and N20, while enhancing conversion to NO: this effect can be explained by competition of NH3 and H20 for some adsorbing sites. REFERENCES

1. A. Wollner, F. Lange, H. Schmelz, H. Knozinger, Appl. Catal. A: General, 94 (1993) 222. 2. G. Bagnasco, G. Busca, F. Giaccio, G. Ramis, G. Russo, M. Turco, 1996 Annual AIChE Meeting, Environmental Catalysis Section, Chicago, November 10-15, 1996. 3. G. Ramis, L. Yi, G. Busca, M. Turco, E. Kotur, R. J. Willey, J. Catal., 157 (1995) 523. 4. M. Turco, P. Ciambelli, G. Bagnasco, A. La Ginestra, P. Galli, C. Ferragina, J. Catal., 117 (1989) 335. 5. A.A. Tsyganenko, D.V. Pozdnyakov and V.N. Filimonov, J. Mol. Struct., 29 (1975) 299. 6. K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination Compounds, 4th ed., Willey, New York, 1986. 7. G. Busca, H. Saussey, O. Saur, J.C. Lavalley and V. Lorenzelli, Appl. Catal., 14 (1985) 245. 8. J.M. Gallardo Amores, V. Sanchez Escribano, G. Ramis and G. Busca, Appl. Catal. B: Environ., 308 (1997) 1. 9. D.N. Sathyanarayana and D. Nicholls, Spectrochim. Acta, 34A (1978) 263. 10. H. Siebert, Anwendungen der Schwingungsspektroskopie in der Anorganischen Chemie, Springer, Berlin, 1966. 11. W.O. George, J.H.S. Green and M.J. Rox, Spectrochim. Acta, 26A (1970) 2007. 12. G. Stedman, Advan. Inorg. Chem. Radiochem., 22 (1979) 113. 13. J.F. Ogilvie, Spectrochim. Acta, 23A (1967) 737. 14. S. Kagami, T. Onishi, and K. Tamaru, J. Chem. Soc. Faraday Trans. I, 80 (1984) 29. 15. G. Centi and S. Perathoner, Appl. Catal. A: General, 132 (1995) 179. 16. M. Turco, L. Lisi, R. Pirone, P. Ciambelli, Appl. Catal. B: Environ., 3 (1994) 133.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

653

Metalloporphyrin-catalysed oxidation of azonaphthol dyes: the mechanism of oxidative bleaching by oxoiron(IV) porphyrins in aqueous solution George Hodges', John R. Lindsay Smith *aand John Oakes b "Department of Chemistry, University of York, York YO1 5DD, United Kingdom bUnilever Research, Port Sunlight Laboratories, Quarry Road East, Bebington, Wirral, Merseyside L63 3JW, United Kingdom 1. INTRODUCTION

Peroxidases are ubiquitous, haem containing enzymes that are present in plants, in some animal tissues and in micro-organisms. They catalyse a variety of biosynthetic and degradative oxidations of phenols and amines by hydrogen peroxide or hydroperoxides. Horseradish peroxidase (HRP) and other peroxidases, particularly those present in white rot fungus Phanerochaete chrysosporium, have also been studied as potential catalysts for the oxidative degradation of aqueous solutions of azo dyes and phenolic environmental pollutants. 19 The generally accepted mechanism for phenolic oxidations catalysed by horseradish peroxidase (HRP), the most thoroughly studied peroxidase, involves an initial two-electron oxidation of the iron(III) resting state of the enzyme by the hydroperoxide (or hydrogen peroxide) to give compound I (HRP I) which is subsequently reduced back to the iron(III) haem, in two one-electron steps, via compound II (HRP II) (Scheme 1). The phenoxyl radical intermediates react to give products, which depending on their structure, may themselves be further oxidised.

Resting state

-

RO2H

S.

0 !1 > ~-'f--F e%

+"

~ f - ~ S-H

Compound II S-H is an oxidisable substrate and S" is a substrate radical Scheme 1

ComDoundI

654 The reactions of peroxidases have been mimicked, using synthetic iron(III) and manganese(III) porphyrins catalysts as models for the haem prosthetic group, and these simplified systems have been used to obtain mechanistic information about the oxidations) ~ At York we have been investigating the oxidation of phenols by oxoiron(IV) and oxomanganese(IV) porphyrin models for HRP II in aqueous solution. ~5'~6This research has shown that, at pH 7.6, oxidation proceeds by hydrogen atom abstraction by the oxometallo(IV) species from the phenol rather than electron transfer from the phenolate anion. Preliminary studies with commercial azo dyes has revealed that these are also substrates for oxoiron(IV)porphyrins. 17 In this paper we report the results from our investigations of the mechanism of oxidation of 1-phenylazo-2-naphthol-6-sulfonates by oxoiron(IV) tetra(2,6dichloro-3-sulfonatophenyl)porphyrin (1) in aqueous solution.

2. RESULTS AND DISCUSSION 2.1. Selection of azo dye substrates

The azo dyes used in this study were 1-phenylazo-2-naphthol-6-sulfonate (2) and seven derivatives with substituents in the meta or para positions of the phenyl ring (3-9). These were selected as representative 1-azo-2-naphthol sulfonate dyes and because the substituents on phenyl ring would allow a systematic study of the mechanism of their oxidation by metalloporphyrin-catalysed systems in aqueous solution. Dye 2 is commercially available (as Acid Orange 12) and was purified by recrystallisation whilst the others (3-9) were prepared by standard diazonium ion/2-naphthol coupling reactions. The purities of all the dyes were checked by TLC, MS and ~H NMR spectroscopy. Table 1 reports the measured pK~ values of all the azo dyes used in this study.

SO 3"

c1 , N" -O3

~

X % C1 .

H

NaO3S

(1)

X = H (2), 4-NO 2 (3), 3-CF3 (4), 4-COMe (5), 4-C1 (6), 4-CHMe 2 (7), 4-Me (8), 4-OMe (9)

655 Table 1 pKa values of substituted 1-phenylazo-2-naphthol-6-sulfonate dyes and second order rate constants for their reaction with oxoiron(IV) tetra (2,6-dichloro-3-sulfonatophenyl)porphyrin in aqueous solution: pH, 6.93; let, 0.05 mol dm 3 Dye (2-9)

Substituent constant b pK~~

X ,

k2(dm3molqs"1)

o

~-

,

o"

,

4-NO2

10.62

69

0.81

0.79

4-MeCO

10.81

103

0.47

0.36

0.060

3-CF3

10.31

116

0.53

0.52

-0.014

4-C1

10.36

492

0.24

0.11

0.011

H

10.66

522

0

0

0

4-Me2CH

10.81

2770

-0.15

-0.28

0.009

4-Me

10.71

2884

-0.14

-0.31

0.015

4-OMe

10.69

18938

-0.28

-0.78

0.018

(a) Measured in deionised water (40 ~ C) and corrected for ~t = 0.05 mol dlTl'3; TM (b) o values, except o4Mo2~, from ref. 23; 04Me2Cnand o + values from ref. 24; o" values from ref. 25.

2.2. Generation of oxoiron(IV) porphyrin In a previous study, 19 w e have shown that both anionic and cationic iron(III) porphyrins, on reaction with oxidants such as t-butyl hydroperoxide or peroxyacids, form relatively stable oxoiron(IV) species in aqueous solution. The porphyrin chosen for this investigation was the sterically hindered, anionic iron(III) tetra(2,6-dichloro-3-sulfonatophenyl) porphyrin (FemTDCSPP), since this is unable to form potentially unreactive la-oxo-dimers2~ and its oxoiron(IV) derivative (1) is the most stable of the anionic iron porphyrins studied previouslyJ 9 Furthermore, the analogous unsulfonated compound is an excellent oxidation catalyst in organic solution.2~, 22 Addition of a 2.5-fold molar excess of 3-chloroperoxybenzoic acid (3CBPA) to an aqueous solution (pH 6.93) of FemTDCSPP leads to a rapid bathochromic shit~ (~-m~x412 to 420 rim) of the Soret band in the UV-Vis spectrum of the porphyrin, arising from the complete conversion to the oxoiron(IV) species (1) (Figure 1). For the kinetic studies (30~ C, pH 6.93 and ionic strength 0.05 mol dm'3), however, it was important that all the peroxyacid had been consumed before the substrate was added. As in our earlier investigations, we assumed a 2"1 stoichiometry of Fer"TDCSPP to 3CPBA (Scheme 2) and employed a half-molar equivalent of the peroxyacid. Under these conditions there was a 70% conversion of FemTDCSPP to 1. Since the reaction between FemTDCSPP and 3CPBA is fast, the remaining 30% of iron(HI) porphyrin shows that

656 all the oxidant was consumed, presumably in competitive side reactions of the oxoiron(IV) porphyrin n-cation radical and 3CPBA. In the absence of added substrate, there is a quantitative conversion of I back to FeUJTDCSPP in a relatively slow first order reaction (tin 4.5 min).

e:

0.08~

s,. [

t'

? i

t

i

l

0.04-

;, t,

/ i

'

i'."

,

i'

a

/

t

..'

i J

../ .-9 ,r

4

--~--'..7 "~,...~....~.=.____ ~.

-~ ~ --

~----'-'----~'':.. ,.

.;o

860 k/nm

Figure 1. UV-Vis specman of an aqueous solution (pH 6.93) of lx10 .6 mol dm"3FemTDCSPP, (-) before and (--) after the addition of 2.5xl 0 .6 mol dm3 3CPBA

R~ OH

RCO~

O

-

~

OH

~

~

2%..~

Scheme 2 2.3. Characterisation of the kinetic equation for the oxidation of Acid Orange 12 by oxoiron(IV) tetra(2,6-dichloro-3-sulfonatophenyl)porphyrin The reactions were carried out in a thermostatted cuvette (30 ~ C) and were followed by monitoring the decay of the Soret band of the oxoiron(IV) porphyrin (Xma~420 rim) with a Hewlett Packard 8452 diode array spectrometer. The standard method involved preparing the oxoiron(IV) species by the addition of a half molar equivalent of 3CPBA to lxl0 6 mol dm3 Fe"~TDCSPP in 0.025 mol dm3 aqueous phosphate buffer (pH 6.93) and ionic strength 0.05 mol dm"3(maintained with NaNO2). The reaction was initiated by the addition of a solution of the dye (f'mal concentration 5xl 0 .6 to 5xl 0 .5 mol dm3) in the same aqueous buffer. The disappearance of the oxoiron(IV) porphyrin followed first order kinetics for more than 3 half-lives (Figure 2) and the measured kobsvalues were found to be directly proportional to the concentration of the dye (Figure 3). The slope of the plot is the k2, the second order rate constant

657 for the oxidation of Acid Orange 12 by the oxoiron(IV) species, and the intercept, kb, is the dye independent, first order rate constant for the background reduction of 1. The overall rate law for decay of I follows the kinetic equation 1. -d[l]/dt = kobs[1] where kohs= kb + k2[1][Dye] 1 -14.6 -14.8 9 -15.0 13_ > -15.2 (D LL. -15.4 0 -15.6

t"

J 9 -15.8 -16.0 -16.2 20

10

30

40

Time / sec Figure 2. First order plot for reduction of OFervTDCSPP (from lxl 0 -6 mol dm "3FemTDCSPP) by 5x10 5 mol dm 3 Acid Orange 12, pH 6.93; la = 0.05 tool dm 3 ; 30 ~ C 3.0 2.5 ,~ ,',2,

2.0 1.5

J

A 1.0 .5

0.0

X

J

0.0

1.0

2.0

3.0

4.0

5.0

[Acid Orange 12] / 10.5 mol dm 3 Figure 3. Dependence of pseudo-first order rate constant for reduction of OFervTDCSPP by Acid Orange 12 on [Dye], pH 6.93; ~t = 0.05 mol dm'3; 30 ~ C 2.4. The effect of substituents, on the phenyl group of Acid Orange 12, on the second order rate constants of dye oxidation Linear kobsvs dye concentration plots, equivalent to Figure 2, and hence k 2 values were also obtained for the seven derivatives of Acid Orange 12 with substituents on the m e t a or p a r a position of the phenyl ring (Table 1). A Hammett analysis of the rate data reveals a good correlation of the log k2 values against ~ with a p value of-1.66 (R = 0,981) (Figure 4), whereas

658 the corresponding plot using o gives a poorer correlation (p = -2.06 with R = 0.936). Unlike previous studies on the oxidation of phenols with cationic oxoiron(IV) and oxomanganese(IV) porphyfins, ~5'16with the dye oxidations there is no significant correlation with o', suggesting that any radical character in the latter transition state is small and dominated by the polar substituent effect.

2.0 1"5t 1.0"

,~ 0.5, 0.0. -0.5.

4-COMe

-1.0. -1.5

'

0

I

-05

'

'

0'0

,

0'5

10

o+

Figure 4. Correlation of log k 2 v s . G+ for the oxidation of 1-phenylazo-2-naphthol-6-sulfonate dyes by OFervTDCSPP 2.5. Kinetic isotope effect studies on the oxidation of l-(4-methylphenylaz~)-2-naphthol-6sulfonate in deuterated buffer The rate of oxidation of 1-(4-methylphenylazo)-2-naphthol-6-sulfonate by 1 was followed as described above using D20 in place of water.When the trace amounts of HzO, arising from the buffer salts and other proton sources, had been taken into account the substrate was calculated to be >99% in the deuterated form. Under these conditions the pD value of the reaction mixture (obtained by adding 0.38 to the value obtained with the pH meter)26 was 7.50. This increase from the value of 6.93 obtained in 1-I20 arises from an increase in pKa of the buffeting salts in DzO) 8 The pI~ of the dye (10.71 in water at ionic strength 0.05 mol dm3) also increases (calculated value 11.14 in I)20) 27 this compensates for the change from pH 6.93 to pD 7.50 and as a result there is a negligible effect on the degee of ionisation of the dye. The reaction between I and the deuterated dye was shown to be first order in both the dye and the oxoiron(IV) species with a second order rate constant, after correcting for the selfreaction of 1, of 1829 + 49 dm 3 mol "~s~. This leads to a kinetic isotope effect, kn/kD, of 1.58. This value is comparable to but larger than that observed for H atom abstraction from 4fluorophenol (kH/k, 1.32) by oxoiron(IV) tetra(2-N-methylpyridyl)porphyrin at pH 7.7.15

659

2.6. The stoichiometry of the oxidation of Acid Orange 12 The stoichiometry of the oxidations was examined with the most reactive of the dyes used in the present study, 1-(4-methoxyphenylazo)-2-naphthol-6-sulfonate, to minimise problems from the slow self-reaction of 1. This revealed that the reaction requires four oxoiron(IV) species to oxidise each molecule of dye. The consequence of this result is that the true k2 value for the reduction of I by an azonaphthol dye in this study is a quarter of the value shown in Table 1. This correction does not, however, affect the Hammett correlations and p values. 2.7. The mechanism of azonaphthol dye oxidation by oxoiron(lV) porphyrins Our recent studies on phenol oxidation by oxoiron(IV) and oxomanganese(IV) porphyrins in aqueous solution (pH 7.6) have shown that the initial step in these reactions is H atom abstraction to generate a phenoxyl radical (Scheme 3).ts'~6The kinetic isotope effect measured in the present study indicates that H atom abstraction also occurs in the rate determining step of the oxidation of azonaphthol dyes by 1. However, although the structures of azonaphthol dyes are normally shown as azo compounds, in aqueous solution they are in a rapid dynamic equilibrium with their hydrazone tautomers; the latter isomer being the dominant species (Scheme 4)fl This complicates kinetic studies on the dyes, since the substrate is in effect a mixture of two compounds. Consequently one or both the tautomers may be the active form of the substrate providing the H atom for the oxidant. It is important to note that, irrespective of which tautomer is the reactive substrate, H atom abstraction leads to a common azonaphthoxyl radical intermediate and subsequent reactions of this species should be independent of the initial tautomerism (Scheme 5).

ArOH

+

O~:e !

-~ArO"

+

HO

Scheme 3

xxN

Azo-tautomer.

------

SO3Na

N,, N

Hydrazone-tautomer.SO3Na Scheme 4

The position of the azo-hydrazone equilibrium has been studied using a range of spectroscopic methods which show that it is sensitive to solvent polarity, solvent H-bonding and

660 to substituents on the phenyl ring: electron-withdrawing groups favour the hydrazone and elcctron-donors the azo tautomer. 29 A further difference between simple phenols and the azonaphthol dyes used in this study is the strong internal H-bonding that occurs in both of the dye tautomers. This results in a large downfield shitt of the 1H NMR signal associated with this proton and an increase in pK a(Table 1). Examination of the measured second order rate constants for the dye oxidations shows that the reaction is significantly favoured by electron-donating substituents: the 4-methoxylated compound (9) reacts 274 times faster than the 4-nitro analogue (3). Furthermore, the excellent linear Hammer correlation of log(second order rate constant) with o ~ and the p value of-1.66 conf'trm this conclusion and show that all the substrates are oxidised by the same mechanism. However, the size of the p value is larger than that obtained from the earlier oxidation of phenols (p-1.10) 15and is larger than expected for a mechanism involving H atom abstraction in the rate determining step. A possible explanation is that the substituents have mo effects on the reaction which reinforce each other: they influence H atom abstraction by the electrophilic oxidant, as noted previously this is favoured by electron-donation,~5'16and they shift the position of the azohydmzone equilibrium. Berrie et al. 3~ carded out an NMR study on the tautomeric equilibria of an analogous series of unsulfonated dyes in chloroform and measured the percentage tautomer distribution for each dye. Their data show that the azo tautomer is favoured by electron donation and vice v e r s a , furthermore, the tautomer distributions can be used to calculate equilibrium constants and to examine the influence of substituents on the equilibrium. The data from eight dyes (omitting the 4-cyano compound) give a Hammett p value of-0.46 (vs 6+, R = 0.930). In agreement with this study in chloroform, UV-Vis spectroscopy shows that in aqueous solution (pH 6.93) the hydrazone is the major tautomer for both 3 and 9 and that the former, with an electronwithdrawing 4-nitro substituent, is almost entirely in this form (strong absorbance at 486 nm) whereas the latter, with an electron-donating 4-methoxy group, is a more equal mixture of tautomers (absorbances at 410 nm and 502 nm). It follows that if the azo tautomer is the reactive compound, then electron-donation by the substituent will increase the measured rate of reaction by increasing the proportion of this species in the reaction mixture. The p value for H atom abstraction (PH.~) is then given by equation 2, where Pobsis the measured p value and Ptaut is the p value for the azo dye tautomerism. PH.abs = Dobs- Ptaut

2

From the p value obtained in this study and from the value calculated from the work of Berrie et al., PH.absis -1.20. Although this p value from equation 2 is close to that obtained from our earlier study of phenol oxidation by oxoiron(IV)tetra(2-N-methylpyridyl)porphyrin (p-1.10), it is important to note that P~ut was obtained from a study in chloroform whereas the present investigations used aqueous buffer (pH 6.93). Spadaro and Renganathan, s from their studies on the peroxidase-catalysed oxidation of azo dyes with hydrogen peroxide, have proposed a mechanism involving the initial formation of the azonaphthoxyl radical. Subsequent one-electron oxidation and reaction with water result in the cleavage of the azophenyl group from the 1-position of the naphthol ring and loss of dye colour. Based on this mechanism and the 4:1 stoichiometry of the oxidation, we suggest the mechanism

661 in Scheme 5 to account for the oxidation of azonaphthol dyes by oxoiron(IV) porphyrins.

"0

oFivp

0

2>N= N X

X

SO 3-

SO 3-

SO 3OFeWP H20

~

O ~---N:NH

X

+ O

+ O

12OFeWp

~

H20

SO 3-

SO 3-

X

Scheme 5 ACKNOWLEDGEMENTS We thank Unilever Research and the University of York for fmancial assistance and the EPSRC Mass Spectrometry Service, Unversity of Wales, Swansea for measuring the electrospray mass spectrum of iron(Ill) tetra(2,6-dichloro- 3-sulfonatophenyl)porphyrin. REFERENCES

1. 2. 3. 4. 5.

K.E. Hammel and P. J. Tardone, Biochemistry, 27 (1988) 6563. K. Valli and M. H. Gold, J. Bacteriol., 173 (I 991) 345. A. Paszczynski and R. L. Crawford, Biochem. Biophys. Res. Commun., 2178 (1991) 1056. J. Spadaro, M. H. Gold and Renganathan, Appl. Environ. Microbiol. 58 (1992) 2397. A. Paszczynski, M. B. Pasti- Grigsby, S. Goszczynski, R. L. Crawford and D. L. Crawford, Appl. Environ. Microbiol. 58 (1992) 3598. 6. M.B. Pasti-Grigsby, A. Paszczynski, S. Gozczynski, D. L. Crawford and R. L. Crawford, Appl. Environ. MicrobioL.58 (1992) 3605.

662 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18 19. 20. 21. 22. 23. 24, 25. 26. 27. 28. 29. 30.

C. Muralikrishna and V. Renganthan, Biochem. Biophys. Res. Commun. 197 (1993) 798. J.T. Spadaro and V. Renganathan, Arch. Bioehem. Biophys. 312 (1994) 301. M. Stibrova, B. Asfaw, E. Frei and H. H. Schmeiser, Collect. Czech. Chem. Commun. 61 (1996) 962. R. Labeque and L. J. Marnett, J. Am. Chem. Soc., 111 (1989) 6621. T.G. Traylor J. P. Ciccone, J. Am. Chem. Soc., 111 (1989) 8413. J.R. Lindsay Smith and R. J. Lower J. Chem. Soc. Perkin Trans. 2, (1991) 31. T.C. Bruice, Acc. Chem. Res. 24 (1991) 243. T.G.Traylor, S. Tsuchiya, Y.-S. Byun and C. Kim, J. Am. Chem. Soc. 115 (1993) 2775. N. Colclough and J. R. Lindsay Smith, J. Chem. Soc. Perkin Trans. 2, (1994) 1139. N.W.J. Kamp and J. R. Lindsay Smith, J. Mol. Catal. A: Chem., 113 (1996) 131. G. Hodges and J. R. Lindsay Smith, unpublished observations. R.A. Robinson and R. H. Stokes, in Electolyte Solutions, Butterworths, London, 1965. S.J. Bell, P. R. Cooke, P. Inchley, D. R. Leanord, J. R. Lindsay Smith and A. Robbins, J. Chem. Soc. Perkin Trans. 2, (1991) 549. T.C. Bruice, Acc. Chem. Res., 24 (1991) 243. P.S. Traylor, D. Dolphin and T. G. Traylor, J. Chem. Soc. Chem. Commun., (1984) 279. B. Meunier, Chem. Rev., 92 (1992) 1411. N.B. Chapman and J. Shorter in Correlation Analysis in Chemistry: Recent Advances, ed. O. Exner, Plenum Press,New York, 1978, p.439. Y. Okamoto and H. C. Brown, J. Am. Chem. Soc., 80 (1958) 4979. D.R. Arnold in Sustituent Effcts in Radical Chemistry, eds. H. G. Viehe, Z. Janonsek and R.Merenyi, Reidel, Dordrecht, 1986, p 171. T.H. Fife and T. C. Bruice, J. Phys. Chem., 65 (1961) 1079. N. Isaacs, Phycal Organic Chemistry, Longman, Belfast, 1987. P. Ball and C. H. Nicholls, Dyes Pigments, 3 (1982) 5. Y. Onari, Bull Chem. Soc. Jpn. 58 (1985) 2526. A. H. Berrie, P. Hampson, S. W. Longworth and A. Mathias, J. Chem. Soc. Section B, (1968) 1308.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

663

VOC's ABATEMENT: PHOTOCATALYTIC OXIDATION OF T O L U E N E I N V A P O U R P H A S E O N A N A T A S E TiO~ C A T A L Y S T V. Augugliaro a, S. Coluccia b, V. Loddo ~, L. Marchese b, G. Martra b, L. Palmisano ~, M. Pantaleone ~ and M. Schiavello a aDipartimento di Ingegneria Chimica dei Processi e dei Materiali, Universith degli Studi di Palermo, Viale delle Scienze, 90128 Palermo, Italy bDipartimento di Chimica I.F.M., Universit/l degli Studi di Torino, Via P. Giuria 7, 10125 Torino, Italy

Photocatalytic oxidation of toluene has been carried out in a gas-solid regime by using polycrystalline anatase TiO2. A fixed bed continuous photoreactor of cylindrical shape was used for performing the photoreactivity runs; the Pyrex glass photoreactor was irradiated by a medium pressure Hg lamp. Air containing toluene and water vapours in various molar ratios was fed to the photoreactor. Toluene was mainly photooxldised to benzaldehyde although benzene, benzyl alcohol and traces of benzoic acid and phenol were also detected. The presence of oxygen was essential for the occurrence of the photoreaction while water played an important role in the mamtainance of the catalytic activity. The results obtained in a preliminary Fourier transform infrared (FT-IR) investigation indicate that toluene is weakly stabilised on the TiO2 particles by hydrogenbonding between the aromatic ring and surface hydroxyl groups.

1. I N T R O D U C T I O N Volatile organic compounds, VOC's, are an important class of mr pollutants usually found in the atmosphere of all urban and industrial areas. Toluene is one of these compounds and, due to its noxious nature, several strategies have been identified in order to reduce its presence in indoor and industrial emissions. Among the methods effective to oxidize toluene, heterogeneous photocatalysis is one of the most attractive, due to the mild conditions under which this process is usually carried out. Photocatalysis has been largely used for the photooxidation of many organic molecules in the hquid-solid regime [1-5], but a few papers report photoreactions in the gas-solid regime [6-10]. The photocatalytic oxidation of toluene in the presence of water was performed m the gas-solid regime by Ibusuki and Takeuchi [6] at room temperature by using TiO2. They found that the presence of water was beneficial in order to achieve the almost complete photo-omdation of toluene to CO2 and HzO, in fact,

664 only very small amounts of benzaldehyde, which is the main product of toluene partial oxidation, were detected. The present paper reports the results of the toluene photooxidation reaction using polycrystalline anatase TiO2 as catalyst. The photoreactivity runs were carried out in a continuous photoreactor fed by a mixture of air, toluene, and water in various molar ratios and irradiated in the near-LW region. The influence of toluene concentration, gas flow rate, and water presence on the photoprocess performance was investigated. A preliminary investigation of the interaction between toluene and the catalyst surface was carried out by Fourier transform infrared (FT-IR) spectroscopy.

2. E X P E R I M E N T A L The set-up of the experimental apparatus is reported in Figure 1. The reactivity experiments were carried out in a flow apparatus using a Pyrex cylindrical reactor whose dimensions were: internal diameter, 1 cm; external diameter, 1.2 cm; and height, 30 cm. A porous frit at the bottom of the cylinder was used to support the fixed bed of powder and to distribute the inlet gaseous mixture. The reactor was vertically positioned inside a thermostatted chamber and was irradiated through a circular window made on a wall of the chamber and covered by a Pyrex glass sheet. An aluminum parabolic reflector was located behind the photoreactor in order to mcrease the illumination. The radiation source was a 400 W medium pressure Hg lamp (Polymer GN ZS, Helios Italquartz) put at approximate 30 cm from the reactor. The radiation power impinging on the photoreactor was measured by usmg a radiometer UVX Digital and its mean value was 5 mWcm 2. The catalyst was polycrystalline TiO2 (Merck, anatase, BET surface area: 10 m2gl); the powder was classified by sieving and the fraction with particle size in the 45-90 ~m range was used. The amount of catalyst was 8 g, corresponding to a fixed bed height of ca. 10 cm. The reactant mixture was generated by bubbling air through saturators containing water and toluene at room temperature. For some runs benzoic acid was used instead of toluene. The gas flows were then mixed and fed to the photoreactor. The temperature of the reactor was always 413 K. The gas flow rates were m the 0.17-10 cm3s-1 range and the toluene molar fraction ranged from 4.0x10 4 to 1.3x10 2. The water molar fraction was held constant to 2.5x10 2. The catalyst was irradiated only when steady state conditions were achieved in the system, i.e. after about 24 h from the beginning of the photoreactor feeding. The runs lasted 170, 350 or 470 h. The gas leaving the photoreactor was periodically analyzed by a gas chromatograph (Varian, Vista 6000), equipped with FID detector. A 0.1% AT-1000 on Carbograph column (2 m x 2 ram) and a Porapak QS column (2 m x 2 ram) were used. For some experiments the gas exiting from the photoreactor was continuously bubbled in liquid acetomtrile. The resulting solution was analyzed by high pressure liquid chromatography (HPLC) (Varian 9050) in order to detect

665 compounds produced in small quantities. At the end of each run, the catalyst was held for 24 h in twice distilled water or acetonitrile in order to dissolve products adsorbed on the surface; the resulting solutions were analyzed by HPLC.

(b)

(c) ~1

(a)

_i

(h)

ill

(e)

vent

(i) GC ._1

(mt

(1)

I~176I

(o) (n)

Figure 1. Photoreactivity set-up. (a) Air cylinder, (b) control valves, (c) bottle with water, (d) bottle with toluene, (e) switch valves, (f) thermostatted chamber, (g) parabolic reflector, (h) cylindrical photoreactor, (i) lamp, (1) power supply, (m) water filter, (n) gas chromatograph, (o) bubbling bottle containing acetonitrile. Thick line was electrically warmed in order to avoid product condensation.

For selected runs the gas at the outlet of the photoreactor was bubbled in a saturated aqueous solution of barium hydroxide in order to trap CO2 as BaCO3. The IR spectra were obtained with a Bruker IFS 48 spectrometer. The catalyst powders, as self-supporting pellets, were placed m an infrared cell allowing adsorption-desorption experiments to be carried out in situ. Prior to the adsorption of toluene, the cell was evacuated (1.0xl0 -~ Torr) at room temperature.

3. R E S U L T S AND D I S C U S S I O N Blank reactivity tests were performed at the same experimental conditions used for the photoreactivity experiments but in the absence of catalyst, oxygen or light. Other runs were carried out by using COz or Nz instead of 02. No reactivity

666 was observed in all these cases so that it may be concluded that 02, catalyst, and irradiation areneeded for the occurrence of the photoprocess. The photoreactivity results showed that the reactor reaches steady state conditions after a long period of time (ca. 70 h) from the beginning of the irradiation. At steady state conditions the mare photooxidation product was benzaldehyde but also benzyl alcohol and traces of benzoic acid and phenol were detected at all the experimental conditions used. During the transient period, benzene together with CO2 were also produced. No significant evidence of CO2 production was observed at steady state conditions. These results obviously indicate that at the experimental conditions used the photoprocess does not give rise to a complete degradation of toluene. In Figure 2 the experimental data of benzaldehyde steady state production rate are reported as a function of the gas flow rate. An increase in the reaction rate occurred above a flow rate of 1.7 cm3s1, and remained constant at higher values. This indicates that external mass transfer limitations occur for flow rates less than 1.7 cmas ' . Because of this all the runs carried out for investigating the influence of toluene concentration were performed at a flow rate of 2.5 cm3s -1.

CD

A v

6

4

~~

3

.~~

2

"~

1

o 0

5

10

F l o w r a t e [cm~s -1] Figure 2. Benzaldehyde steady state production rate versus the gas flow rate. Toluene molar fraction: 4.0x10-4. Catalyst amount: 8 g. Photon flux: 5 mWcm 2.

The inlet toluene concentration greatly affected the benzaldehyde production rate. The runs performed at higher molar fraction of toluene (4.0x10 4, 7,0x10 4, and 1.3x10 -2) showed an increasing benzaldehyde production rate (4.1x10 a, 6.2x10 "a, and 3.4"x10 2 ~molsl).

667 The occurrence of catalyst deactivation and of the role played by water, was investigated by performing very long reactivity runs at the flow rate of 0.42 cm3s 1. Deactivation is not affected by the presence of mass transfer limitations, while the use of a low flow rate allows variations of outlet gas composition to be detected more accurately. Figure 3 reports the fractional conversion of toluene to benzaldehyde versus irradiation time for two long runs carried out in the presence and in the absence of water vapour. For the run in the presence of water, a maximum conversion of 0.19 (corresponding to an

O O

0 r cD

0.1

A

0

AA A m

(D ,.Q

A O

m

A

0.01 0.1

1

10

100

1000

T i m e [h] Figure 3. Toluene fractional conversion to benzaldehyde versus irradiation time for runs carried out in the presence of water vapour (m) and in the absence of water vapour (A). Flow rate: 0.42 cmasl; toluene molar fraction" 1.3x10 -~. Catalyst amount: 8 g. Photon flux: 5mWcm -2.

oxidation rate of 5.6x10 -2 ~mols -1) was achieved after 2 h, while a steady state conversion of 0.08 was achieved after 70 h of irradiation (corresponding to an oxidation rate of 2.4x10 -2 ~mols-i). No decrease of the photoreactivity was observed after 350 h. A different behaviour can be observed for the run carried out in the absence of water. Indeed, a maximum conversion of 0.1 (corresponding to an oxidation rate of 2.9x10 -2 ~mols -1) was reached after ca. 12 h. After that time, the photo-reactivity continuously decreased down to negligible values. In Figure 4 the fractional conversion of toluene to benzene is reported versus irradiation time for the same runs reported in Fig. 3. The presence of water was beneficial for benzene production, but benzene virtually disappeared after 3-4 h of irradiation, independent of the presence of water.

668 In order to determine if the catalyst deactivation was irreversible, a run was performed where first the wet reagent mixture was fed to the irradiated photoreactor and time was allowed for the achievement of a constant photoreactivity level. Then the dry reagent mixture was fed for a prolonged time; and finally the wet reagent mixture was again fed. In Figure 5 the results obtained in this run are reported as toluene fractional conversion to benzaldehyde versus irradiation time. In the absence of water a sharp decrease of toluene conversion occur from 0.08 to 0.04 after ca. 6 h and thereafter from 0.04 to 0.01 after ca. 180 h. When water vapour was again added to the reaction mixture, the photoreactivity initially increased but then slowly decreased until a

0.005 o o o0

-mmq

mm

o c~

mm

c~

A o

mm A i

0

1

m

m A

m I

2

nI

|

ILl 3

n I

n I

4

Time [h] Figure 4. Toluene fractional conversion to benzene versus irradiation time for the runs reported in Figure 3.

constant value of conversion was achieved. For this run benzene was produced only by the fresh catalyst in the first hours of irradiation; the partially spent catalyst did not produce benzene when water vapour was again present. The results reported in Figure 5 indicate that water is an essential reagent for the photoprocess. The highest activity is shown by the fresh catalyst working m the presence of water vapour; m the absence of water the catalyst progressively deactivates. The catalyst deactivation seems to be a partially reversible process; the restoration of water m the reaction mixture allows a partial recovery of the activity whose final level, however, is less than that of the fresh catalyst.

669 0.1 O O r~

.r,-I

',

;>

! ! ! ! ! ! ! !

-

O r

;= ~9

O r

- H~O

+ H20

t~ (D

(D ! !

~9 !

.l

!

!.

~ . !

.!

~.

l . ~

.~

.~

l.

! . ~

.!

~.

l . ~

.J

~.

l

~

~

O

150

250

350

450

Time [hi Figure 5. Effect of water on the photoreactivity.

Figure 6 reports FT-IR spectra obtained at various experimental conditions. The admission of toluene onto the catalyst causes the depletion of the band originally present in the background spectrum at 3670 cm -1 (curve a), due to a stretching mode of surface hydroxyl groups [11]. This band is transformed to a complex and much broader absorption in the 3600-3450 cm -I range (curve b). On the basis of such behaviour it can be concluded that surface OH groups are revolved in the adsorp.tion of toluene onto the catalyst, the resulting organic molecules are probably stabilized on the surface by hydrogen bonding between the aromatic ring and the hydroxyl groups [12]. IR bands assignable to adsorbed toluene appear in spectrum b, namely in the 3200-2800 cm 1 (CH stretchings) and 2000-1300 cm 1 (summation bands, ring stretchings). By simply outgassing at room temperature these bands disappear (spectra not reported) and the original spectral band of the surface hydroxyls is progressively fully restored (Figure 6, reset). Such reversibility suggests a weak character of the observed interaction between toluene and the catalyst. By taking into account all the results above reported, the foUowmg reaction mechanism for the production of benzaldehyde can be suggested. Under photoexcitation of the semiconductor oxide with band gap irradiation, electron-hole (e-h) pairs are photogenerated: TiO2 + hv -~ TiO2 + e- + h §

(1)

The pairs, once separated, can induce chemical redox transformations with the species adsorbed on the surface, subject to thermodynamic constraints. It is generally assumed that surface hydroxyls act as hole traps by producing OH" radicals:

670

2.0 1.3 1.2

!.5 "~ e~

1.0 0.9

l.O

0.8

r,~

0.7

<

3800 3700 3600 3500 3400 3300

0.5

Wavenumber [ cm -1 ]

0.0 3500

3000

2500

2000

1500

Wavenumber [ cm -1] Figure 6, IR spectra of toluene adsorbed on TiO2 catalyst: a) spectrum of the catalyst outgassed 30' at room temperature; b) after admission of 3 Torr of toluene vapour. Inset: a,b) the same as in the mare layer; c-h) desorption of toluene by outgassmg at room temperature for increasing times.

OH

(2)

+ h § -~ OH'.

The reactivity results indicate that oxygen is needed to sustain t h e toluene photooxidation; therefore, while OH groups act as traps for the photo-holes, adsorbed oxygen species act as traps for free photo-electrons and give rise to very reactive species by the following reactions [ 13]:

O2(gas phase) ~

O2(ads)

(3)

(ads)

(4)

02 (ads)+ H § -~ 02H"

(5)

202H"

(6)

O2(ads) "~" e -

---> 0 2

--~ 02

+

H202

H202(.~) +O2-(.~) --> OH- +

OH" + 02.

(7)

671 The radical species are very reactive and may attack toluene molecules according to the following reactions: OH"

+ C6H~CH3(.~)

-+ H20 + C6H5CH2"

C6HsCH2" + 02(,d~) --> C6HsCH200" C6HsCH2OO"

+ e

~

C~H~CHO + O H .

(8) (9) (10)

The formation of benzyl alcohol may occur by the following reaction: C6HsCH2" + OH"

~

C6H~CH2OH.

(11)

The small amount of benzyl alcohol found as a product is understandable since two radicals are involved in reaction (11). Concerning the simultaneous appearance of benzene and C02, the experimental results suggest that the breakage of the bond between the CH3 group and the ct carbon of toluene does not occur. In this last case, in fact, CH4 and/or CH30H molecules should be produced by the reaction between CH3 radicals and OH groups present on the catalyst surface but neither CH4 or CH3OH were ever detected in our experiments. On this basis the occurrence of the following reaction steps may be suggested: C6HsCHO(.~) + OH" C6H~CO" + 0 2 ( ~ ) ~ C6H5C000"

-~ C6H~CO" + H20 C6H5C000"

+ C6HsCHO(a~)-~

C6H5C0" + C6HsCOOOH(a~)

(12) (13) (14)

C6HsCOOOH(.~) + C6HsCHO(.~)---> 2 C6H~COOH(.~)

(15)

C6H~COOH(~) --> C6H6 + C02.

(16)

According to reactions (12)-(16), CO2 results from the oxidation of toluene to benzoic acid whose traces were found in our experiments and its subsequent photodecarboxylation. It must be reported that some runs carried out by using benzoic acid at the same experimental conditions used for toluene photooxidation, afforded benzene and CO2 in large amounts. It is well known [7, 8] that ethanoic acid is easily photodecarboxylated in a gas-solid regime in the presence of irradiated polycrystallme semiconductor oxides. The small amounts of phenol are probably due to an attack of benzene by OH radicals (see eqn.(16)). The progressive deactivation of the catalyst in the absence of water could be due to some surface dehydroxylation and/or to the formation of stable intermediate species which can not evolve in the absence of water so that they remain strongly adsorbed on the catalyst surface. It is worth reporting that the

672 catalyst was found of ochre colour at the end of the run and this colour completely disappeared by irradiating the catalyst in liquid water. The presence on the catalyst surface of sites with different oxidation strength can explain the role played by water and the deactivation-activation pattern exhibited by the catalyst. The partial restoration of activity when water was restored to the reacting ambient could be due to the rehydroxylation of surface sites. The finding that benzene is produced only in the first hours of irradiation while the partially spent catalyst never produced benzene, may be justified by considering that the fresh catalyst has higher oxidant properties due to the presence of highly oxidant hydroxyl groups. Benzaldehyde, instead, can be obtained both on these sites and on less oxidant sites. The experimental results suggest that the more oxidizmg sites react irreversibly under irradiation. By concluding, the photo-oxidation of toluene mainly to benzeldehyde and benzene (as a transient product) was proved to occur in mild conditions in gassolid regime. The presence of sites with different oxidant properties is proposed and a preliminary FTIR mvestigation indicates a weak interaction between toluene molecules and TiO2 surface. ACKNOWLEDGEMENT The authors wish to thank the 'TIinistero delrUmversith e della Ricerca Scientifica e Tecnologica" (Rome) for financially supporting this work.

REFERENCES 1. M. Schiavello (ed.), Photocatalysis and Environment. Trends and Applications, Kluwer, Dordrecht, 1988. 2. V. Augugliaro, L, Palmisano, A. Sclafani, C. Mmero and E. Pelizzetti, Toxicol. Envir. Chem., 16 (1988) 89. 3. E. Pelizzetti and N. Serpone (eds.), Photocatalysis. Fundamentals and Applications, Wiley, New York, 1989. 4. E. Pelizzetti and M. Schiavello (eds.), Photochemical Conversion and Storage of Solar Energy, Kluwer, Dordrecht, 1991. 5. D. F, Ollis and H. A1-Ekabi (eds.), Photocatalytic Purification and Treatment of Water and Air, Elsevier, Amsterdam, 1993. 6. T. Ibusuki and K. Takeuchi, Atmos. Environ., 20 (1986) 1711. 7. L. Palmisano, M. Schiavello, A. Sclafani, S. Coluccia and L. Marchese, New J. Chem., 12 (1988) 847. 8. T. Matsuura and M. Anpo (eds.), Photochemistry on Solid Surfaces, Elsevier, Amsterdam, 1989. 9. M.L. Sauer and D. F. Ollis, J. Catal., 158 (1996) 570. 10. A. J. Malta Vidal, J. Soria, V. Augugliaro, V. Loddo, Chem. Biochem. Eng. Quart., 1997 (in press). 11. C. Morterra, J. Chem. Soc. Faraday Trans. I, 84 (1988) 1617. 12. E. A. Paukshits and E. N. Yurchenko, Russ. Chem. Rev., 52 (1983)242. 13. R. I. Bickley, G. Munuera and F. S. Stone, J. Catal., 31 (1973) 398.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

673

Oxidation Processes on Stoichiometric and Nonstoichiometric Hydroxyapatites H. Hayashi', H. Kanai b, Y. Matsumura c, S. Sugiyama', and J.B. Moffat d "Department of Chemical Science and Technology, Faculty of Engineering, The University of Tokushima, Minamijosanjima, Tokushima 770, Japan ~Faculty of Life Science, Kyoto Prefectural University, Simogamo, Sakyo-ku, Kyoto 606, Japan COsaka National Research Institute, AIST, Midorigaoka, Ikeda, Osaka 563, Japan dDepartment of Chemistry and the Guelph-Waterloo Center for Graduate Work in Chemistry, University of Waterloo, Waterloo, Ontario, Canada N2L 3G 1

1. INTRODUCTION Solids whose crystallographic structures are stable at high temperatures, permit ionic exchange and changes in their elementary composition while remaining structurally invariant are of considerable interest for utilization as and fundamental studies of heterogeneous catalysts. Solids with these properties but additionally with catalytic functionalities which are altered by one or more of the aforementioned changes are particularly useful in understanding the nature of catalytically active sites. Calcium hydroxyapatite [Ca~o(PO4)6(OH)2] (Fig. 1) is an inorganic solid which exhibits these characteristics [ 1]. It does not readily lose hydroxyl groups and the lattice is believed to be stable 9. I

.,

~" ,<' "

.O I

/.~

0l.

~

I

,..

9..,,

e - ~ ~ ~

,

b

~. I I

"~<" t ,,

.01 ,.c.,,_.

I I

/

ol '

~

_

_

--

-

w~

~. I

"k,

"v" 9 %'-

J"

ca,a," X

01 ca

..~6~--tl,,'cD

.tiC)

J"

"~ ~. "

9

9

lr"

..

I

I T ],, . . . . .

" ~ / vl

o,,-,

I I

"~k

""

9

.~ ' ""

,i

~

S

,~'" I

Figure 1. Stoichiometric calcium hydroxyapatite structure projected on a, b, plane. (Suzuki)

674 up to 1000~ [2]. X-ray diffraction and transmission electron microscopic investigations have established its structural characteristics [3-7] and a recent application of atomic force microscopy has shown that the crystal surface structure is highly ordered with atomic spacings of 0.68 and 0.43 nm [8]. Hydroxyapatite has a hexagonal structure constructed from columns of Ca and O atoms which are parallel to the hexagonal axis. Three oxygen atoms of each PO 4 tetrahedron are shared by one column, with the fourth oxygen atom attached to a neighbouring column. The hexagonal unit cell of hydroxyapatite contains 10 cations located on two sets of nonequivalent sites, 4 on site (1) and 6 on site (2). The calcium ions on site (1) are aligned in columns, while those on site (2) are in equilateral triangles centered on the screw axes. The site-(1) cations are coordinated to six oxygen atoms belonging to different PO4 tetrahedra and also to three oxygen atoms at a large distance. The site-(2) cations are found in cavities in the walls of the channels formed between the cations and O atoms. The hydroxyl groups are situated in these channels and probably form an approximately triangular coplanar arrangement with the Ca ions.

2. EXPERIMENTAL

Calcium hydroxyapatites (CaHAp) were prepared from Ca(NO3)2.4H20 (BDH AnalaR) and (NH+)2HPO+ (BDH AnalaR) according to the method previously described [9]. Strontium hydroxyapatites (SrHAp) were similarly prepared from the corresponding strontium compound (Wako Pure Chemical). Lead cation was ion-exchanged into the hydroxyapatites following a procedure in the literature [ 10]. Compositions of the prepared solids were obtained from ion chromatographic analysis (Dionex 4500 i) of the solutions remaining from the synthesis, and/or atomic absorption spectrometry and/or inductively coupled plasma (ICP) spectrometry (SPS-1700, Seiko). Reactions were performed in a conventional fixed-bed continuous flow reactor operated under atmospheric pressure. Details of the method of operation and analytical procedures have been provided previously [11,12]. Surface analyses by XPS were obtained with a Shimadzu ESCA-1000AX spectrometer under conditions described previously [ 11]. Powder X-ray diffraction (XRD) patterns were recorded with a Rigaku RINT 2500 X using monochromatized MgKa radiation. ESR spectra were recorded with a JES-TE-300 spectrometer at 9.30 GHz and room temperature [ 12].

3. RESULTS AND DISCUSSION Calcium hydroxyapatites of various stoichiometries catalyze the partial oxidation of methane with oxygen at 600~ to carbon monoxide, carbon dioxide, formaldehyde, hydrogen and water. While the stoichiometric composition (Ca/P=l.65) produced predominantly carbon dioxide the molar ratio of H2/CO~ in the products was as high as 1.3. With decrease in the Ca/P ratio the H2/CO~ ratio decreased while the selectivities to CO7. and CO decreased and increased, respectively, the latter at least down to a Ca/P value of 1.61, after which it remained approximately unchanged. Stoichiometric CaHAp apparently predominantly catalyzes the reaction CH4 + 02-" C02 + 2H2

(1)

675

while the process (2)

CH4 + O2 -" CO + H 2 + HzO

occurs on the nonstoichiometric analogue. Since a band at 727 cm l (attributed to PzO~) appears in the infrared spectra of the nonstoichiometric CaHAp but not in that for the stoichiometric composition such species are tentatively assumed to be responsible for the selective formation of carbon monoxide. Separate experiments with Ca2PzO7 yielded selectivities of approximately 20% for H2CO although the conversion of CH4 was quite low. In contrast to the aforementioned, with strontium hydroxyapatite the selectivities are essentially independent of the composition of the catalyst, as is the H2/COx ratio of 1.5. The partial introduction of lead as cations in the hydroxyapatites shifts the methane conversion process to favour oxidative coupling products (Fig. 2). 15 100 80

141

g lo

.;

+o .;'.

g

+o m

20 0

CO 0

0.05

~

~'I 0.10

0

0.30

P b / l P b + C a ) molar ratio in sample Figure 2. Oxidative coupling of methane with dioxygen over lead-calcium hydroxyapatite. Reaction conditions: reaction temperature, 700~ catalyst, 0.30g; partial pressure of methane, 29 kPa; dioxygen, 4 kPa; total flow rate, 0.9 dm 3 hr~; time on stream, 3 hr. Although the introduction of small quantities of lead produced dramatic shifts in the selectivities to CzH4 and CzH6 these became relatively constant for compositions with Pb/(Pb + Ca) equal to 0.02 or greater. Concomitantly the selectivity to CO decreased to vanishingly small values. The conversion of methane and the various product selectivities achieve approximately constant values for surface concentrations of lead equal to approximately 3 mol % which corresponds to a Pb/(Pb + Ca) ratio of 0.02, suggesting that the changes produced by the introduction of Pb result primarily from its presence in the surface, rather than the bulk. Lead species in the surface apparently assist in the stabilization of methyl radicals, while the formation of ensembles of these species hold such stabilized radicals in sufficiently close proximity to permit lateral interaction and the formation of Cz hydrocarbons [ 13].

676 Further information on the catalytic properties of stoichiometric and nonstoichiometric CaHAp may be obtained from studies on the adsorption and dehydrogenation of methanol. With stoichiometric CaHAp methanol decomposes at 600~ to produce predominantly carbon monoxide (Table 1) whose selectivity diminishes as the Ca/P ratio decreases while those to formaldehyde and dimethyl ether increase. Infrared spectra show that methoxy groups are formed on the surface of both the stoichiometric and nonstoichiometric catalysts. The results from temperature-programmed desorption experiments together with those from infrared spectroscopy suggest that the acidic sites found on the nonstoichiometric CaHAp catalyze the dissociative adsorption whereas the basic sites on the stoichiometric analogue catalyze the C-H bond scission and formation of CO and H2. Table 1 Decomposition of methanol" on stoichiometric and nonstoichiometric calcium hydroxyapatite Hydroxyapatite Composition

Conversion of

Ca/P

Methanol (%)

H2CO

CO

C02

1.51 1.55 1.61 1.65

85.2 96.6 95.0 100.0

51.1 36.8 9.3 0.2

11.1 46.1 60.7 79.1

0.0 0.6 3.8 10.4

" 600~

Selectivityb (%)

3h, Pcmo.=2kPa, F = 0.9 dm 3 hl; b CH4, (CH3)zO and C2 are also produced

Studies of the oxidation of CO may also be advantageously employed to obtain further information on the hydroxyapatites. Over nonstoichiometric CaHAp (abbreviated as Ap-N) the rate of CO oxidation was independent of P(Oz), while with Ap-S the rate was higher and dependent on P(O2) (Fig. 3). Both CaHAp samples were dependent on P(CO). Um

lop 2 5

T.:

Ap-S

ao

Pco " 2 2 kPa _

E

10 1 E

5

(:~

Ov

o

o

o

'I

5

l

lO

Pressure of 02 / kPa

Figure 3. Rate of formation (600~ of carbon dioxide from carbon monoxide on stoichiometric (Ap-S) and nonstoichiometric (Ap-N) hydroxyapatite for various P(O2).

677 Evidence for the formation of paramagnetic species on Ap-S which are apparently absent on Ap-N can be obtained from ESR spectra (Fig. 4). New ESR signals, whose g-values suggest the 12.0051 A2.0048 ~ll

2.0262

2.0247 2.0265

I

2.0035 2.0053 Ap-N

. . . . . . . . .

I 2.0239

327

1.9862

A

I

"/ ~,*-'.w~

I 2.0049

I

330

"

2"0160

1.9871 ........

I

333

~

I

ii

I 1.9836

ill ~

"'

~III illl I

k/ "'

~

,.0""

I 2.0102

336 327

Magnetic field / mT

I

330

il~ I I

~i

,,'

ill

II

II

qll

I 2 0013 I

333

336

Magnetic field / mT

Figure 4. ESR signals for hydroxyapatites Figure 5. ESR signals for Ap-S after adsorption evacuated at 600~ Microwave power, of oxygen for 1 day at room temperature followed 0.2 mW. by evacuation for 1 min. Dashed:simulated presence of 02, appear after exposure of Ap-S to 0 2 but these do not appear with Ap-N [ 12] (Fig. 5). The g-values of the ESR lines found with Ap-S and Ap-N after evacuation at 600~ are similar to those expected where an electron is trapped in an oxygen vacancy [14,15], which may result from dehydration of hydroxyl groups [ 15,16],

Calo(PO4)6(OH)2-, Caio(POa)6(OH)z.zxOxDx + xH20

(D = vacancy, x < 1)

(3)

The interaction between molecular oxygen and 02- ions in the columns of the hydroxyapatite may result in the formation of 02 species and O centers, 02 + 0,,2- = 02 + 0~

(4)

The absence of 02 radicals on the nonstoichiometric hydroxyapatite may be attributed to the interaction of 02- in the column with P2074 to form two PO43 groups with the concomitant elimination of 02 ions in the column:

02- + P2074 --, 2PO43

(5)

The 02 species formed on the stoichiometric hydroxyapatite can react with carbon monoxide, CO + 02 -" C02 + O

(6)

678 which is then dependent on the pressure of oxygen through equation (5). In turn the highly active species O [ 18] will also react with carbon monoxide, (7)

CO + O + 0 , -- C02 + O,2.

Earlier and continuing work from our laboratories has shown that the introduction of small quantities of tetrachloromethane (TCM) into the feedstream for methane partial oxidation or oxidative coupling produces beneficial increases in the selectivities to Ca+ hydrocarbons, particularly ethylene, with many heterogeneous catalysts [19]. Consequently it is of interest to examine the effect of this additive with other oxidation processes, such as the conversion of CO to COz. The addition of small quantifies of tetrachloromethane to the feedstream in the CO oxidation process inhibits the formation of CO2 on stoichiometric HAP but has relatively little effect on a nonstoichiometric sample (Fig. 6). With both catalyst samples the inhibition is enhanced with

20

)It

15-

I=

0

-0

e,,i,

i._ t)

:D,

.65

10-

i=

o o

0 0

III

Ap-1.51 0

:1

A P - 1.51

0

+ ~::Cla i 1

i

'

2

Time-on-stream

',

'i

3

4

/ h

Figure 6. Oxidation of carbon monoxide with and without carbon tetrachloride over hydroxyapatites at 600~ increasing time-on-stream. Both XRD and XPS analyses demonstrate that the inhibition results from the interaction of TCM with the surface to convert the hydroxyapatite, at least partially to its chlorinated analogue, chlorapatite, which suppresses the oxidation of CO to COz Earlier work from one of our laboratories has shown that, with 12-tungstophosphoric acid as a heterogeneous catalyst, the partial oxidation of methane in the presence of small partial pressures

679 of TCM produces high selectivities to methyl chloride. For comparison purposes similar experiments have been performed with strontium hydroxyapatites ion-exchanged with lead (Table 2). Table 2 Conversion of Methane and Product Selectivities for Strontium Hydroxyapatite with and without ion-exchanged lead Selectivities (%) or Conversion (%) SrHAp Time-on-stream (h)

Pbo.26SrHAp

0.5

Selectivity to CO Selectivity toCOz Selectivity toCH3C1 Conversion ofCH 4 T = 773 K, P(CH4) = 28.7 kPa, TCM (0.17 kPa).

6

0.5

6

52.7 (100) 100 (45.9) 0 (0) 0 (20.9) 47.3 (0) 0 (0) 100 (100) 100 (5.7) - (0) - (54.1) - (0) - (73.4) 0.4 (0.3) 0.1 (0.6) 3.1 (2.1) 3.0 (0.8) P(O2) = 4.1 kPa. Values in brackets obtained in the presence of

With SrHAp and in the absence of TCM the products are approximately divided between CO and COz. On addition of a small partial pressure of TCM to the feedstream the CO2. vanishes in favour of CO at relatively short times-on-stream. For longer times-on-stream the selectivity to CO decreases while methyl chloride appears and little or no CO2 is evident in the product stream.

:

9

s

!

-

J

A A

.c_ 25

"A

I

" A

0

&

35

A

'

'

A

A

/.5 25 2(9 / degrees

I

&~&

35

/.5

Figure 7. XRD patterns of SrHAp (A) and Pbo.26SrHAp (B) previously employed in obtaining the results shown in Table 2 but after 6 h on-stream. Symbols: Open circles, strontium hydroxyapatite. Open triangles, strontium chlorapatite. Closed triangles, lead chlorapatite.

680 In contrast, with Pbo.26SrHAp and for short times-on-stream little or no CO is found in the product either with or without TCM and in the presence of TCM the selectivities to both CO and methyl chloride are vanishingly small. However for longer times-on-stream substantial selectivities to both CO and methyl chloride are observed to the detriment of those to CO2. The induction period for the formation of methyl chloride and the increase and decrease of selectivities to CO and CO2 respectively, provideevidence for the incorporation of TCM or fragments thereof on and in the catalyst. XRD and XPS measurements show that C1 is introduced into the surface and bulk of PbSrAp, apparently forming the corresponding chlorapatite, although not precluding the existence of chlorine in additional configurations (Fig. 7). Concomitantly a portion of the cationic lead is reduced to the metallic state.

4. CONCLUSIONS 1. The catalytic properties of hydroxyapatites are dependent upon the nature of the cation and the ratio of the number of cations to the number of phosphorus atoms. 2. With calcium as the cation the oxidation of methane on stoichiometric CaHAp produces predominantly COz and H 2 while with the nonstoichiometric catalyst the principal products are CO and H 2. 3. The formation of carbon monoxide is attributed to the presence of P2074 in the stoichiometric CaHAp. 4. With strontium as the cation the selectivities to the products are independent of the Sr/P ratio of the catalyst. 5. Substitution of relatively small fractions of the calcium cations by lead depresses the selectivity to COx while generating significant selectivities to C2§ hydrocarbons as well as increasing the conversion of methane. 6. The surface lead species are believed to stabilize the methyl radicals formed from methane while the formation of ensembles of surface lead facilitates the generation of Cz hydrocarbons. 7. Since the rate of oxidation of CO to form CO2 on the stoichiometfic CaHAp(S) is higher than that with the nonstoichiometrioc form (N) such differences may contribute to the disparity of the selectivities to CO on the S and N forms of CaHAp. 8. Paramagnetic species, 02 form from 02 on the S form of CaHAp but not on Ap-N. The former species may contribute to the conversion of CO to CO2. 9. Introduction of tetrachloromethane (TCM) to the feedstream in the CO oxidation process inhibits the formation of CO 2 on stoichiometric CaHAp but has relatively little effect on a nonstoichiometric sample. The effect of TCM results from the interaction of chlorine with the surface of the catalyst to form chlorapatite. 10. In the presence of TCM and with SrHAp as the catalyst only CO and no CO2 is formed from CH4 at short times-on-stream, while at longer times on stream methyl chloride begins to form to the detriment of CO.

5. ACKNOWLEDGEMENTS This work was partially funded by a "Grant for Natural Gas Research" of The Japan Petroleum Institute to S.S. and the Natural Sciences and Engineering Research Council of Canada to J.B.M.

681 We cordially thank Professor Satohiro Yoshida of Kyoto University for providing the ESR simulator program.

6. REFERENCES 1. J.C. EUiott, Structure and Chemistry of the Apatites and Other Calcium Orthophosphates, Elsevier, Amsterdam, 1994. 2. D.E.C. Corbridge, The Structural Chemistry of Phosphorus, Elsevier, Amsterdam, 1974. 3. H. Ji and P.M. Marquie, J. Mater. Sci. Lett, 10 (1991) 132. 4. W.J. Landis, J. Moradian-Oldak and S. Weiner, Connect. Tissue Res., 25 (1991) 181. 5. W.J. Landis and M.J. Glimcher, J. Ultrastruct. Res., 63 (1978) 188. 6. M.I. Kay, R.A. Young and A.S. Posner, Nature, 204 (1964) 1050. 7. H.C.W. Skinner, H.T. Hunt and J. Griswold, J. Phys. E. Sci. Instrum., 13 (1980) 74. 8. L.M. Siperko and W.J. Landis, Appl. Phys. Lett., 61 (1992) 2610. 9. E. Hayek and H. Newesely, Inorg. Synth., 7 (1963) 63. 10. A. Bigi, M. Gandolfi, M. Gazzano, A. Ripamonti, N. Roved and S.A. Thomas, J. Chem. Soc. Dalton Trans., (1991) 2883. 11. S. Sugiyama, T. Minami, H. Hayashi, M. Tanaka and J.B. Moffat, J. Solid State Chem., 126 (1996) 252. 12. Y. Matsumura, H. Kanai and J.B. Moffat, to be published. 13. Y. Matsumura, S. Sugiyama, H. Hayashi and J.B. Moffat, Catal. Lett., 30 (1995) 235. 14. J.H. Lunsford and J.P. Payne, J. Phys. Chem., 70 (1966) 3464. 15. H. Matsuhashi and K. Arata, J. Phys. Chem. 99 (1995) 11178. 16. T. Kijima and M. Tsutsumi, J. Amer. Ceram. Soc., 62 (1979) 455. 17. G.R. Fischer, P. Bardhan and J.E. Geiger, J. Mater. Sci. Lett, 2 (1983) 577. 18. M. Che and A.J. Tench, Adv. Catal., 31 (1982) 77. 19. J.B. Moffat, S. Sugiyama and H. Hayashi, Catal. Today, (in press).

This Page Intentionally Left Blank

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

683

Oxidative Coupling of Methane in Solid Oxide Fuel Cells Guo Xiu-Mei * Kus Hidajat Chi-Bun Ching Department of Chemical Engineering, National University of Singapore, Singapore 119260 Chen Hong-Fang Chemical Engineering Department, Tianjin University, Tianjin 300072, China 1. INTRODUCTION A large amounts of chemicals are produced via an exothermic heterogeneous catalytic reaction pathway today. Due to the high exothermicity of these reactions, large amounts of thermal energy are generated. In conventional reactors, the enthalpy is converted into heat, a low quality type of energy. Fuel cells have the advantages of high efficiency, low pollution, simplified system, high power density, which make them very attraction for utility and industrial applications. However, in fuel cells, the enthalpy is converted directly into electrical power, but fuels such as H2, CO, CH4 are converted into the products of total oxidation, i.e. CO2 and H20, which are not more valuable than the fuels themselves. Moreover, CO~ is considered to be responsible for about half of the global warming rate. Obviously, if electrical power and useful chemicals can be cogenerated in a fuel cell, it will be highly desirable. Vayenas and his co-workers [1, 2] have already studied cogeneration of electrical power and NO, styrene and so on before in the type of solid oxide fuel cells. Since the pioneering work of Keller and Bhasin[3], which showed that methane can be converted into ethane and ethylene via oxidative coupling of methane(OCM), a strong effort has been conducted to develop more efficient catalysts and to understand the reaction mechanism. However, the low C2 yield ( maximum between 25% and 30%), the cost of 02, the problem of heat removal, product separation and etc. need to be studied further before a profitable commercial process can be developed. Cogeneration of electrical power and useful chemicals opens up a new field for the study on the OCM. Besides the above advantages, it can separate oxygen from air, simplify the separation processes and decrease the amount of the gases out the reactors. In this present study, the cogeneration of ethane and ethylene and electrical power is investigated in Ag-YSZ and 1wt%Sr/La203-Ag-YSZ solid oxide fuel cells.

684 2. EXPERIMENTAL

2.1 Experimental Apparatus A schematic diagram of the solid oxide fuel cell is shown in figure 1. A electrodecatalyst is deposited on the surface of the oxygen ion conducting solid electrolyte (8% mole percent yttria stabilised zirconia, YSZ) disk. The anode is exposed to the reactant CH4, the cathode is exposed to an oxygen-containing gas and the electrical power is produced in a nearly reversible fashion. In the cathodic cell, Oeu) is reduced to 0 2. on the cathode. Then 0 2. is transported through the YSZ in a lattice with anion vacancies to the anode, on which oxygen transported is activated and reacts with CH4 to form C2H6,C2H4, CO2, etc. and discharge simultaneously.

CH4~

C2H4,C2H6, "

resistence //box

O2 Figure 1 Schematic diagram of the solid oxide fuel cell The configuration of the solid oxide fuel cell is shown in figure 2. It consists of a YSZ tube ( 1 lmm OD, 9mm ID) opened at both ends and enclosed in a quartz tube( 20mm OD, 17mm ID). The porous Ag, porous metal oxide lwt%Sr/La203 and Ag are deposited on the inside and outside surface of YSZ tube as cathode and anode. During the experiments, CH4 diluted in N2 is fed in the annular space between the quartz and the YSZ tube at a total flow rate of 100ml/min. The fuel cell is operated under atmospheric pressure. The temperature is controlled within 1-~2 K by a EUROTHERM808 temperature controller. Voltage and current generated are measured by two HP 34401A multimeters. N72 resistance box is used to change the load. Current collection is via Ag wire. Analysis of reactants and products is performed using a gas chromatograph with two columns packed with Porapak Q and molecular sieve 5A. The main detectable products are C2H6,C2H4,CO2, H2, H20. The amount of CO is negligible. Carbon atom balance does not exceed +2%, which indicates that no

685 consistent loss of material, no significant formation of other oxygenated species, and coking of the catalyst. CH4 conversion and C2 selectivity are calculated by: (1)

XCH4 : (FOcH4 -- F CH4) / FOcH4 Sc 2 -- 2(Fc2H4 -k-Fc2H6) / (F~

(2)

-- FCH4 )

metal oxide+Ag YSZ

\_j CH4' N2

V

Ag /1

......

" "i"

/ C2H 6, C2H4, CO 2, H20, N2

3

~_J

V

Figure 2 The configuration of solid oxide fuel cell 1-quartz tube; 2-YSZ tube; 3-Ag wire; 4-multimeter used as galvanometer; 5-multimeter used as voltmeter; 6-resistance box

2.2 Electrode-catalyst Preparation The solid oxide membrane consists of three layers: YSZ layer, electrode layer and catalyst layer. The YSZ layer (YSZ tube) is composed of 8mo1% yttria in zirconia, which is available commercially. The porous electrode layer is made of AgNO3, prepared according to the following procedure. First, the YSZ tube is cleaned by rinsing in distilled water in an ultrasonic bath, followed by drying and calcining for 4h at 1223 K. After the YSZ tube is cooled down naturally, it is slightly heated and coated using AgNO3 in nitrocellulose/butyl acetate, followed by drying and calcining. The procedure is repeated several times in order to achieve an electrode resistance of less than 0.30f2. The catalyst layer is composed of l wt%Sr/La203, prepared from Sr(NO3)2, La203. At first, l wt%Sr/La203 is prepared by impregnation and ground to powder. Then, the powder product of 1wt%Sr/La203 is added to the nitrocellulose/butyl acetate and applied

686 to the external surface of the Ag-YSZ followed by drying and calcining in air.

tube using the method shown in Figure 3,

Figure 3 Coating apparatus 1-rotating axle; 2-coating cell; 3-Ag-YSZ; 4-coated mixtures; 5-plug

3. R E S U L T S AND DISCUSSION

3.1 Effect of composition of catalyst-electrode Table 1 summarises the C2 selectivity, CH4 conversion and current generated in the AgYSZ and 1wt%Sr/La203-Ag-YSZ solid oxide fuel cells. It can be seen that C2 selectivity in lwt%Sr/La203-Ag-YSZ solid oxide fuel cell is obviously higher than that in Ag-YSZ solid oxide fuel cell, but CH, conversion and current generated in l wt%Sr/La203-AgYSZ solid oxide fuel cell is lower than that in Ag-YSZ solid oxide fuel cell. It indicates that coating the l wt%Sr/La203 on the Ag-YSZ surface is favourable for producing the active oxygen species responsible for C2 hydrocarbons. It demonstrates that the composition of catalyst-electrode make a great effect on the OCM reaction and current generated. Most metal oxides that have been shown to be good catalysts for OCM reaction are poor electrical conductors. On the other hand, although metals show good electrical conductivity, they show poor selectivities and activites to C2 hydrocarbons. Therefore, it is greatly attractive if the advantages of the metal oxides and metals can be made use of fully at the same time. The future work should be done further in this area. The following studies focus on the OCM reaction in 1wt%Sr/La203-Ag-YSZ solid oxide fuel cell. Table 1 Effect of composition of electrode-catalysts on fuel cells Composition of ElectrodeCH, C2 Catalysts Conversion(%) Selectivi~'(% ) Ag-YSZ 3.2 44.0 2.4 82.9 1wt%Sr/La~O3-A~-YSZ

Currcnt(mA) 56.2 48.5

687 (Reaction conditions: T=730 C, R=0.1ff~, anode: Q=100ml/min, PcH4=20.3kPa, cathode: Q=250ml/min, Po2=20.3kPa) 3.2 Effect of Current Generated

Figure 4 shows the relationships between current generated and C~ selectivity and CH4 conversion in lwt%Sr/La203-Ag-YSZ solid oxide fuel cell. In order to compare with the results in a conventional reactor, the current generated in the solid oxide fuel cell is converted into oxygen flux according to the Faraday's law:

Jo2-- I/4F

(3)

in which F stands for the Faraday constant. The current is adjusted by putting a load in the outer circuit. It can be seen that the CH, conversion increases and C2 selectivity decreases with a rise in oxygen flux both in the conventional reactor and in the fuel cell. It indicates that the products of C2 hydrocarbons are more reactive than methane and much easier to form the total oxidation products. A comparison of the results achieved in the fuel cell and in the conventional reactor demonstrates that the C2 selectivity obtained in the fuel cell is 10%- 20% higher than that obtained in the conventional reactor. It indicates that it is much easier for oxygen transported to produce the active oxygen species than for oxygen supplied by gas phase oxygen. CH, conversion in the fuel cell is slightly lower than that in the conventional reactor. In addition, it can be observed that the difference of C2 selectivity between the fuel cell and the conventional reactor is reduced slightly with increasing oxygen flux. This may be due to the fact that the reaction producing gas phase oxygen is enhanced with a rise in the current generated. Moreover, it can be found that the results in solid oxide fuel cell are different from the results of Otsuka et al[6]. They observed that the C2 selectivity increased with a rise in current using BaCO3-Au as the electrode-catalyst in the solid oxide fuel cell. It indicates that there is a dramatic effect of the properties of the electrode-catalyst on the reaction results. Figure 5 shows the current-voltage and current-power plots at constant temperature, feed composition, total flow rate. It is noted that the voltage obtained under open-circuit reaches 0.84v and decreases quasi-linearly with increasing current. The thermodynamic activity of oxygen absorbed on the catalyst surface during reaction can be written as [7]: a = (0.21) 1/2 exp(2FE /

RT)

(4)

in which R is gas constant, F is Faraday constant. The oxygen activity a on the membrane surface calculated is less than PO21/2measured. That means that the overall reaction is limited by the adsorption of oxygen. Figure 5 also shows that ohmic overpotantial is the dominant source of polarisation at the temperature range. The electrical power output increases with current until it reaches a point above which it decreases. In other words, there is an optimal load so that power output

688

becomes maximum. This optimal load equals the sum of electrolyte resistance and electrode resistance. 100 80

r 09

4

oo

S

60

[]

Sc2, reactor

40

x

Xch4, reactor,,.~

J"

~

"

,

X o 2 .Ix

"

o-e.

20 0 0

i

J

5

10

0 15

Jo2 10e-8 mol/s Fig. 4 C2 selectivity and CH4 conversion as a function of oxygen flow rate in the fuel cell and in the conventional reactor (T=730 C, anode: Q=100ml/min, Pch4=20.3kPa, cathode: Q=250ml/min,Po2=20.3kPa)

1

>

10

0.8

8

0.6

6

0.4

4

0.2

2

0-0

"13

g, r

0 10

20

30

40

50

I /mA

Fig. 5 Current-votage and current-power (reaction conditions are the same as Fig.4)

3.3 Effect of Temperature Figure 6 shows the effect of temperature on C~ selectivity and CH, conversion at constant feed composition and total flow rate. It can be seen that the C2-selectivity initially increases with a rise in temperature and then appears to reach a plateau and for

689

T>750 ~ it starts decreasing. The C2-selectivity shows maxim at 750 ~ There is the optimum temperature for the synthesis of C2 hydrocarbons. It also indicates that the total oxidation reaction is much easier to occur under high temperatures. The CH4 conversion shows a increase with increasing temperature. Figure 7 shows the effect of temperature on the current generated. It can be seen that the current generated increases with temperature. This is due to the increased ionic conductivity of the YSZ at the higher temperatures. 100

5

90 ``9` o

80

r O9

70

3o

X

x 2

60

x

50 660

o-.,e,

Xch4

h

,

i

I

i

b

680

700

720

740

760

780

0 800

T/C

Fig.6 C2 selectivity and CH4 conversion as a function of temperature( R=0.1 ohm, anode: Q=100ml/min, Pch4=20.3 kPa, cathode: Q=250ml/min, Po2=20.3kPa) 100 4,

Exp.

80 <

E

60 40 20

650

i

I

700

750

800

T/C

Fig.7 Current generated as a function of temperature (reaction conditions are the same as Fig.6)

690 Figure 8 shows the effect of temperature on the ethylene-to-ethane ratio. It is noted that the ratio increases with increasing temperature. The result is consistent with the result in a conventional reactor. It indicates that ethylene is formed from ethane as a secondary product. The reaction of ethane oxidative dehydrogenation is accelerated with a rise in temperature,

2.4 9 Exp.

"r0 "1r

0

1.6

1.2 0.8 0.4

0 650

700

750

800

TIC Fig.8 Effect of temperature on C2H4/C2H6 ( reaction conditions are the same as Fig.6)

3.4 Effect of Feed Composition Figure 9 shows the effect of feed composition on C2 selectivity and CH4 conversion. N2 is used as a diluent. It can be seen that the C2-selectivity increases and CH4 conversion decreases with a rise in feed ca4 concentration. This may be due to the fact that the amount of oxygen transported depends on the current generated, while the current generated is low. Figure 10 shows the effect of feed composition on current generated. The current generated increases weakly with increasing feed methane concentration. The amount of oxygen limits the reaction.

691

90

5

.

l

Sc2

x Xch4

, 4

85

9

g 03 80

9

x

2 ~ 1

75

0

20

40

60

0 100

80

ych4/% Fig. 9 C2 selectivity and CH4 conversion as a function of feed methane mole fraction (reaction condition: T=730 C, R=O.1 ohmi, anode: Q=lOOml/min, cathode: Q=250ml/min, Po2=20.3kPa) 60 Exp.

.

56 < 52 E -48

C

44 40

0

~

k

i

J

20

40

60

80

100

ych4/% Fig.lO Current generated as a function of feed methane mole fraction( reaction conditions are the same as Fig.9) 4. C O N C L U S I O N S 1. The O C M reaction can be carried out in a solid state electrochemical reactor with 0 cogeneration of electrical energy when l wt~Sr/La203-Ag is used as the electrodecatalyst.

692 2. The CH4 conversion, C2-selectivity and current generated is functions of temperature, feed composition and properties of the electrode-catalyst. 3. The C~-selectivity can exceed 90% but the power output as well as CH4 conversion are low. NOMENCLATURE E Fi I J P~ Q R Sc2 T V P XCH, yCH4

open-circuit voltage mole flow rate current oxygen flux i component partial pressure total flow rate resistance C2-selectivity temperature voltage power CH4 conversion CH, mole fraction

Greek Symbols c~ oxygen activity

/v / ml/min /mA / ml/s /Pa ml/min f2 /% /C /v /w /% /%

a t m 1/2

REFERENCES (1) Wynveen, R.O., Fuel Cells, Reinhold, New York(1963) (2) Vayenas, C. and R.Farr, Science, 208, 593(1980) (3) Vayenas, C. and J.N. Michaels, J. Catal., 85,477(1984) (4) Lee, J.S. and S.T. Oyama, Catal. Rev.-Sci. Eng., 30, 249(1988) (5) Omata, K., et al., Appl. Catal., 52, L1(1989) (6) Otsuka, K., et al., Catal. Today, 6,587(1990) (7) Stoukides, M., Ind. Eng. Chem. Res., 27,1745(1988)

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

P A R T I A L O X I D A T I O N OF M E T H A N E MEMBRANE REACTOR.

693

TO SYNTHESIS GAS IN A FAST FLOW

M. Alibrando and E.E. Wolf, Chemical Engineering Department, University of Notre Dame, Notre Dame, IN, USA. INTRODUCTION

The catalytic partial oxidation of methane to syngas was first reported by Pettree as early as 1946 [ 1]. More recently, several investigators have studied the reaction mostly on Ni supported catalysts [2-4] and on Ru oxides [5]. In the past several years, Schmidt and coworkers [6-8] have studied the reaction using various noble metals in a monolith reactor. In this configuration, high methane conversions and syngas selectivities were achieved in a reactor operated under autothermal conditions Despite the impressive results achieved in partial oxidation, a major difficulty still exists. For autothermal behavior to be attained a low methane/oxygen ratio mixture must be fed to the reactor. When such mixture contacts a high temperature surface, the potential for an explosion exists. The goal of this work was to design a reactor which will allow for a low methane/oxygen ratio mixture to be fed at high flow rates while separating the methane and oxygen feeds. Based on our previous studies concerning the distribution of oxygen during the methane oxidative coupling reaction [9], we designed a configuration that uses a porous membrane reactor. The reactor consists of a porous permeable ceramic membrane tube placed concentrically inside a quartz reactor (Figure 1). Methane is fed to the reactor on the shell side and flows in the annular space between the membrane tube and the reactor wall. Oxygen is fed to the tube side and flows exclusively inside the membrane tube as the upper portion has been made impermeable by a ceramic glaze. The end at the lower portion of the membrane tube is closed, leaving the permeable length at the bottom of the membrane tube as the only exit for the oxygen. The permeable length is surrounded by catalyst on the outside of the tube allowing the oxygen and methane to be mixed only over the catalyst surface. Because 100% oxygen conversion is achieved, the possibility of developing a gas mixture with high oxygen content is very small. Therefore, a high overall flow rate with low methane/oxygen ratio can be attained while maintaining a low local concentration of oxygen, eliminating the possibility of an explosion. The objective of this study was to determine the conversion and selectivity of the methane partial oxidation reaction when using high feed rates and low methane/oxygen feed ratios in the membrane reactor configuration. A Rh supported catalyst was chosen because Rh has been shown to be one of the most active and selective catalysts for methane partial oxidation [6-8]. A 3% RhffiO2 was the most active catalyst, which ignited at 320~ in a fixed bed microreactor when using methane and oxygen feed rates of 500 and 250 cc/min respectively. It yielded a methane conversion of--70% and a CO selectivity of 85%[10]. It was also found that 100% oxygen conversion is achieved in all cases and that the ignition temperature could be even lower for lower methane/oxygen feed ratios. Experiments were performed initially in the fixed bed reactor so that results obtained in the membrane reactor could be compared to those obtained in the fixed bed reactor. EXPERIMENTAL

Catalyst Preparation. The 3% Rh/TiO2 catalyst was prepared by the wet impregnation technique. The required amount of RhC13.2H20 (Sigma Chemical) was dissolved in deionized water. Then after a few minutes the TiO2 (mostly anatase) was

694 added. The solution was stirred for approximately 30 minutes, then slowly heated so that all of the water evaporated. The resulting solid was then calcined in oxygen for 2 h at 600~ XRD, XPS, and surface area measurements were then performed to characterize the fresh and used catalysts. Procedure. Before ignition, the catalyst is externally heated to the expected ignition temperature under the flow of methane corresponding to the experiment. ,, . -,4-02 Upon reaching the expected ignition temperature, the reaction is ignited by II Ili~, ~e~mocouple initializing the flow of oxygen. Ignition occurs almost simultaneously with the CH4.--~-~ introduction of oxygen. Therefore, for safety reasons, the oxygen concentration is Glazed Portion maintained at a low value until the temperature stabilizes. After ignition, the Reactor Wall 9 external heater is turned off as the heat (quartz) Tube generated by the reaction is enough to sustain the reaction under autothermal behavior. Analysis. The effluent concentration Portion was analyzed by gas chromatography (GC) Catalyst Bed using a 5 m. 1/8" molecular sieve column to determine CH4, 02, H2, and CO. CO2 was determined using and Infrared analyzer. Water was condensed in an ice trap and further removed by using a drierite trap. Conversions were calculated as the difference between the outlet and inlet molar flow rates of methane or oxygen, divided by their inlet Fig. 1. Schematic diagram flow rates. The CO and H2 selectivities were calculated as the molar flow rates of CO and of the.membrane reactor. H2 in the effluent divided by the total amount of carbon oxides and H2 and H20 in the effluent respectively. Water flow rates were calculated by mass balance.

[::J l

RESULTS

AND

DISCUSSION

In all experiments with the 3% Rh/TiO2 catalyst the reaction ignited at 340~ only 20~ higher than the fixed bed experiments. As with the fixed bed results, 100% oxygen conversion was achieved in every experiment and there was no evidence of hydrocarbon production. Upon ignition, the temperature of the reactor increased until it reached a steady state value that depended on operating conditions. The first variable studied was the effect of catalyst loading. In these experiments methane and oxygen feed rates were maintained at 500 and 250 cc/min respectively. The results (Figure 2) show that methane conversion remains fairly constant around 60% and decreases only slightly even when the amount of catalyst was almost doubled. Oxygen conversion remains constant at 100%. CO selectivity also remains nearly constant at around 80%, being highest reaching 90% when using 60 mg of catalyst. The hydrogen selectivity shows the most dependence on catalyst loading. Hydrogen selectivity reaches a maximum of 82% using 60 mg and decreases slightly as loading increases or decreases. The optimal yield in the membrane reactor occurs when the catalyst loading is just large enough to

695

100

r

=

0

0

W

.

9

.

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o~ 75 C

o

=l,m

100 -

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-!

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=~t 5 0

> r

o

o

-250

25

0

40

0 Methane Conversion

I

50

9 Oxygen Conversion Temperature I I I J 0 60 70 80 90 100

Catalyst Loading (mg)

v

25-

0 CO Selectivity H2 Selectivity

O~ 40

I

50

I

60

I

70

I

80

I

90

Catalyst Loading (mg)

Figure 2. Effect of catalyst loading completely cover the permeable section of the membrane tube. The minimum amount of catalyst that can be used in the membrane reactor is 50 mg because this amount is required to completely surround the permeable portion of the membrane. Conversion is highest at loadings which are just over this minimum loading. At the higher loadings, methane conversion and hydrogen selectivity both decrease, which is an indication of oxygen reacting with the hydrogen rather than methane. It is clear that oxygen is the limiting factor in methane conversion. Increasing the catalyst loading exposes only a small fraction of the additional catalyst to the methane/oxygen mixture because the volume of catalyst is higher than the permeable section of the membrane tube. Simulation results in a fixed bed reactor (not shown) predicts that the reaction occurs only in a small region at the bed entrance. In the case of the membrane reactor, the active region is located near the wall of the membrane tube. Accordingly, increasing the catalyst loading does not significantly increase the active region as shown by the experimental observations. The catalyst temperature, like methane conversion and CO selectivity, does not vary too much in these experiments. It remains constant near 515~ throughout regardless of the catalyst loading. At the higher loadings, methane conversion decreases so it would be expected that the temperature would decrease. The results show that the temperature does not change at the point at which it is measured. However, it is expected that the temperature gradient in the reactor is affected by the increasing amount of catalyst. The membrane reactor yields slightly lower conversions and selectivities than the fixed bed reactor. Nonetheless, the safer reaction environment that the membrane reactor provides, permits us to study the reaction at conditions that could be hazardous in a fixed bed reactor. In the membrane reactor, it is possible to lower the methane/oxygen feed ratio without the potential for an explosion. Conversions and selectivities obtained when varying the methane/oxygen feed ratio are shown in Figure 3. In these experiments 60 mg of catalyst

100

696

02

are used, and the methane feed rate is held constant at 500 cc/min while varying the flow rate.

900 100 ~

10s

600 ~ "r, o.75'

75-

.>,

0

I

400 = ~50

5025-

It 1 .~3

Methane Conversion Oxygen Conversion Temperature ~

2.~

~

3.

Methane/Oxygen Feed Ratio

200

25 0 I

1

1.5

H2 Selectivity I

2

I

2.5

I

3

3.6

Methane/Oxygen Feed Ratio

Figure 3. Effect of methane/oxygen feed ratio Although oxygen conversion is complete at 100%, methane conversion, syngas selectivity and temperature are affected by changes in the methane/oxygen feed ratio. Methane conversion varies from a high value of 64% at a ratio of 1/1 to a low value of 44% at a ratio of 3/1. Selectivities vary from 22 to 82% for hydrogen and 67% to 90% for CO. Both CO and hydrogen selectivities are significantly lower at the lower feed ratios, particularly when the ratio is less than 2/1. This is an indication that the complete combustion reaction begins to occur. This is corroborated by the large amount of water condensed in the traps and the larger amount of CO2 detected in product gas analysis. This effect can be explained by the fact that at the lower feed ratios, more oxygen is available in the reactor to oxidize CO and H2 to form CO2 and water. At higher feed ratios, the concentration of oxygen is low and hence methane reacts to form CO instead of CO2. It should be noted however that the CO and hydrogen selectivities remain nearly constant around 90 and 80% ,respectively, for values of CH4/O2 above 2/1. The decrease in temperature with increasing methane/oxygen ratio also is consistent with the role of the oxidation of the products as the reason for the lower selectivities. At feed ratios less than 2/1, the steady state temperature rises above 700~ but at ratios equal to or greater than 2/1 the temperature is in the 500~ range. The higher temperature can be accounted for by the reaction of hydrogen and oxygen to form water and the occurrence of the complete combustion reaction, which has a higher heat of reaction than the partial oxidation reaction. In addition, as methane conversion decreases, the heat generated decreases and the temperature decreases.

697 Although 100% conversion of oxygen is achieved at low feed ratios, the methane conversion levels off at about 65%. While the residence time is higher at the low C H j O 2 ratio, we expected higher methane conversion since oxygen is no longer the limiting reagent in this range of feed ratios. Additional oxygen should result in more methane being converted, but instead it only results in more of the products being converted to CO2 and water. At higher feed ratios, the methane conversion begins to decrease because the amount of oxygen in the reactor decreases. In addition, at higher feed ratios the temperature is lower, further decreasing methane conversion. At a 2/1 feed ratio, which is the stoichiometric ratio of the direct partial oxidation reaction, methane conversion and both CO and hydrogen selectivities are maximized. The effect of varying the total feed rate was studied using 60 mg of catalyst and a methane/oxygen feed ratio of 2/1. The total feed rate varies from 300 to 1200 cc/min corresponding to residence times of about 3 to 0.8 ms respectively. The results ( Figure 4) show that, contrary to expectations, upon doubling the flow rate from 300 to 600 cc/min, methane conversion and H2 and CO selectivities remain nearly constant at approximately 65%, 70%, and 80% respectively. In addition, as in the previous results, oxygen conversion remains at 100%. Methane conversion, which is nearly constant at 65% at feed rates of 600 cc/min and below, falls to only 35% at 1080 cc/min. The CO and H2

IOE

9 W

A

9

A W

;

9 9

9 9

750

100

!

-.I

o~75

75

C 0

,m

C

P

0

~50 "~ u}

C

0

o

250"-"

o

25

0 MethaneConversion I OxygenConversion 0 Temperature

0 200 4()0

I

600

25

CO Selectivity H2 Selectivity

0

800 1000 1200

0

i

200

Total Feed Rate (cc/min)

400

I

600

I

I

800 1000 1200

Total Feed Rate (cc/min)

Figure 4. Effect of Total Feed Rate selectivities remain nearly constant throughout the whole range of flow rates, even when methane conversion decreases. The temperature increases at the higher feed rates, even when a decrease in methane conversion occurs. The temperature is nearly constant, around 515~ at the lower feed rates, but increases to 680~ at 1080 cc/min. This is because, although methane conversion is decreasing, overall syngas production is increasing due to the larger total flow rate. Thus, the temperature increase results from an increase in heat generation due to higher rate of syngas production. It should be noted that higher feed rates were attempted, but in such

698 cases the temperature rose out of control due to the large amount of heat evolved from an increased amount of syngas production. The effect of flow rate observed in a fixed bed reactor is somewhat different than what is achieved in the membrane reactor. In the fixed bed reactor, methane conversion remains nearly constant at higher flow rates, but in the membrane reactor, the conversion decreases significantly. This is probably due to the fact that in the fixed bed reactor, the reaction occurs in a narrow region at the top of the bed, but in the membrane reactor, the reaction occurs close to the membrane wall. In the membrane reactor, as the total flow rate increases, the amount of oxygen near the outside reactor wall does not increase, but the amount of methane throughout the bed does increase. In addition, there is a greater area for mixing in the fixed bed reactor, which becomes more important to maintaining high conversions at high flow rates. It is this greater reaction area which also results in lower hydrogen selectivities in the fixed bed. In that case, there is a greater probability that hydrogen will react with oxygen to form water. A detailed elementary step model similar to that proposed by Hickman and Schmidt [7] was developed to interpret the fixed bed results. The model assumes that methane adsorbs dissociatively forming adsorbed carbon and adsorbed hydrogen. Oxygen adsorbs dissociatively and the reaction of adsorbed carbon and adsorbed oxygen yields CO. Recombination of adsorbed hydrogen yields dihydrogen which desorbs into the gas phase. Further oxidation of CO yields CO2 and the reaction of adsorbed hydrogen and adsorbed oxygen yields adsorbed hydroxyls which recombine or react with adsorbed hydrogen to yield adsorbed water. Using this model we were able to reproduce the results reported by Schmidt and Hickman using their reaction parameters. After accounting for the differences between the monolith and fixed bed reactors we attempted to fit the model to our fixed bed results using the same reaction parameters as Hickman and Schmidt [7], we found a significant discrepancy between the model predictions and our experimental fixed bed results. The discrepancy derived mainly from the fact that the reaction temperatures in our experiment are significantly lower than in the those reported by Schmidt group using a monolith reactor. The model predicted that under our reaction conditions the surface will be covered mainly by CO and under predicted the methane conversion (30% instead of 70%) and CO selectivity (20% instead of 90%). Upon conducting a sensitivity analysis we found that the model prediction fitted our results fairly well when the CO desorption energy was lowered from the value of 31.6 kcal/mol used by Hickman and Schmidt to 25.5 kcal/mol. Hence we conclude that in our case the TiO2 support increases the rate of CO desorption. One of the main results of the model is that it showed that the reaction occurred in a narrow region near the entrance of the bed, which explains the lack of sensitivity of the results to the changes in loading, flow rate and methane/oxygen ratio. The fixed bed reactor model is one dimensional and thus its solution is not so involved. In the case of the membrane reactor it is necessary to use a two dimensional model to account for the radial flow of oxygen at the wall. Work is underway to solve this model, however, the results of the one dimensional model are qualitatively useful to interpret the membrane reactor results. After studying the effect of catalyst loading, methane/oxygen feed ratio, and total feed rate, the effect of operating the reactor at temperatures above the autothermal steady state temperatures was studied. Because only 65% methane conversion was achieved in the membrane reactor, an attempt was made to increase the conversion by raising the reaction temperature using external heating. In these experiments, 500 cc/min of methane and 250 cc/min of oxygen were reacted over 60 mg of catalyst. As the temperature is increased to 650~ the oxygen conversion remains at 100%, methane conversion remains around 60% and the CO and hydrogen selectivities remain around 90 and 75% respectively. Due to 100% oxygen conversion, there is not enough oxygen throughout the catalyst bed to react with the methane and no further methane conversion is achieved. Oxygen is consumed near the membrane wall and does not reach the outer, external wall of the reactor, resulting in

699 unconverted methane. An alternative method of increasing the methane conversion is for the unreacted methane to participate in an additional reaction. A potential candidate is the reaction between methane and CO2 (dry reforming). So a specified amount of CO was added to the methane feed. It was found that the methane conversion and syngas selectivity decreased with the addition of CO2. Methane conversion decreased from 64% to 41% at a carbon dioxide feed rate of 145 cc/min, and CO and hydrogen selectivities decreased from 90% and 82% to 37% and 62% under these conditions. The decrease occurs because the added gas dilutes the methane concentration in the feed. This smaller CH4 concentration allows for a larger amount of CO2 to be produced because the local methane/oxygen concentration is relatively low. In addition, the reaction temperature remained around 510~ which is below the temperature at which the dry reforming reaction occurs. The fixed bed reactor can achieve CO and hydrogen selectivities of 93% with methane conversions around 70%. In the membrane reactor the highest methane conversion is about 65% with CO and hydrogen selectivities of 90% and 82% respectively. One of the reasons for this difference lies in the different concentration profiles in the membrane and fixed bed reactors. In the fixed bed, methane and oxygen are mixed throughout the reactor at the top of the catalyst bed, but in the membrane reactor, the oxygen concentration is much greater near the membrane tube than towards the outer wall of the reactor. The concentration profile determines the rate of reaction and thus the rate of heat evolution and consequently the temperature distribution in the bed, which is also an important factor in determining conversion. Further increases in methane conversion were attained by using an additional bed downstream from the membrane bed. In addition, the reactor temperature was increased so that the second bed operates at temperatures higher than the autothermal operation. This allows for the dry reforming reaction to occur in the second bed thus increasing the conversion of the methane not consumed in the first bed. In this case the highest methane conversion was about 90% with CO and H2 selectivities of about 90% when the external temperature is 700~ Similar results can be attained without heating if a third feed of 02 is added between the membrane bed and just before the second fixed bed. In this case, the temperature increase is realized by the partial oxidation reaction with no major loss of selectivity. At this point only a few catalysts have been studied in the membrane reactor. XRD and XPS analyses have been conducted mainly for the 3% Rh/TiO2 catalyst. XRD results show that after impregnation and calcination the TiO2 phase is mainly anatase. Some small peaks corresponding to metallic Rh are also detected. The XRD pattern for the catalyst after reaction shows that the TiO2 phase has changed from mainly anatase to rutile. It should be noted that the catalyst did not deactivate during the duration of these experiments. Furthermore, the results were reproducible when tested several times using the same batch of catalysts. So the phase change of the support does not seem to affect the activity of the Rh catalyst. XPS results indicate that the Rh 3d peak is located at 309.1 eV, while the reported binding energy for metallic Rh is 307.2 eV. After reaction, there is a significant change in binding energy to 306.7 eV. These results indicate that there is a change in the oxidation state of Rh during reaction. The reported binding energies of Rh203 and RhC1 are 308.8 and 310.1 eV respectively. The specific state of Rh cannot be determined by these ex-situ experiments and work is underway to establish the state of the catalyst surface after reaction. The changes occurring during reaction do not appear to significantly affect the catalyst activity and selectivity, at least during the duration of these experiments. BET surface areas for the calcined but unreacted 3%Rh~iO2 was 117.5 m2/g whereas the value for a 0.3%Rh/TiO2 catalyst was 43.7 m2/g. These results, obtained in a Quantachrome apparatus, were confirmed by several measurements. It is not clear, however, why the surface area of the catalyst with higher Rh loading is so high.

700 Hydrogen chemisorption results, obtained by the pulse method, indicated that the dispersion of Rh on the calcined 3%RhffiO2 catalyst was 56.8 % which corresponded to an average crystallite size of 1.9 nm. It should be noted that this value is rather low, probably because some of the surface was in an oxidized state as shown by the XPS results. In conclusion, the catalytic partial oxidation of methane on a RhffiO2 catalyst has been demonstrated to occur at high conversions and selectivities under fast flow conditions when a membrane is used to separate the methane and oxygen feeds. This configuration allows operation of the reactor at low methane/oxygen ratios and high feed rates while eliminating the possibility of a flame flashback leading to an explosion. The low methane/oxygen ratio and the high flow rates are the key factors to attain autothermal behavior. The RhffiO2 catalyst exhibited a low ignition temperature and did not exhibit deactivation during the duration of these experiments. REFERENCES

1. Prettre, M. Eichner, C. and M. Perrin, Trans. Fraday Soc., 43 (1946) 335. 2. Gavalas, G. R.Phichticul, C. and G. E. Voecks, J. Catal., 88 (1984) 54. 3. Blanks, R. E., Witrig, T. S.and D. A. Peterson, Chem. Eng. Sci., 45 (1990) 2407. 4. Vermeiren, W. J. M., Blomsma, E. and P. A. Jacobs, Catal. Today, 13 (1992) 427. 5. Ashcroft, A. T. Cheetham, A. K, Foord, J. S., Green, M. L. H., Grey, C. P., Murrell, A.J. and P. D. F. Vernon, Nature, 344 (1990) 319. 6. Hickman, D.A. and L.D. Schmidt, J. Catal 138 (1992) 267. 7. Hickman, L.D. and L.D,. Schimidt, AICHE J. 39 (1993) 1164. 8. Torniamen,P.M., Chu,X., and L.D. Schmidt, J. Catal, 146 (1994) 1. 9. Santamaria, J.M., Miro, E.E., and E.E. Wolf. 10. Shiraha, T., M.S. Thesis, Unversity of Notre Dame (1995).

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

701

Sustainable N i / B a T i O 3 Catalysts for Partial O x i d a t i o n of M e t h a n e to Synthesis Gas R. Shiozaki, a A. G. Andersen, b T. Hayakawa, S. Hamakawa, K. Suzuki, M. Shimizu and K. Takehira* National Institute of Materials and Chemical Research, Tsukuba Research Center, AIST, Higashi 1-1, Tsukuba, Ibaraki 305, Japan. aChemical Research Department, Institute of Research and Innovation, Takada 1201, Kashiwa, Chiba 277, Japan. bNycomed Imaging, P.o.box 4220 Torshov, 0401 Oslo, Norway. Ni/BaTiO3 catalyst has been prepared by solid phase crystallization (SPC) method and used successfully for partial oxidation of CH4 into synthesis gas at 800~

The catalyst was

further tested for 75 hrs with no observable degradation and negligible coke formation. The

SPC method bestows the catalyst with high Ni dispersion over the perovskite as well as strong metal-support interaction between Ni and the perovskite, resulting in both high activity and high sustainability against coke formation.

I. I N T R O D U C T I O N Recently, intensive studies have been carded out on the catalytic partial oxidation of CH4 to synthesis gas [1-3]. This process has advantages over the conventional steam reforming of CH4 to make synthesis gas, as the latter process is highly endothermic and produces synthesis gas having a H2/CO ratio > 3. The partial oxidation of CH4, expected to afford synthesis gas having H2/CO ratio of about 2, makes methanol synthesis an ideal follow-up process. In the partial oxidation of CH4 to synthesis gas, coke formation over the catalyst frequently takes place, resulting in catalyst deactivation. Claridge et al. [4] observed that the relative rate of coke formation follows the order Ni>Pd>> Rh, Ru, Pt, Ir. Nickel catalysts are highly effective for partial oxidation of CH4 to synthesis gas, but they are unsatisfactory with respect to coke formation [5]. From the industrial view point, Ni is preferable as the active species compared to expensive precious metal such as Rh, Pd or Ru. High dispersion of metal species over catalyst may reduce coke formation [6]. Nickel-supported catalysts are conventionally

702 prepared by wet impregnation of different supports. This method is not fully reproducible and may give rise to some unhomogeneity in the distribution of the metal on the surface. Therefore, a new method of catalyst preparation able to produce homogeneous distribution of nickel is proposed, i.e., use of the precursors containing homogeneously distributed nickel inside in the structure, followed by further calcination and reduction, may result in the formation of well dispersed and stable metal particles. This method gives a higher reproducibility and easier characterization in comarison to the analogous samples prepared by wet impregnation. Here, we named this method "solid phase crystallization (SPC) method". We have reported sustainability and high activity of Ni/(Ca, Sr)TiO3 in the oxidation of CH4 to synthesis gas [7-10], where the Ni catalyst supported on (Ca,Sr)TiO3 perovskite was prepared by the SPC method. However, the precursor was not homogeneous and contained two types of nickel species; one was dissolved in the Ti site in (Ca, Sr)TiO3 and another was separated as NiO from the perovskite structure. Nickel dissolved in the perovskite structure was reduced to well dispersed and stable metal species, while NiO was reduced to large size of Ni metal particles after the reduction. The well dispersed and stable nickel particles seem to be important for operation at low coke formation over the catalyst during the reaction. BaTiO3 can contain higher amount of nickel in the Ti site compared to CaTiO3 and give rise to more completely dispersed metal on the surface during the SPC preparation. Here, we report the preparation and thermal evolution of new perovskite containing well dispersed and stable nickel on the surface, resulting in the sustainable activity for the partial oxidation of methane.

2. EXPERIMENTAL

Catalysts of the composition BaTi 1-xNixO3-5, 0'~_x<0.4, were prepared according to the Pechini patent [ 11]. Ethylene glycol, cytric acid, tetraisopropyl orthotitanate, barium carbonate and nickel nitrate were used as starting materials [9,10]. Moreover, several model catalysts with related stoichiometfic Ba-Ti-O composition and combined with nickel oxide, i.e., Ba2Ti O4.0.3NIO, BaTiO3o0.3NiO, BaTisO1 l O0.3NiO and TiO2.0.3NiO, were also prepared in the same way. The solution obtained from the starting materials was evaporated at 80-90~ to form a viscous liquid, followed by two step decomposition by heating at 200~ for 5 hr and 500~ for 5 hr, and finally calcination at 850~ in air for 10 hr. Only TiO2~

was finally

calcined at 955~ for 1 hr to ensure the formation of rutile crystal structure [12]. Cooling was slowly performed in air by turning off the furnace. The powders produced by the above method were tested in a U-shaped fixed bed reactor in a mixture of CH4 (1.0 loh-1) and air (2.5 l.h-1). The catalyst bed contained 300 mg of catalyst

703 material in 2 ml of quartz wool to avoid sintering and clogging of the reactor. The catalytic experiments were started by introducing CH4 and air at room temperature, and carded out by increasing the reaction temperature from room temperature to 800~ at the rate of 2.5~

1.

If not specifically mentioned in the text, gas samples for analysis of products were taken after 30 minutes equilibration at each temperature. Moreover, the reactions for testing the catalyst life were carried out at 800~ for 75 hr under the same conditions. Product gases from catalytic testing were analysed by a gas chromatograph of Shimadzu GC-8A equipped with TCD detector. The selectivities to C2 compounds, CO2, CO and H2 were calculated based on the numbers of carbon and hydrogen atom in CH4. The catalysts were characterised by XRD (MXP-18: MAC Science Co.), SEM (Hitachi S 800), TEM (JEM-2000FX: JEOL with EDS: Northern), TGA/DTA (Shimadzu DTA 50 and TGA 50), IR (JASCO VF/IR 7000) and XPS (PHI-5000 with Mg Ktx radiation). Surface area of the catalyst was measured with a Micromeritics model 2200.

3. R E S U L T S

3.1 BaTil.xNixO3.8 catalysts XRD analyses of the BaTil_xNixO3-8 catalysts (x= 0, 0.01, 0.05, 0.1, 0.2, 0.3 and 0.4) showed the formation of BaTiO3 in all the samples. The first weak sign of excess NiO is seen for x--0.2, and is clear for higher nickel concentrations. The absence of NiO peaks at lower concentrations may be due to solubility of Ni in the perovskite, as well as too small crystallites to give diffraction signal. Reflections from Ba2TiO4 can be seen increasing with the nickel concentration from x--0.2. The presence of Ba2TiO4 is not surprising as there will be an excess of Ba on A-sites beyond the point of saturation of Ni on B-sites. After the catalytic testing, the XRD patterns still show BaTiO3 to be the main constituent of the catalyst. In these diffractograms we see neither NiO nor nickel metal, except small traces of NiO at the highest nickel concentration (x=0.4). The second phase in the diffractograms is BaCO3, which is seen increasing with the original nickel content of the material. Ba2TiO4 disappeared after the catalytic testing. NiO which was visible in diffractograms after calcination was negligible after the catalytic testing. We believe substantial part of nickel on the surface has been reduced to its metallic form during the catalytic testing. The nickel metal particles are probably too small to give reasonable signals in XRD. BaTil_xNixO3_8 (x = 0, 0.01, 0.05, 0.1, 0.2, 0.3 and 0.4) were tested for the CH4 oxidation. The 02 conversion was high for all experiments at 750 and 800~

The CH4

conversion increased with temperature and with nickel content. For low nickel content and low

704 temperature the catalysts showed a certain selectivity (=30%) for C2-products (ethane and ethene). However, the main product under the same conditions was CO2. At high nickel content and 800~ the catalysts selectively produced synthesis gas. Visual observation, carbon mass balance of the gaseous products, as well as TGA of used catalysts revealed no or very low (< 0.1%) coke formation during the catalytic testing. Total carbon (xC: CO +CO2+CH4 unreacted) in the effluent gas was 40.3 mmol~ -1 which well coincided with the value calculated from the CH4 gas flow 1.0 loh-1 controlled at 30~

3.2 Ba/Ti/Ni compounds By controlling the ratio of Ba/Ti in the citrate method mentioned above, Ba2TiO4, BaTiO3 and BaTi5011 were obtained as a single phase in each composition. The three mixed oxides and TiO2 were combined with nickel at the ratio of 1/0.3 by the citrate method and tested for the CH4 oxidation. A summary of the observations in the XRD patterns for the related compounds before and after the catalytic testing is given in Table 1. BaTiO3o0.3 NiO shows BaTiO3 and traces of NiO Table 1 Surface area and XRD data of Ba-Ti-Ni mixed metal oxide catalysts. Catalyst

Surface Areaa) m2~ -1

Ba2TiO4~

Observed Phases before testing

after testing

2.3

Ba2TiO4 + NiO

BaTiO3 + BaCO3

BaTiO3.0.3NiO

8.9

BaTiO3 + NiOb)

BaTiO3

BaTi5011~

7.1

TiO2~

0.6

BaTi5011 BaTi5011 NiTiO3(illumenite)+TiO2(rutile) TiO2(ruffle)+Ni

a) fresh catalyst. b) Traces. of NiO peaks were observed. before the catalytic testing, while neither NiO nor Ni metal is observed after the testing. In the case of Ba2TiO4o0.3NiO, NiO can be weakly seen before the testing and dissappeared after the testing. When the catalyst was used for the catalytic testing for 60 min at 800~

BaTiO3 was

the main compound together with BaCO3; still, no nickel was observed. During the testing, Ba2TiO4 was decomposed into BaTiO3 and BaCO3 in the atmosphere containing CO2. BaTi5 O11~

showed the peaks of BaTisO11, and neither nickel species nor substantial change

705 was observed in the XRD patterns before and after the testing. Use of TiO2o0.3NiO resulted in the formations of NiTiO3 (illumenite) and TiO2 (rutile) after one hour calcination at 955~ After the catalytic testing, XRD showed only ruffle and metallic nickel. In this case, Ni metal was clearly observed as in the cases of Ni/(Ca, Sr)TiO3 and Ni/T-AI20 3 [9-10]. IR spectroscopy studies confirmed a strong C=O stretching band at 1431 cm -1 on the Ba2TiO4o0.3NiO after the testing due to BaCO3 formation. The same band, much weaker, was observed on all catalysts before and after the testing, probably due to surface carbonate groups attached to barium in the perovskite. From the SEM observations before the testing, BaTiO3o0.3NiO is composed of small particles (<0.1 ktm), TiO2~

is of sintered agglomerate of small crystals (0.2-0.3 ~tm) of

rutile, and plate-like crystals are formed in Ba2TiO4~

After the testing, BaTiO3*

0.3NiO showed no obvious change in the morphology, whereas the surface morphology of Ba2TiO4o0.3NiO was drastically changed and the small particles clearly appeared. Plate-like crystal before the testing was completely changed into the porous structure having many crevices after the testing by the phase transfer from Ba2TiO4 to BaTiO3 and BaCO3. In TiO2o0.3NiO after the testing, small particles (<0.1 ~tm) probably of Ni metal were observed on rutile TiO2 surface suggesting the decomposition of NiTiO3 to Ni metal and TiO2. The present process is thought to involve first the oxidation of part of CH4 to H20 and CO2, followed by the reforming reaction of CH4 with H20 and CO2 [13]. Therefore, the catalyst bed may be composed of two zones, the first zone is composed of NiO over the perovskite or Ni in the perovskite under the oxygen-rich atmosphere and the second is of Ni metal under the reducing atmosphere. TGA was performed on catalysts before and after the testing in flowing air (20 mlomin-1). Around 400~ a weight increase probably due to oxidation of Ni metal is observed on used catalysts, while fresh catalysts show flat curves in the whole temperature range. The weight increase of used catalysts of both BaTiO3o0.3NiO and Ba2TiO4~

amounts to about 0.7% of the total mass of the used sample. Assuming

oxidation of Ni metal to NiO being the only reaction taking place, this corresponds to 10% of all the nickel being in the metallic form after the testing. The sample of used TiO2o0.3NiO shows a weight increase in TGA of 4.75%, corresponding to 60% of all the Ni being in the reduced form. The difference between the TGA behavior of the rutile sample and the others is also reflected in the XRD results. Nickel probably forms larger crystals on rutile than on the barium compounds. TGA of used sample of Ba2TiO4o0.3NiO shows an additional feature at temperatures above 800~

namely a weight loss of 8%. XRD of this catalyst after the testing

revealed large amounts of BaCO3, and we assume the weight loss is corresponding to the thermal decompostion of BaCO3, releasing CO2 to the air. 8% weight loss thus indicates that

706 Table 2 CH4 oxidation with air over a series of Ba-Ti-Ni mixed metal oxide catalysts. Catalyst

Temperature Conversion / %

Ba2TiO4~

BaTiO3o0.3NiOb) BaTi5Oll~ TiO2~

Selectivity / %

/~

CH4

02

C2

CO2

CO

H2

750

41.0

94.9

35.6

62.7

1.6

2.9

800

38.0

94.9

22.2

70.0

7.8

10.4

800a)

65.8

94.9

0.2

24.5

75.3

75.8

750

94.8

94.9

n.d.

7.5

92.5

85.3

800

98.0

94.9

n.d.

5.6

94.4

91.9

750

30.2

88.3

10.2

83.4

6.4

10.8

800

32.7

93.8

11.3

82.7

6.0

8.2

750

94.7

94.9

n.d.

7.8

92.2

91.4

800

98.1

95.0

n.d.

5.6

94.3

92.4

a) sampled after 60 min. b) C2 products were observed up to 650~ 35% of the catalyst after the testing is composed of BaCO3. Thus, it is likely that Ba2TiO4oNiO was quantitatively converted into BaTiO3~

and BaCO3 during the catalytic testing.

Catalytic data for the Ba/Ti/Ni compounds used in the CH4 oxidation are summarized in Table 2. BaTiO3o0.3NiO showed high conversions of both 02 and CH4, low selectivity for CO2 and high selectivity for synthesis gas at high temperatures. BaTi501 l~

showed a

high yield of CO2 as well as small amount of C2 products at all temperatures. Ba2TiO4~ gave CO2 as the main product, but showed a relatively high selectivity for C2 products. During the testing at 800~ for 60 min, the selectivity gradually changed with time and the catalyst substantially turned its characteristics into producing synthesis gas, but not as efficiently as BaTiO3o0.3NiO.

3.3 Long term testing of BaTiOa.0.3NiO The best performing catalyst, BaTiO3-0.3NiO was subject to a continuous reaction test and showed a constant activity during 75 hours at 800~ (Fig. 1). ~:C showed almost constant value close to 40.3 mmol.h -1. After 75 hours of testing, the reactor was filled with nitrogen and cooled according to normal procedures. Finally, a temperature programmed oxidation (TPO) experiment was performed by heating the reactor from room temperature to 950~ at a

707

1 O0

>,

98 Conv.

>

O9

Oo Conv.

96

"0 ttO

-

..--

j:?

C'X,,,,~_Select.

~43-13.o

L_

>

~ H 2 Select.

0

0

90

I

o

20

,

I

,

40

I

,

60

80

Reaction Time / h

Figure 1 Conversion and selectivity as a function of time for the long term test on BaTiOa.0.3NiO. rate of 2.5~

-1, with an air flow of 41 ml.min -1. Off-gases were analysed and the rate of

CO2 formation is plotted. No significant amount of CO2 was observed during this test with BaTiO3-0.3 NiO; CO2 formation observed around 450~ may be due to coke formed over the catalyst which corresponded to 0.02 wt% carbon on the catalyst. In the CH4 oxidation over TiO2o0.3NiO, the catalyst behaved initially like BaTiO3.0.3NiO. After extended time, however, the CH4 conversion decreased, as well as the selectivities to H2 and CO, and in contrary CO2 increased with time. EC showed no significant change during the reaction. After 75 hours of testing an amount of CO2 corresponding to 0.41 wt% of the catalyst being coked. TiO2o0.3 NiO was thus enough sustainable against coking, and therefore the other deactivation mechanism, i.e., sintering or oxidation of Ni metal, can be considered. The CH4 oxidation over the other conventional Ni/q,-A1203 or Ni/o~-Al203 catalyst under the same conditions as above have been tested. After 150 hr of the reaction, 29.3 and 16.8 wt% of coke were observed over Ni/y-A1203 and Ni/t~-Al203 catalyst, respectively [9,10].

4. D I S C U S S I O N The CH4 oxidation over BaTiO3o0.3NiO catalyst showed variable dominating reaction mechanisms at various temperatures. Deep oxidation to CO2 occurs together with oxidative CH4 coupling below 700~

while the reaction is completely dominated by synthesis gas

708 Table 3 Surface metal composition of BaTiO3.0.3NiO before and after the catalytic test determined by XPS. Catalyst

BaTiO3"0.3NiObefore

Content / % Ba

Ti

Ni

40.2

45.3

14.5

BaTiO3.0.3NiOafter

39.5

46.3

14.2

Calculated

43.5

43.5

13.0

production around 800~

Synthesis gas production is probably related to the presence of

metallic nickel on the surface of the catalyst. By XRD, trace NiO was observed before the catalytic testing, but no Ni metal was observed after the testing. However, oxidation of the used catalysts (BaTiO3.0.3NiO and Ba2TiO4.0.3NiO) in TGA gave a strong indication of the presence of metallic nickel as shown in 3.2. The presence of Ni on the surface of catalyst is also evident from XPS measurement of BaTiO3o0.3NiO before and after the catalytic testing. The compositions of metals were calculated from the spectra by using the method reported by Penn[ 14] and are shown in Table 3. The amount of Ni was obtained by summing Ni 0, Ni 2+ and Ni 3+, since surface of Ni metal over the catalyst was possibly oxidized by exposing the samples after the testing. BaTiO3~

after the testing still shows the same value of Ni

surface amount as before the testing. Thus, both TGA and XPS measurements show strong evidence of Ni metal on the catalyst surface, and moreover no observation of Ni metal in XRD suggests the formation of fine Ni metal particles with high dispersion. In the present study, XRD measurements simply indicate that the solubility of Ni is 1020% on B-sites of BaTiO3. The solubility will probably decrease by lowering the temperature and the partial pressure of oxygen. In this respect, Ni may precipitate to the surface during the use of the catalyst under reducing atmospheres. The precipitated Ni will yield a high dispersion of Ni on the suface of the perovskite crystallites, and contribute to the formation of an active and sustainable catalyst. TEM observation of BaTiO3~

after the testing actually shows

high dispersion of Ni species (Fig. 2). Ni is apparently present at the located area as seen as dark spots (A and B in Fig. 2a), where EDS analyses show the coexistence of Ba and Ti together with Ni. At the site C, Ni is not observed, but Ba and Ti are detected by EDS. TEM image of the dark spot at high magnifications (Fig. 2b) suggests that the spot is composed of the agglomerate of fine Ni particles (diameter of particle < lnm). TEM image of TiO2~

709

a Figure 2. TEM images of BaTiO3~

b

c

(a and b) and TiO2o0.3NiO (c) after the catalytic

testing. after the testing as a comparison (Fig. 2c) clearly shows the formation of large size of Ni metal particles as black spots (A and B), which are mainly composed of Ni by EDS analyses. Thus, Ni is quite well separated from TiO2 as the support on TiO2o0.3NiO after the testing, while BaTiO3 perovskite can strongly hold fine Ni metal particles on the surface. In the latter case, nickel is obviously distributed in fine particles, and we interpret this in terms of a more favourable interaction between nickel and the support. Such phases indicate that there may be an increased "affinity" between the nickel and the support during catalytic action. When nickel is precipitated to the surface, a surplus of Ba will occur on A-sites of the perovskite material. Barium is one of the elements that are known to keep the level of coking on metals down. In the present set of experiments, coke is not formed to any obvious extent. There are three advantages related to the present BaTiO3o0.3NiO catalyst prepared by the SPC method:

1) Precipitation from solid solution gives optimum dispersion of Ni on the surface. A main fraction of Ni has been dissolved in the lattice during calcination, and has precipitated during the catalytic testing. 2) Ni has better "affinity" to B-sites in the perovskite than to most traditional supports, like A1203 and TiO2. This results in a better dispersion and low driving force for metal sintering. 3) Low coking is traditionally achieved by promoting the catalyst with alkaline or alkaline earth metals. The rationale behind this practice is probably to kill acidic sites, which are known

710 to produce coke. Abundant barium is probably easily accessible for killing acidic sites on the surface, especially when it is released together with nickel from the bulk of the material. The main problem connected to practical use of this kind of perovskite catalyst in an operating plant is probably low surface area and the tendency to form fine particles during the catalytic action. This must be solved by using a binder or a certain support, like ct-A1203. The latter study has just been undergoing in our laboratory.

REFERENCES

1 A.T. Ashcroft, A. K. Cheetham, J. S. Foord, M. L. H. Green, C. P. Grey, A. J. Murrell and P. D. F. Vernon, Nature, 344 (1990) 319. 2

D.A. Hickman and L. D. Schmidt, Science, 259 (1993) 343.

3

V.R. Choudhary, A. M. Rajput and V. H. Rane, J. Phys. Chem., 96 (1992) 8686.

4

J.B.Claridge, M. L. H. Green, S. C. Tsang, A. P. E. York, A. T. Ashcroft and P. D. Battle, Catal. Lett., 22 (1993) 299.

5

M. Audier, A. Oberlin, M. Oberlin, M. Coulon and L.Bonnetain, Carbon, 19 (1981) 217.

6

C.H. Bartholomew, Catal. Rev.-Sci. Eng., 24 (1982) 67.

7

T. Hayakawa, A. G. Andersen, M. Shimizu, K. Suzuki and K.Takehira, Catal. Lett., 22

8

K. Takehira, T. Hayakawa, A. G. Andersen, K. Suzuki and M. Shimizu, Catal. Today, 24

(1993) 307. (1995) 237. 9

T. Hayakawa, H. Harihara, A. G. Andersen, A. P. E. York, K. Suzuki, H. Yasuda and K. Takehira, Angew. Chem., Int. Ed., 35 (1996) 192.

10 T. Hayakawa, H. Harihara, A. G. Andersen, K. Suzuki, H. Yasuda, T. Tsunoda, S. Hamakawa, A. P. E. York, Y. S. Yoon, M. Shimizu and K. Takehira, Appl. Catal., in press. 11 M. P. Pechini, US Patent 3,330,697 (1967); M. S. G. Baythoun and F. R.Sale, J. Mater. Sci., 17 (1982) 257. 12 Gmelin: Handbuch der anorg. Chem. Bd. 41 (1951) 228. 13 D. Dissanayake, M. P. Rosynek, K. C. C. Kharas and J. H. Lunsford, J. Catal., 132 (1991) 117. 14 D. R. Penn, J. Electron Spectrosc., Relat. Phenom., 9 (1976) 29.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

711

Synthesis of Early Transition Metal Carbides and their Application for the Reforming of Methane to Synthesis G a s A. P. E. York, J. B. Claridge,t C. Mhrquez-Alvarez, # A. J. Brungs, S. C. Tsang * and M. L. H. Green*

The Catalysis Centre, Inorganic Chemistry Laboratory, University of Oxford, South Parks Road, Oxford, OX1 3QR, U.K. Binary and ternary group V and VI transition metal oxides were converted to their carbides using temperature programmed reduction (TPR). When ethane was substituted for methane in the TPR, transmission electron microscopy indicated that the oxide-carbide transformation proceeded topotactically. The carbides synthesised were tested as catalysts for methane reforming with carbon dioxide (dry reforming), giving carbon monoxide and hydrogen (synthesis gas) as the main products; of the binary carbides, only molybdenum and tungsten were active and stable, displaying activities comparable to supported noble metal catalysts. No carbon deposition was observed on post-catalytic carbide samples. Molybdenum and tungsten carbides were also active for methane reforming with air (partial oxidation ) and water (steam reforming). No synergistic effects were evident in the ternary carbides, and only those with a high ratio of Mo/W to group V metal exhibited significant catalytic activity and stability.

1. INTRODUCTION The catalytic properties of transition metal carbides and nitrides have attracted much attention in recent years. Indeed, over the past 20 years, it has been established that these materials, particularly those of the group VI transition metals, exhibit catalytic properties analogous to precious metals [1-2]. Recently, with the advent of synthetic methods for the production of high specific surface area metal carbides/nitrides (Sg _ 220 m 2 g-l) [3-9], studies have been undertaken that reveal the high activity of these materials for a number of reactions, including Fischer-Tropsch synthesis [ 10], alkane isomerisation [ 11-14], hydrodenitrogenation (HDN) [15], hydrodesulphurisation (HDS) [16] and ammonia synthesis [17]. Since the group VI transition metals are abundant and relatively cheap, it has been suggested that they can replace the scarce and expensive noble metals for a number of catalytic applications. In this paper, further research has been carried out concerning the temperature programmed reduction (TPR) synthesis method pioneered by Boudart and co-workers t Currentaddress: Department of Chemistry and Biochemistry, University of South Carolina, Columbia, SC 29208, U.S.A. Current address: The Catalysis Research Centre, Department of Chemistry, University of Reading, Whiteknights, Reading, RG6 6AD, U.K. ~9 On leave from: Instituto de Cat~ilisisy Petroleoquimica, CSIC, Campus Cantoblanco, 28049 Madrid, Spain Corresponding author

712 [3,18,19]. In the original work methane was used as the gaseous carbon source. It will now be shown that using ethane as the carbon source results in the formation of metal carbides with higher surface areas than those produced by TPR with methane. Further, the use of ethane allows the carbiding reaction to proceed in a topotactic manner and provides a potential method for synthesising a variety of carbide phases. Due to the large amount of methane reserves that currently exist, the conversion of methane to more useful chemicals is a matter of great importance. One of the most widely used commercial processes involves the reforming of methane into carbon monoxide and hydrogen, which is commonly known as synthesis gas or 'syngas' [20-22]. Unfortunately, current industrial nickel based catalyst systems have a tendency to promote coke formation, which results in a loss in catalyst activity [20]. Therefore, much research has been done over the years to reduce the levels of carbon deposition, including the use of sulphur passivated nickel catalysts [23] and noble metal catalysts [24-26], although these systems have the disadvantages of low activity and high costs respectively. Initial investigations by York et al. [27] have recently demonstrated that molybdenum and tungsten carbides are highly active methane reforming catalysts. In this paper, the use of carbide catalysts for the reforming of methane has been extended to include the group V metals, niobium and tantalum, and the results obtained are compared with molybdenum and tungsten. In addition, following the recent synthesis of high surface area ternary metal carbides by Claridge et al. [28], these materials have also been tested for the dry reforming of methane with carbon dioxide, in order to ascertain whether the ternary structures provide any synergistic effects. The ternaries examined were W9Nb8047, W3Nb14044, WNb12033, Mo2Ta2Oll, Mo3Nb2Ol4 and Mo3Nb14Oa4.

2. EXPERIMENTAL 2.1. Catalyst Synthesis and Characterisation The oxide precursors were high purity commercial MOO3, WO3, Ta205 and V205 powders (Johnson-Matthey, Puratronic| 99.998%) and Nb2Os, W9NbsO47, W3Nb14044, WNb12033, Mo2Ta2Oll, Mo3Nb2Ol4 and Mo3Nb14044. Nb205 and the mixed W-Nb oxides were prepared by crushing together the appropriate amounts of WO3 and Nb205, before peUetising, and calcining the samples in a platinum crucible at 973 K for 10 days, followed by 1473 K for a further 4 days, as described by Viccary and Tilley [29]. The mixed oxides containing molybdenum were prepared by intimately mixing various combinations of MoO3 and Nb205 or Ta205 powders, and firing in a sealed silica ampoule at 1073 K for 7 days (MoxNbyO,) or 1173 K for 15 days (Mo2Ta20~l), as described by Ekstrom [30]. The composition of the mixed oxides was verified by XRD. The carbide materials were prepared by temperature programmed reduction (TPR) of the corresponding oxides, under flowing 20% v/v CH4~2 or 10% v/v C2H6/H2 (150 ml min-l), and at a heating rate of 1 K min-~ from room temperature to 900-1223 K, depending on the metal oxide and the hydrocarbon used. Normally the catalysts were prepared in situ and tested immediately. The post-synthesis high surface area carbides are readily and exothermically oxidised by air at room temperature and, therefore, passivation in flowing 1% O2/N2 for 10 hours at room temperature was carried out before exposure to the atmosphere and characterisation.

713 X-ray diffraction (XRD) patterns of the oxide and carbide powders, in the 3 ~ to 70 ~ 20 range, were acquired with a Philips PW 1729 diffractometer using Cu Ktx radiation at 40 kV and 30 mA. Samples were mounted on an aluminium plate, and the aluminium peaks were used as a reference. High resolution transmission electron microscopy (HRTEM) was carried out using a Jeo14000EX TEMSCAN electron microscope, with an accelerating voltage of 400 kV. The N2 BET isotherms were obtained using an all glass high vacuum line and catalyst surface areas calculated. Methane (Union Carbide, >99.95%), ethane (BOC, C.P. grade), hydrogen (BOC, C.P. grade), argon (BOC, C.P. grade), oxygen (BOC, C.P. grade) and carbon dioxide (BOC, C.P. grade) were used as received without further purification.

2.2. Apparatus The apparatus used in this work was a modified version of the commercial Labcon microreactor described previously [31 ]. Briefly, the reactor was built using 1/8" and 1/16" o.d. 316 stainless steel tubing and 316 stainless steel Swagelok fittings throughout. The catalyst sample was placed between two quartz wool plugs in the centre of a 4 mm i.d. silica tube, and inserted into a vertical Severn Science tube furnace. This was heated to the required reaction temperature controlled from a Eurotherm 905 temperature controller. For safety reasons, in experiments carried out at elevated pressures the silica tube was placed inside a steel tube. Inlet gas flow rates were controlled using Brooks 5850TR mass flow controllers, and the water flow was controlled using a Bronkhorst LIQUI-FLOW liquid flow controller with a helium over pressure and operated from a Bronkhorst HighTech EPA2 control module. The exit gas stream from the reactor passed through a Tescom two stage back-pressure regulator to allow elevated pressure experiments to be carried out. All the pipework was heated to prevent condensation of the products.

2.3. Product Analysis Product analysis was carried out using a Hewlett-Packard 5890II gas chromatograph, fitted with both a thermal conductivity detector, and a methanator/fiame ionisation detector. Separation of the products was achieved using a 3m Porapak Q packed column, with argon carrier gas. Reference data and pure component injections were used to identify the major peaks, and response factors for the products and reactants were determined and taken into account in the calculation of the conversion and product distribution. In all cases stoichiometric gas mixtures were used and carbon balances were better than 97%. Conversions and yields were calculated as follows: Conversion, C[CHa]or C[CO2] = (% conversion of CH4 or CO2 into all products). Carbon monoxide yield, Y[CO] = (CO in products)/((CO2 + CH4) in reagents) x 100.

3. RESULTS AND DISCUSSION

3.1. Synthesis of High Surface Area Metal Carbides Before the transition metal carbides could be assessed as catalysts it was necessary to prepare samples of high surface area. Boudart and his co-workers revealed that by using a CHaffI-I2 mixture and temperature programmed reduction (TPR) it was possible to prepare high surface area binary carbides. Further work by Claridge et al. showed that by substituting C2H6

714 for CH4 in the gas mixture during TPR carbidation, a further increase in surface area could be achieved. Table 1 shows that with this technique a wide variety of group V and VI transition metal oxides, with BET surface areas <3 m 2 g~, could be converted into the carbide whilst increasing the surface area up to 174 m 2 g-1. Table 1 also shows the carbide phases synthesised and the final temperature needed to ensure that the reaction had reached completion. Amongst the group VI metals, MoO3 gave a h.c.p, carbide, while with WO3 the phase formed was dependent on the reactant gas; h.c.p, with ethane, and simple hexagonal with methane. This demonstrates that by modifying the TPR conditions slightly, it may be possible to access different phases of the materials. Group V metal oxides required higher temperatures both for CH4 and C2H6 TPR and formed f.c.c, structures. Ta205 was unique in that it underwent no reaction at all in C2H6. The mixed ternary oxides were also able to be carbided via CH4 and C2H6 TPR, being converted to their analogous single phase f.c.c, carbides; XRD and TEM [28] confirmed that no phase segregation was occurring, since no molybdenum or tungsten carbides could be observed. It was interesting to note that under C2H6, the temperatures required for carbidation of the ternaries was lower than that for the binary group V metal oxides. This would seem to infer that a synergistic effect exists, with the group VI metal reacting initially with the C2H6 to form a reactive carbon intermediate which could then react with the group V metal. The above reactions were followed by on-line mass spectrometry, and the products examined in detail by TEM; the results of these studies will be published elsewhere [28].

Table 1 Summary of some of the characteristics of the metal carbides. CH4 TPR Oxide Carbide phase: CH4/C2H6 (structure) Final T / K S 8 / m 2 g~ precursor MoO3 I]-MozC (h.c.p.) 1020 91 WO3 o~-WC/WzC (s.h./h.c.p.) 1150 39 V205 VCx (fc.c.) 1173 44 1173 62 Nb205 NbCx (fc.c.) 1223 54 Ta205 TaCx (fc.c.) WNb12033 WNbl2Cx(f.c.c.) 1173 76 W3Nb14044 W3Nbl4fx(fc.c.) 1173 87 W9Nb8047 W9NbsCx(fc.c.) 1173 88 1173 52 Mo2Ta2Ol~ Mo2Ta2Cx (f..c.c.) x = 0.73-0.97 [VCx]; 0.70-0.99 [NbCx] and [TaCx] [32].

C2H6 TPR Sg ]m2g-1

Final T / K 900 900 1023 1023 no reaction 900 900 900 900

174 71 76 113 107 126 102 84

It has been reported previously that simple oxide to nitride reactions appear to proceed in a topotactic fashion [ 18,33]. TEM studies of the binary and ternary metal carbides showed that the carbide materials maintained the overall morphology of the oxide precursor, but that the carbide particles were polycrystalline, as shown in Figure 1. Electron diffraction patterns of samples prepared by CH4 TPR gave rings patterns, as expected for these types of material. However, electron diffraction patterns of materials prepared by C2H6 TPR exhibited indexable spot patterns, and after measurement of a large number of patterns on crystallites of W9Nb8047, it was found that a preferred orientation exists which can be related to the original

715 oxide (see Figure 1); this gives strong evidence for topotacticity. Initially it was thought that the higher carbide surface areas achieved with ethane as the reactant gas were due to the lower temperature needed for the reaction compared with methane, and thus to less sintering of the sample. However, the topotacticity of the reaction may also be an explanation for this phenomenon.

o.oYo 6 00 |

C2H6/ H2 TPR (900K)

9

o.oYo~ .|

|247247 9l

[00 1] W9NbsO47

J,

[011 ] MC

Figure 1. Transmission electron micrographs and electron diffraction patterns of W9Nb8047 and its carbide formed by C2H6TPR.

3.2. Catalytic Reforming of Methane to Synthesis Gas Table 2 shows the initial results for the catalytic dry reforming of methane using bulk carbides of niobium, tantalum, molybdenum and tungsten, prepared by CH4 TPR. The conversions and yields obtained over 13-Mo2C and ~x-WC are very similar to those predicted by thermodynamic considerations, and thus these materials are efficient catalysts for methane dry reforming. At atmospheric pressure, deactivation was observed over both catalysts after about 8 hours on stream. Examination of the post-catalytic samples by powder XRD (Figure 2) revealed that as the reaction proceeded the active 13-Mo2C was oxidised and converted to

716

MOO2, which is inactive for methane reforming; similar XRD results were seen in the case of c~-WC, which was converted to WO2 during the reaction. However, when the reaction was carried out at elevated pressure no deactivation was seen, and the Mo2C and WC catalysts were stable for the duration of the experiments (>140 hours). Furthermore, no traces of M002 or WO2 could be seen in the XRD patterns of the post-reaction samples, and no carbide phase changes were observed.

Table 2 Results for the dry reforming of methane over group V and VI transition metal carbides. (GHSV = 2.87 x 103 h -1, CH4]CO2 = 1) Catalyst T/K P / bar C[CH4]/% C[CO2]/% Y[CO]/% H2/CO NbCx 1223 8.0 t 67.6 77.3 72.4 0.82 1373 8.0 83.7 96.3 90.0 1.33 TaCx 1223" 8.0 t 54.7 61.5 58.1 0.67 1123 1.0 t 92.4 92.5 92.5 0.93 I]-Mo2C 1223 1.0 t 98.8 95.9 95.9 0.92 1123 8.3 62.5 75.9 69.5 0.78 1223 8.3 83.3 89.5 86.5 0.88 1123 1.0 t 92.0 93.1 92.6 0.94 c~-WC 1123 8.3 62.7 75.4 68.6 0.79 * catalyst deactivates; * initial result could not be obtained.

1400 1200

A1

>~ 800 600

9 MoO2 " 13-Mo2C

,, 9

Ai

9

9

(c) _.zz____3

*

400

.

All

9 i AI

. .* . . .

9

*

9

9

II

200

20

3'o

4'0

5'0

6'0

7'o

Bragg Angle (20) Figure 2. XRD patterns of (a) [~-Mo2C as prepared by CH4 TPR, (b) [~-Mo2C post-dry reforming (8 bar, 1123 K), and (c) ~-Mo2C post-dry reforming (1 bar, 1123 K).

717

The primary cause for deactivation in current reforming catalysts, such as supported nickel, is the formation of carbon on the catalyst during reaction [20]. However, when postreaction samples of ~-MozC and a-WC were examined by HRTEM, no observable carbon deposition had occurred on the catalyst surface during the reaction. In addition, activity studies demonstrated that Mo2C had a methane dry reforming activity similar to an active supported noble metal catalyst, namely 5% Ir/Al203 [27]. Table 2 also shows that the group V transition metal carbides, NbCx and TaCx, are less active than the carbides of the group VI metals. At 1223 K, both NbCx and TaCx deactivated rapidly, even at elevated pressure, while catalyst stability and high reactant conversions were achieved at 1373 K with NbCx, where autothermal gas-phase reactions are likely to play a significant role in the reaction and carbon deposition was observed, as is demonstrated by the high H2/CO ratio. The reason for the low activity of these materials is the relative ease of their conversion back to the oxide, i.e. the rate of carbidation over these materials is slower than the rate of oxidation. Since MozC and WC were found to be active for the dry reforming of methane, they were also tested as catalysts for the partial oxidation of methane with air and methane steam reforming. As before, the product distributions were close to those predicted by thermodynamic equilibrium calculations, and deactivation occurred at atmospheric pressure due to oxidation of the carbide to inactive MoO2; in fact, with 02 (air) as the oxidant the catalyst deactivation was too fast for the initial activity to be measured. However, Table 3 shows that at elevated pressures the catalysts were highly active for both reactions, while no catalyst deactivation was observed for the duration of the experiments (>72 h), even with the strongest oxidant, i.e. oxygen. This is surprising, since the high surface area carbides are extremely reactive towards oxygen, even at room temperature, but is probably due to the relative rates of carbidation/oxidation during the reaction. At atmospheric pressure the oxidation is fastest, while the carbidation is favoured by elevated pressures.

Table 3 Catalytic results for the partial oxidation and steam reforming of methane using ~-WC. (GHSV = 5.2 x 103 h ~ (with air); 2.8 x 103 h -1 (with H20)) Catalyst CH4/ P/ T/ C[CH4]/% Y[CO]/% oxidant bar K ~-MozC 2:1 (air) 4.0 1173 94.7 89.6 2:1 (air) 8.7 1073 73.5 59.2 2:1 (air) 8.7 1173 88.5 81.3 a-WC 2:1 (air) 8.7 1173 87.5 78.2 ~-Mo2C 1:1 (H20) 1.0* 1223 91.5 90.1 1:1 (H20) 8.3 1223 81.8 77.7 8.3 1223 81.9 77.8 a-WC 1:1 (H20) catalyst deactivates

~-MozC and H2/ CO 2.01 2.07 2.01 2.04 3.08 3.07 3.06

718

3.3. Dry Reforming of Methane Using Ternary Metal Carbides as Catalysts Following the development of synthetic routes to produce single phase group V and VI ternary carbides, and the observation that the carbidation temperature of the ternaries was lower than the group V binary carbides, it was postulated that the synergistic combination of the group V and VI transition metals may result in the synthesis of highly active and stable carbide catalysts for methane dry reforming.

Table 4 Results for methane dry reforming over the ternary metal carbides (C2H6 TPR; T = 1223 K, p = 8 bar, GHSV = 5 x 103 h-1) Catalyst precursor C [ C H 4 ]/ % C[CO2]/ % Y[CO] / % MoO3 83.3 89.5 86.5 Mo2Ta2Oll 86.8 89.8 88.3 Mo3Nb2Ol4 82.0 89.9 86.0 Mo3Nb14044 47.7 64.3 56.0 W9Nb8047 73.3 92.1 82.7 W3Nb14044 26.1 22.8 24.4 WNb12033 22.4 25.5 24.0

H2/CO 0.88 0.84 0.90 0.60 0.83 0.78 0.63

4OOO

3O0O

(-. :3

0

23OO

o

i

i

I

~

(c~

1000

J '

II0

'

I

20

'

i

3o 20

,

i

40

,

i

50

'

i

m

,

i

7o

Figure 3. XRD patterns for (a) MoETaECx post-catalysis (1223 K, 8 bar), (b) Mo2Ta2Oll as prepared, and (c) Ta2Os.

At first glance, the results in Table 4 indicate that some of the ternary metal carbides, namely those synthesised from Mo2Ta2Oll, Mo3Nb2O14, and W9Nb8047, are highly active for

719 methane dry reforming, with reactant conversions and yields very similar to those obtained with [~-MozC. Further, the catalysts appeared to be stable during lifetime investigations of more than 50 hours (not shown). However, a comparison of the pre- and post-reaction XRD patterns revealed that all the single phase ternary carbides are phase separating during the reaction. In fact, as was the case with the binary metal carbides, the group V metal carbides are oxidised, leading to the formation of the group V metal oxide (Nb205 or Ta2Os), while the group VI metal carbide is stable under these conditions (8 bar). Figure 3 shows that the postreaction XRD pattern of the carbide of MozTazOll has a strong resemblance to the pattern for pure Ta2Os; no evidence for any tungsten oxides was observed, and it is probable that the tungsten is present in the sample in the form of tungsten carbide, which is characterised by weak and broad peaks, and thus cannot be seen. Similar XRD patterns were also obtained for the other ternary carbides, and these observations were also supported by TEM and energy dispersive X-ray analysis (not shown). Therefore, it would appear that the stability of some of the ternary carbides is actually due to the presence of stable Mo2C or WC. This is further confirmed by the observation that only those carbides with a high ratio of group VI to group V metal are active for dry reforming, while in the other ternaries, such as WNblzCx, the amount of tungsten is small and the catalyst surface mainly consists of the metal oxide; therefore, the number of catalytically active carbide species at the surface is very small.

4. CONCLUSIONS It has been demonstrated that the use of ethane in the temperature programmed synthesis of transition metal carbides results in the formation of materials with higher surface areas than with methane TPR. In addition, the conversion process with ethane appears to proceed in a topotactic fashion. Of all the materials synthesised, only molybdenum and tungsten carbide were active and selective for the stoichiometric dry reforming, partial oxidation and steam reforming of methane to synthesis gas. These materials deactivated at atmospheric pressure, but were very stable when elevated pressures were employed, and no carbon deposition was observed on the catalysts. Niobium and tantalum carbide were inactive for the dry reforming of methane, due to their rapid deactivation by oxidation. Some of the ternary carbides were found to be active and stable for dry reforming, but the catalyst activity could be attributed to the presence of molybdenum or tungsten carbide. In fact, the group V metal appeared to play no role in the processes contributing to the catalytic reforming of methane to synthesis gas. The observations presented in this paper are potentially important, since these materials offer the chemical industry cheap and abundant, non-coking reforming catalysts, where previously the choice was between nickel, which promotes carbon formation, and expensive noble metals.

REFERENCES 1. J.M. Muller and F. G. Gault, Bull. Soc. Chim. Fran., 2 (1970) 416. 2. R.B. Levy and M. Boudart, Science, 181 (1973) 547. 3. L. Volpe and M. Boudart, J. Solid State Chem., 59 (1985) 348. 4. J.S. Lee, S. T. Oyama and M. Boudart, J. Catal., 106 (1987) 125. 5. D. Zeng and M. J. Hampden-Smith, Chem. Mater., 4 (1992) 968.

720 6. 7. 8. 9. 10. 11.

12.

13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33.

M.J. Ledoux and C. Pham-Huu, Catal. Today, 15 (1992) 263. J. Lemaitre, B. Vidick and B. Delmon, J. Catal., 99 (1986) 415. L. Leclercq, M. Provost, H. Pastor, J. Grimblot, A. M. Hardy, L. Gengembre and G. Leclercq, J. Catal., 117 (1989) 371. F. Sherif and W. Vreugdenhil, in "The Chemistry of Transition Metal Carbides and Nitrides", S. T. Oyama (ed.), p.414, Blackie Academic & Professional, Glasgow, 1996. G.B. Raupp and W. N. Delgass, J. Catal., 53 (1979) 361. C. Pham-Huu, M. J. Ledoux and J. Guille, J. Catal., 143 (1993) 249. M. J. Ledoux, C. Pham-Huu, A. P. E. York, E. A. Blekkan, P. Delporte and P. Del Gallo, in "The Chemistry of Transition Metal Carbides and Nitrides", S. T. Oyama (ed.), p.373, B lackie Academic & Professional, Glasgow, 1996. F. H, Ribeiro, R. A. Dalla Betta, M. Boudart, J. Baumgartner and E. Iglesia, J. Catal., 130 (1991) 86. F. H. Ribeiro, R. A. Dalla Betta, M. Boudart and E. Iglesia, J. Catal., 130 (1991) 498. J. C. Schlatter, S. T. Oyama, J. E. Metcalfe III and J. M. Lambert Jr., Ind. Eng. Chem. Res., 27 (1988) 1648. J. S. Lee and M. Boudart, Appl. Catal., 19 (1983) 207. L. Volpe and M. Boudart, J. Phys. Chem., 90 (1986) 4874. L. Volpe and M. Boudart, J. Solid State Chem., 59 (1985) 332. J. T. Wrobleski and M. Boudart, Catal. Today, 15 (1992) 349. J. R. Rostrup-Nielsen, in "Catalysis, Science and Technology", J. R. Anderson and M. Boudart (eds.), vol. 5, p. 1, Springer, Berlin, 1984. S. T. Sie, M. M. G. Senden and H. M. H. Wechem, Catal. Today, 8 (1991) 371. S. C. Tsang, J. B. Claridge and M. L. H. Green, Catal. Today, 23 (1995) 3. J. R. Rostrup-Nielsen, J. Catal., 85 (1984) 31. J. R. Rostrup-Nielsen, J. Catal., 31 (1973) 173. A. T. Ashcroft, A. K. Cheetham, M. L. H. Green and P. D. F. Vernon, Nature, 352 (1991) 225. J. B. Claridge, M. L. H. Green, S. C. Tsang, A. P. E. York, A. T. Ashcroft and P. D. Battle, Catal. Lett., 22 (1993) 299. A. P. E. York, J. B. Claridge, A. J. Brungs, S. C. Tsang and M. L. H. Green, J. Chem. Soc., Chem. Commun., (1997) 39. J. B. Claridge, A. J. Brungs and M. L. H. Green, Chem. Mat., submitted. M.W. Viccary and R. J. D. Tilley, J. Solid State Chem., 104 (1993) 131. T. Ekstrom, Acta Chem. Scand., 25 (1971) 2591. J. B. Claridge, M. L. H. Green, S. C. Tsang and A. P. E. York, Appl. Catal., 89 (1992) 103. E. K. Storms, The Refractory Carbides, Academic Press: New York, 1967, vol. 2. S. T. Oyama, R. Kapoor, H. T. Oyama, D. J. Hofmann and E. Matijevic, J. Mater. Res., 8 (1993) 1450.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

721

Partial o x i d a t i o n of m e t h a n e to s y n t h e s i s gas using L n C o O 3 p e r o v s k i t e s as catalyst precursors R. Lag&, G. Bini b, M.A. Pefia" and J.L.G. Fierro" 'llnstituto de Catfilisis y Petroleoqu/mica, CSIC, Campus UAM, Cantoblanco, 28049 Madrid, Spain [FAX +34 1 585 4760; E-mail: [email protected]] bDipartimento de Chimica lndustriale e Materiali, University of Bologna, Vie. Risorgimento 4, 40136 Bologna, Italy

In this work Co supported catalysts have been prepared by the reduction of the perovskites LnCoO~ (Ln=La, Pr, Nd, Sm and Gd) to produce Co"/Ln20~. Detailed TPR and XRD studies showed that the perovskite NdCoO3 is reduced in two steps, first to NdCoO2s and further to Co~ and in both stages it was demonstrated that the reoxidation with Oe is capable to recover the perovskite structure. TPO experiments with reduced Ln-Co-O (Ln = La, Nd, Sm and Gd) catalysts indicated that reoxidation takes place in two steps: first oxidation of the supported Co" to the spinel Co~O4 (Co:+Co3+204) and further the oxidation of the Co :+ to Co 3+ with a simultaneous solid state reaction with Ln20~ regenerating the perovskite structure. It was observed that the temperature for the second oxidation step is strongly dependent on the nature of the lanthanide increasing in the following order La > Nd > Sm > Gd. This trend seems to be determined by the thermodynamic stability of the parent perovskite. These catalysts (Co"/Ln:O3) have been tested for the partial oxidation of methane to synthesis gas showing remarkable differences in activity [1]. The system Gd-Co--O showed exceptionally better performance for CO and H2 production whereas the activity for the other catalysts decreased in the following order Sm-Co-O > > Nd-Co-O > Pr-Co-O. The catalyst La-Co-O was active for methane combustion and only traces of CO and H: were observed. It was found that these differences are determined by the lanthanide which plays a fundamental role on the stability of the catalyst.

1. INTRODUCTION Perovskites of the type LnMO 3 (Ln = lanthanide and M - transition metal) may offer interesting features as precursors for supported metal catalysts. For example, careful reduction can be carried out in order to produce a finely dispersed transition metal over the sesquioxides Ln:O~. Also, the flexibility in the perovskite composition allows the preparation of compounds of the type LnM'~_,,M"xO~ (M' and M" = different transition metals) or Ln~_,,A,,MO~ (A = for example an alkaline or alkaline earth metal), which show the unique possibility as precursor of producing well dispersed bimetallic catalyst or doped metal catalyst. The research on CH4 conversion has recently focused on the direct oxidation to synthesis gas:

722

CH4 + 1 / 2 0 2 - - >

CO +

2H2

( A N~ ,, 2,~8~ = - 35.6 kJ/mol)

(1)

The major advantages of this route over the steam reforming are the H:/CO ratio of ca. 2 suitable for downstream processes and the exothermicity of the reaction which eliminates the need for a fuel gas [2]. Many catalysts for the partial oxidation of methane to synthesis gas, consisting of supported metals such as Ni, Co and Fe and noble metals Pd, Ir, Rh, Ru, Pt, etc have been described in the literature [2-9]. Hayakawa et al. [ 11] studied the perovskite Ca~_,,Sr,,TiO3 mixed with nickel oxide for methane oxidation at 1028 K, both before and after pre-treatment with methane. Before methane pretreatment the catalyst produced mainly C2 hydrocarbons and CO2 but no syngas, whereas after methane pre-treatment at 1048 K for 1 h, the mixture became very selective to synthesis gas giving 70.9% CH4 conversion with 94% selectivity to CO and H 2. Based on XRD analyses it was proposed that the methane pre-treatment produced Ni" supported on the perovskite which was responsible for the synthesis gas formation. Similar results were obtained for cobalt and iron (8). Slagten and Olsbye [10] studied the systems La-M-O (M = Co, Ni, Rh and Cr) for the partial oxidation of methane to syngas. They observed very high activity for the system La-Rh-O whereas the catalyst La-Co-O (which was a mixture of LaCoO3, La203 and Co304) produced mainly CO:. If the catalyst La-Co-O was kept at 1073 K after 30 h reaction the activity changed to give mainly CO which they assigned to the in situ reduction of cobalt. In this work we have studied the solid state reactions taking place with supported cobalt metal catalysts prepared by the reduction of LnCoO 3 precursors under the experimental conditions of the partial oxidation of methane.

2. EXPERIMENTAL All the perovskites were prepared by the method of the amorphous citrate percursor [12] and were present as pure phases with no contaminants of LnzO3 or cobalt oxides according to the powder XRD analyses. The catalytic tests were carried out in a quartz fix bed microreactor with 50 mg of catalyst and a 2:1"4 mixture of CH4:O:'Ar at 166.6 ml/min. All the catalysts were reduced in situ with a H2 (33%):Ar mixture at 1023 K for 3.5 h prior the reaction. Two thermocouples were used, one outside the reactor to control the furnace temperature and the other one inside in contact with the catalyst to measure the bed temperature. X-ray photoelectron spectroscopy (XPS) analyses were carried out in a ESCALAB 200R spectrometer equipped with a MgKa 120 W X-ray source (hv = 1253.6 eV). After reaction the catalysts were all quenched to room temperature under argon and immediately drenched in isooctane. Despite this careful procedure to isolate the catalyst it was not possible to completely avoid oxidation of superficial metallic cobalt during the transfering of the sample from the reactor to the XPS spectrometer. Powder XRD were obtained in a Seifert 3000P diffractometer. Temperature-programmed reduction (TPR) and oxidation (TPO) analyses were carried out in a Micromeritics 2900 apparatus. For the Co" area determination, H2 chemisorption measurements were performed in a conventional volumetric adsorption apparatus. A 200 mg sample of the perovskite was reduced at 1023 K for 3.5 h under a stream of 33% H2/N2 and then outgassed at 773 K overnight at pressures lower than 10 -4 Torr. The measurements were carried out at 315 K and the reproducibility was within 10%.

723 3. RESULTS 3.1. Reduction of the Cobalt Containing Perovskites TPR profiles of the cobalt containing perovskites are displayed in Fig. I. All the perovskites showed similar reduction profiles consisting of two sets of peaks at approximately 633 and 833 K. It can be observed that especially for LaCoO3 and PrCoO 3 the second reduction peak shifts to higher temperatures. For example, the reduction peak for SmCoO3 at c a . 785 K is observed at 844 K for LaCoO3. In all cases the hydrogen consumption for the first reduction step (peak at c a . 633 K) was always approximately half of the hydrogen consumption obtained for the second reduction step (peak at 833 K). Careful thermogravimetric reduction experiments (not shown here) demonstrated that the first step is a 1 electron reduction process whereas the second step is a 2 electrons reduction process. To identify the reduced species formed in the TPR steps in Fig. 1, XRD studies of NdCoO 3 A

-k 9

&

Q

__

_a =

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C

v

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~

,~

IC C

t ?'i" 9

At'?

,_ 5,,L_.Y.__.2_J,._..,' -b._.,,,.._...._..,,._~

=

=

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c

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B

Nd-Co-O Pr-Co-O La-Co-O I

400

A tO ~,,.

I

600

,

I

~()0

T e m p e r a t u r e (K)

1()()()

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21)

2 0 (")

Figure 1. TPR profiles of LnCoO3 perovskites. Figure 2. XRD patterns for NdCoO3: (a), as prepared" (b), reduced in H: at 623 K for 15 min; (c), after reduction (623 K, 15 min) and heated under He at 1123 K for 1 h (sintering); (d), reduced in H2 at 778 K for 15 min" (e), after reduction (778 K, 15 min) and sintering; (f), after reduction (778 K, 15 rain), sintering at 1123 K and reoxidation under O2 flow at 973 K. A = N d C o O 3 " B = N d : O 3 ( c u b ) ; C = N d : O 3 ( h e x ) " D = C o O ; E = Co~O4" a n d F = C o ~.

724 subjected to different reduction treatments were carried out (Fig. 2). The XRD pattern displayed in Fig. 2a corresponds to NdCoO~ cubic perovskite structure. Upon reduction at 623 K under H2 flow for 15 min the diffraction peaks are essencially at the same position but with a very strong broadening and in some peaks a splitting is observed (Fig. 2b). This indicates that the perovskite structure is basically the same however distorted due to the creation of anion vacancies. These results suggest that the stoichiometry of the phase formed in the first step 1 electron reduction can be represented as NdCoO2.~. The reoxidation of this sample at 973 K under oxygen flow can easily regenerate the original perovskite structure (XRD not shown here) demonstrating the reversibility of this process. On the other hand, the phase NdCoOz.5 is not stable and at high temperatures (1123 K for 1 h) under He phases of CoO and Nd203 hexagonal are formed (Fig. 2c). When the perovskite was treated with hydrogen at 776 K for 15 min, a 3 electron reduction was observed and XRD analysis showed the presence of Nd203 hexagonal and very weak and broad peaks which could be due to metallic coba!t (Fig. 2d). Upon heating at 1123 K in helium the highly dispersed cobalt metal sinters and shows much stronger lines in the XRD pattern (Fig. 2e). It was also observed the presence of small amounts of CoO, probably formed by the reoxidation of cobalt by water or oxygen still present in the sample during sintering. The reoxidation of the reduced sample in Fig. 2d (before sintering) also reproduced the perovskite original structurewith the same XRD pattern (not shown here). However, after sintering (Fig. 2e) the reoxidation produces a mixture of the phases NdCoO3, Co304 and Nd203 (Fig. 2f), showing that the process is not completely reversible. It was observed that if the sample in Fig. 2f was kept under oxygen flow at temperatures of 1123 K for 2 h a slow solid state reaction takes place regenerating the perovskite NdCoO~ structure. Therefore, the two main reduction steps of NdCoO 3 can be written as" 1 e/mol"

2NdCoO~ + H: ~

3 e/mol"

2 NdCoO 3 + 3H: ~

2NdCoOes

+ H:O

Nd203 + 2Co" + 3H20

(2) (3)

These reactions are reversible and as long as the metallic cobalt remains well dispersed the perovskite structure can be easily regenerated by oxidation. These results are in agreement with the work of Crespin and Hall [ 13], who observed basically the same reduction steps for LaCoO 3. The cobalt metal area of the reduced perovskites was determined by hydrogen chemisorption experiments. The results are shown in Table 1. The chemisorption measurements revealed that the cobalt metallic surface area was similar for all the perovskites. This is supported by the Co/Ln surface ratio (Table 1) obtained by XPS which also suggests similar metallic dispersion. The XPS analyses of the reduced perovskites showed the presence of Co" (778.6 eV) but also a doublet at approximately 780.5 and 796.2 eV which correspond to Co 2p3/2 and Co 2p~/: peaks respectively, for the Co -~+ ion. Shake-up satellite lines with 4.7 eV over the Co 3+ lines were also detected indicating the presence of Co :+ [12]. These oxidised species of cobalt are probably formed by air oxidation during the transference of the reduced sample from the reactor to the XPS spectrometer. Also, Marcos et al. [15] have shown that the reduction of the perovskite LaCoO 3 produced a La203 oxide covered by hydroxyl groups which upon heating and evacuation in the XPS pretreatment chamber partly reoxidises the cobalt crystallites. Powder XRD analysis of the reduced LnCoO~ perovskites showed only the presence of the sesquioxides La20~ (hex), NdzO~ (hex), Sm203 (cub) and Gd203 (cub). The apparent absence of reflections for the metallic cobalt indicates a high metallic dispersion with Co" particles smaller

725

Table 1 Characterization of the reduced perovskites Precursor

BET area (m:/g)"

Co" area (atom/g) ~'

LaCoO3 NdCoO3 SmCoO3 GdCoO 3

8.6 (5.4) 5.6 (4.5) 5.1 4.0 (5.0)

3.77- 10 I'~ 3.98-10 ~'~ 3.69.10 ~'~ 3.21.10 j''

XPS Co/La ratio (reduced cat.y 0.71 0.50 0.58 0.75

XRD (reduced cat.) La203 (hexagonal) Nd203 (hexagonal) Sm203 (cubic) Gd:O3 (cubic)

a BET areas for reduced catalysts are in parenthesis; b Catalysts reduced at 1023 K under H2 (33%) in Ar; c For Co/La ratio measurement all Co signals (reduced, oxidised forms and shakeup lines) were included. than 2 nm. Considering the approximate value of 3.10 t'~ atoms/m: for a close packing arrangement [ 16], the number of surface cobalt metallic atoms obtained for the reduced catalysts is similar to the number of surface cobalt ions on the original perovskite surface (between 2.10 ~~ to 5.10 ~~ at/g for GdCoO 3 with BET surface area of 4.0 m:/g and LaCoO 3 8.6 m2/g, respectively). Therefore, apparently after reduction of the perovskites great part of the cobalt metal is located in the bulk of the material.

3.2. Reoxidation of Co~ TPO experiments were performed for the reduced catalysts in order to investigate the stability of these systems towards gas phase oxygen (Fig. 3). The perovskites were completely reduced in a TPR experiment with a stream of 10% H: in Ar heating from room temperature to 923 K at a rate of 10 K/min. The temperature was then rapidly decreased to avoid sintering of the cobalt metal and TPO analysis performed. TPO profiles showed that the oxidation takes place in two steps. Thc first one at near 473 K (peak I) has an oxygen consumption which suggests that the cobalt is oxidised to two Co ~+ and one Co :+. Therefore, this oxidation probably corresponds to the formation of the cobalt oxide spinel according to the reaction" 3Co" + 202

~

Co304

(4)

The second step (peak II) takes place at a much higher temperature and the 02 consumption suggests the following process:

Co304 --I- 1/4 02 + 3 Ln203 ~

LnCoO 3

(5)

In this step apparently the Co :+ is oxidised to Co 3+ simultaneously with a solid state reaction to regenerate the perovskite structure. The occurrence of these reactions are also supported by a simple experiments involving sequences of TPO and TPR measurements, shown in Fig. 4A and 4B. It can be observed that the TPR profile, obtained with the sample reduced in the first TPR (Fig. 4B, profile a), and reoxidised in the first TPO up to 1123 K (Fig. 4A, profile a) is very similar to the TPR of the fresh sample (Fig. 4A, profile a) suggesting the presence of a SmCoO3 perovskite structure (Fig. 4B, profiles a and b). On the other hand, if the reduced sample (after

726

Nd-Co-O --

c

400

600

800

1000

12(

Temperature (K) Figure 3. TPO profiles for the reduced systems La-Co-O, Nd-Co-O, Sm-Co-O and Gd-Co-O. TPR) is reoxidised in a TPO experiment which is interrupted at 623 K and the sample heated up to 1123 K under He, the TPR analysis showed only a broad peak at ca. 623 K (Fig. 4B, profile c) which is similar to the reduction of the spinel Co304 reported in the literature [17]. It is interesting to note in Fig. 3 that the temperature for the first oxidation step is similar for all the catalysts whereas the temperature of the second step is strongly dependent on the nature of the lanthanide. For La, oxidation occurs at 939 K whereas for Gd a much higher temperature is necessary (1109 K). Also the area and the shape of the second peak varied significantly for the different lanthanides. Very sharp peaks were obtained for Ga and Sm whereas a broad and

02

:11

A

B

gl,

'2 //-"~

r

eq

/ /

0 inl c riu p ted C

I

400

~

1

~

60t)

I

800

T e m p e r a t u r e (K)

~

I

1000

,

f

-

,

120{}

J

I

600

750 T e m p e r a t u r e (K)

Figure 4. Sequences of TPO (A) and TPR (B) experiments with the system Sm-Co-O.

900

727 weak peak was observed for La. The Apcak i/Apcak II area ratio of the first and the second peak from the TPO profiles shown in Fig. 3 are 12.5; 10.0; 9.0 and 7.5 for La-Co-O, Nd-Co-O, Sm-Co-O and Gd-Co-O, respectively. The systems La-Co-O and Nd-Co-O can be observed to deviate significantly from the expected area ratio which is 8.0 considering the stoichiometry of Eqs. [4] and [5]. This suggests that part of the cobalt which in the Gd-Co-O and Sm-Co-O systems is only oxidised at temperatures higher than 973 K, is being oxidised at lower temperatures for La-Co-O and Nd-Co-O. In fact, Crespin et al. [23] showed that if the reduced LaCoO 3 was kept at a temperature as low as 673 K under oxygen the Co" was completely reoxidised regenerating the perovskite structure. 3.2. Influence of the ianthanide on the reduction-oxidation of the system Ln-Co-O As observed by TPR, the nature of the Ln affects the reducibility of Co in the perovskites LnCoO 3. The Goldschmidt tolerance factor t - (r~,+ ro)/[~2 (r(.o+ro) ] obtained for the structures of LaCoO3, PrCoO3, NdCoO3, SmCoO3 and GdCoO3 were 0.899, 0.885, 0.878, 0.867 and 0.857, respectively. These tolerance factors indicate that considering solely geometric factors lanthanum, the largest ion in the series, forms the most stable perovskite structure. This trend is reflected in the TPR results where the perovskite LaCoO3, the most stable structure, is reduced at the higher temperatures, 844 K (Fig. 5). Likewise, the TPO experiments showed that reoxidation of cobalt to form the perovskite structure is more favorable for larger lanthanides. Figure 5 shows a good correlation between the oxidation temperature obtained from the TPO profiles with the Goldschmidt's tolerance factor. Katsura et al. [ 18] studied the thermodynamics between 1473 and 1673 K of the oxidation of iron to the rare earth perovskites according to the reaction:

Fe(s) + 1/2 Ln203 + 3 / 4 0 : (g) ~

(6)

LnFeO3 (s)

They showed that reoxidation to form the perovskite structure is favored in the order La > Nd > Sm > Gd with standard Gibbs energies o f - 2 8 8 . 0 , -274.6, -267.9 and -263.7 kJ/mol, respectively. k;i o

0.90

[]

Pr c]

d

Nd El

.~ 0.88 t_ ..... SITI cl 0.86

Gd

o 780

i

I

800 Reduction

i

I

i

820 Temperature

TPR (K)

I

I

I

840

950

1000 Oxidation

i

Temperature

I

~

1050

I

1100

T P O (K)

Figure 5. Goldschmidt's tolerance factor t versus (a) reduction and (b) oxidation temperatures obtained by TPR and TPO experiments for the perovskites LnCoO~.

728

3.3. Catalytic Testing Among the cobalt containing perovskites GdCoO3, SmCoO~, NdCoO~, PrCoO3 and LaCoO3 tested as catalyst precursors for the partial oxidation of methane the Gd-Co-O system showed exceptionally better performance for synthesis gas formation (Figs. 6A-6C). At 1009 K a steadystate methane conversion of 73% with selectivities of 79 and 81% for CO and H2, respectively, is observed for the catalyst Gd-Co-O. The catalysts Sm-Co-O and Nd-Co-O, of lower activity, show similar steady-state methane conversions in the temperature range studied. On the other hand, the H2 and CO selectivities are much higher over Sm-Co-O. The catalyst La-Co-O is active for the methane combustion and only traces of H2 and CO were observed under the reaction conditions used (Fig. 6A-6C). It is interesting to observe that although the reduced perovskites possess similar cobalt metallic areas, as revealed by H2 80 A

~ ~ - - - - [] . - - - - ' ~ []

' ~ o

60

...__~0

411 211-

~

~

0

91)!I

800 100

75

~

B

o_i ~ o ~ O ~

1000

Gd-Co-O

.______-----o-~ - - - [ ]

~

1100

Sm-Co-O

~

~.. ":

51)

-:'

25

,._co_o

-

o ~

____+__________T_o~

(} 800

Pr-Co-O La-Co-O

900

1000

I 1O0

1000

1100

I00

c ~"

75

~

_

~

o

"= 50 25 800

900

T e m p e r a t u r e (K)

Figure 6. (A), Methane conversion and (B), H2 and (C), CO selectivities in the presence of LnCo-O systems prereduced under a 30% HJAr flow at 1023 K for 3.5 h.

729

Table 2 XRD and XPS analyses of the catalyst after reaction" Precursor

Co/Ln XPS ratio b

LaCoO3 NdCoO.~ SmCoO~ GdCoO~

0.7 0.2 0.3 0.5

XRD LaCoO3 and La203 c NdCoO~ and Nd203 (hexagonal) SmzO3 (cubic) Gd203 (cubic)

'~ After reaction reaction at 1023 K for 19 h; b For Co/Ln ratio measurement all Co signals (reduced, oxidised forms and shake-up lines were included" c traces chemisorption and XPS data, they showed strikingly different catalytic properties. XRD and XPS characterization data of the catalysts after reaction are summarized in Table 2. The XPS Co/Ln ratios suggest that the Co dispersion over the catalyst surface after reaction follows the order LaCo-O > Gd-Co-O > Sm-Co-O > Nd-Co-O. XRD analyses of the used catalysts Gd-Co-O and Sm-Co-O showed similar patterns to the reduced catalysts with very strong and sharp peaks for the sesquioxides Gd203 and Sm203. On the other hand, the XRD analysis of the La-Co-O catalyst after reaction at 1023 K for 19 h clearly showed the formation of the perovskite LaCoO 3. Therefore, it is not surprising that the only reaction products observed were water and carbon dioxide. This agrees with previous works on this perovskite and other forms of cobalt oxide which have been shown to be active catalysts for methane combustion and also for CO and H: oxidation [19]. The high Co/Ln surface ratio determined by XPS for the used catalyst is expected for a perovskite like surface. Slagten and Olsbye [ 10] studied the perovskite LaCoO~ (containing some impurities of La20~ and Co~O4) for the partial oxidation of methane to syngas and observed the production of mainly CO,,. If the catalyst was kept at 1073 K after 30 h on-stream the activity changed to give mainly CO which they assigned to the in situ reduction of cobalt. The XRD for Nd-Co-O after reaction revealed the presence of the phases Nd20~ and also the perovskite NdCoO3. For all used catalysts no clear evidence for the presence of simple cobalt oxides such as CoO, Co,O~ and Co304 could be found by XRD.

4. CONCLUSION This work suggests that the high activity and selectivity of the catalysts Gd-Co-O and Sm-Co-O for the partial oxidation of methane to synthesis gas is due to the stability of the cobalt in its reduced state over the sesquioxides Gd:O3 and Sm20~. In the case of La-Co-O and Nd-Co-O reoxidation of cobalt to the original perovskite structure causes loss of activity and selectivity. TPO experiments with reduced Ln-Co-O (Ln = La, Nd, Sm and Gd) catalysts indicated that reoxidation takes place in two steps: first oxidation of the supported Co ~ to the spinel Co304 (Co2+Co3+204) and further the oxidation of the Co `'+ to Co ~+ with a simultaneous solid state reaction with Ln20~ regenerating the perovskite structure. It was observed that the temperature for the second oxidation step is strongly dependent on the nature of the lanthanide increasing in

730 the following order La > Nd > Sm > Gd. This trend seems to be determined by the thermodynamic stability of the parent perovskite.

Acknowledgements This work was supported by CICYT, Spain (Contract MAT95-0894). One of the authors (R.M.L.) is grateful to the Ministerio de Educaci6n y Ciencia, Spain, for the award of postdoctoral fellowship.

REFERENCES .

2. 3. 4. .

6. 7.

10. 11. 12. 13. 14. 15. 16. 17. 18. 19.

R.M. Lago, G. Bini, M.A. Pefia and J.L.G. Fierro, J. Catal., (1997), in press. B.M. Tindall and M.A. Crews, Hydroc. Proc., 11 (1995) 75. S.T. Tsang, J.B. Claridge and M.L.H. Green, Catal. Today, 23 (1995) 3. A.T. Ashcroft, A.K. Cheetham, J.S. Food, P.D.F. Vernon and M.L.H. Green, Nature, 344 (1990) 319. D.A. Hickman and L.D. Schmidt, Science, 259 (1993) 343. D. Dissanayake, M.P. Rosynek and J.H. Lunsford, J. Phys. Chem., 97 (1993) 3644. D. Dissanayake, M.P. Rosynek, K.C.C. Kharas and J.H. Lunsford, J. Catal., 132 (1991) 117. F. Looji, J.C. Giezen, E.R. Stobbe and J.W. Geus, Catal. Today, 21 (1994) 495. T. Hayakawa, A.G. Andersen, M. Shimizu, K. Suzuki and K. Takehira, Catal. Lett., 22 (1993) 307. A. Slagten and U. Olbsbye, Appl. Catal., 110 (1994) 99. T. Hayakawa, A.G. Andersen, M. Shimizu, K. Suzuki and K. Takehira, Catal. Today, 24 (1995) 237. J.M.D. Tascon, S. Mendioroz and L.G. Tejuca, Z. Phys. Chem. NF, 124 (1981) 109. M. Crespin and W.K. Hall, J. Catal., 69 (1981) 359. T.J. Chuang, C.R. Brundle and D.W. Rice, Surf. Sci., 60 (1976) 286. J.M. Marcos, R.H. Buitrago and E.A. Lombardo, J. Catal., 105 (1987) 95. G.A. Somorjai, "Introduction to Surface Chemistry and Catalysis", Wiley Interscience Pub., New York, 1994. J.G. Choi, Catal. Lett., 35 (1995) 291. T. Katsura, K. Kitayama, T. Sugihara and M. Kimizura, Bull. Chem. Soc. Jpn., 48 (1975) 1809. M. Futai, C. Yonghua and L. Louhui, React. Kinet. Catal. Lett., 31 (1986)47.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

731

Performance o f catalytic properties o f reagent catalyst in the processes such as methane oxidative coupling and h y d r o g e n production by methane conversion M.I. Levinbuk a, N.Y. Usachev b, M.L. Pavlov c, A.U. Loginovd, L.V. Surkova ~ E.M. Savin ~, V.K. Smirnov e, I.V. Ivkova~ aGubkin State Academy of Oil and Gas. 117917, Leninsky prospect 65, Moscow, Russia blnstitute of Organic Chemistry, 117913 Leninsky prospect 47, Moscow, Russia ~

Catalyst Plant, 453210, Ishimbai, Russia

dMoscow State University, 117234, Moscow, Russia eCatachem Company Ltd., 129832, Giliarovskogo 3 l, Moscow, Russia

1. INTRODUCTION More than 400 articles on methane oxidative conversion have been published for the past decade [1]. In order to realize this process in industry a new method based on a welldeveloped production technology has to be created. As it was described in the literature, metals with variable valency can be utilized as catalysts to convert methane into higher hydrocarbons [2]. Introduction of variable valency metals into cracking catalysts initiates methane partial oxidation reactions based on catalytic cracking technology. The catalyst used in this process circulates between the reactor and regenerator thus providing the partial oxidation of the feed in the first vessel while further oxidation of the reduced metal oxide occurs in the second one. We call it as a "reagent catalyst" since it functions as the carrier of the reagent (oxygen in this case) [3]. The feasibility of catalytic cracking technology utilization for the processes of production of gCz-hydrocarbons and pure hydrogen by methane oxidative catalytic conversion are considered in this article. US oil refinery plants will probably suffer from the shortage of hydrogen because the "reformulated gasoline" program demands to decrease the reforming capacity.

2. E X P E R I M E N T A L

The investigated catalysts were prepared using TCC (Thermofor Catalytic Cracking) catalysts as the base. Starting oxides MnO, Mn304, MnzO3 and MnO2 (at 10 wt% of Mn• per dry catalyst) were introduced by a special technique into a gel-forming solution. All the TCC catalysts contained 5% of rare-earth Y zeolite (with a molar ratio S I O 2 / A 1 2 0 3 = 6 obtained by direct synthesis). Nickel incorporating pentasil zeolites (3 - 8 wt % of NiO) were

732 prepared by impregnation of zeolite samples having different molar ratios SIO2/A1203 with a nickel-containing salt solution. Tested samples of the catalysts were exposed to various pretreatments: at air temperatures up to 1073 K without steam and at 1023 K with 100 % steam. Methane coupling reaction in the absence of oxidants in feed was carried out in an impulse microreactor; the charge was 0.1 g of the catalyst; one pulse of methane was 1.17 ml. Methane coupling reaction over TCC catalysts involving magnesium oxides with different starting oxidation levels was studied in a plug-flow reactor (with a volume of 1 ml) at atmospheric pressure, methane feed volume rate of 1000 h -1 (without oxygen) and reaction time of 1 min. The reaction was investigated within a temperature range of 973 - 1073 K. Methane catalytic conversion into carbon and hydrogen was examined over nickelcontaining pentasil zeolites in a vacuum-circulation laboratory unit [4] at a catalyst/feed mass ratio of 5.0 and within a temperature range of 743 - 843 K. The Mn +2 ion intensity variation in TCC catalysts in oxidation - reduction (reactionregeneration) cycle was observed by EPR method on a JeoI-JES-3BS-Q spectrometer (9 GHz) at 77 K and 295 K. Changes in the state of Ni +2 applied on pentasil in oxidation-reduction cycle were detected by oxygen absorption technique at a temperature of 77 K.

3. RESULTS AND DISCUSSION The utilization of metals with variable valency in catalysts in an oxidation-reduction cycle (separately in the reactor and regenerator when used in the catalytic cracking process) allows us to use the following reaction equation so as to depict the methane partial oxidation model: (p+oQCH4 + MeOm --> [3C2H4 + (o~- [~)C02 + pCH4 + MeOm_n

(1)

From the molar ratio of metal oxide to feed for the model reaction (1) a mathematical dependence can be obtained: y = Z * AS/(13/~t) * (L * Mf)/M m

(2)

where y = Gt / (p + or) - methane conversion, Z - catalyst/feed mass ratio, L - the metal oxide catalyst content, Mr, Mm - the molecular weights of feed and metal oxide, AS - variation in the valency of metal oxide. The verity of the equation (2) was confirmed by an experiment on methane conversion over TCC catalyst containing 10 wt% of Mn203 in an impulse microreactor at 1073 K and different catalyst/feed ratios (Table 1). From Table 1 it follows that the selectivity of the yield of C2-hydrocarbons and complete oxidation products depends antibately on the catalyst/feed ratio. In the range of the studied values of the catalyst/feed ratios the minimal yield of the complete oxidation products corresponds to only 2.5 wt% methane conversion.

733 Table 1 Conversion and selectivity of methane coupling vs. pulse number during the reduction of the sample Pulse Number 1

1. Methane conversion, wt% 5.0 120 2. Calculated values of catalyst/feed mass ratio 3. Productyields, wt% C2H6 20 C2H4 10 ZC 2 30 CO 25 CO2 45 ECO + CO2 70

2

3

4

5

4.0 60

3.5 40

3.0 30

2.5 24

25 14 39 22 39 61

31 16 47 16 37 53

38 19 57 10 33 43

45 22 67 5 28 33

Table 2 Conversion and selectivity of metal coupling vs. the type of initial Mn• a TCC catalyst Sample .No Starting Conversion, Product Yield (wt%) MnxOy (wt%) 1 2 3 4

MnO

Mn203 Mn304 Mn02

3.8 4.4 5.6 6.0

introduced into

C2H 6

C2H 4

CO + CO 2

40.2 40.5 41.6 39.8

35.6 38.0 40.4 41.6

24.2 21.5 18.0 18.6

The investigation of the influence of the oxidation level of starting Mn• oxides introduced into TCC catalysts on the yield of methane conversion products (in the absence of oxidants in the feed) was performed in a plug flow microreactor at 973 K (Table 2). From Table 2 it follows that the methane conversion and ethylene/ethane ratio increase, and the yield of the complete oxidation products decreases with the growth of the oxidation level of the starting oxides Mn• The examination of TCC catalyst samples by EPR-method before and after the performance of the methane conversion reaction (reaction/regeneration cycle) disclosed two major signal types (g-factor = 2.00), which do not change their parameters with registration temperature changes (Table 3). As is evident from Table 3, the distinctions between the EPR signals (AH) are connected with only the redox properties of the starting MnxOy oxides introduced into the TCC catalysts. Changes in the integral density of the EPR signals (Mn 2§ ion concentration) in oxidation-reduction cycle point at MnxOy valency variation in the TCC catalyst. Hence the Mn• oxide is the carrier of the reagent (oxygen in the present case) in the reaction-regeneration cycle, which provides controlled selective

734 oxidation of methane into higher hydrocarbons. However, industrial realization of such a process is not promising due to low methane conversion. It seems more interesting to utilize the discovered mechanisms of methane conversion into higher hydrocarbons for catalytic cracking of methane into carbon and hydrogen. Table 3 Parameters of EPR spectra before and after methane conversion reaction vs. the type of initial Mn• v introduced into TCC catalysts Sample .No Initial MnxOy Oxidized samples Reduced samples AH (G) EPR signal AH (G) EPR signal intensity intensity (Mn+2* 102~ MnO 215 6.4 360 46 Mn203 207 20 370 75 Mn304 203 33 437 51 MnO2 200 34 380 130

(Mn+2*lO2~

2.5 748K (Methane)

2 o

. ~ 15

O

773K (Methane)

--

798K (Methane) 823K (Methane)

Aq Fo ,

"

9

~

~

-- ---. 823K (Hydrogen)

~

~

-"

1

,-

~

o

0"

.

-

-

~

-- t - . 798K (Hydrogen)

"'" "'" ::t: ..... ~~ " - - -

-- A--. 773K (Hydrogen)

~ 9

0.5

_

,.o

9149 o'~

-- o - . 748K (Hydrogen)

0

10

20

30

Experiment time (min)

Figure 1 Kinetic characteristics of methane conversion and hydrogen yield at different temperatures of methane decomposition reaction on zeolite with 8 wt% Ni.

Kinetic characteristics of methane conversion and hydrogen yield on nickel-containing pentasil zeolite with a molar SIO2/A1203ratio of 210 are represented in Fig. 1. The distinctive feature of the methane conversion into C and O is the absence of any feed oxidation products in the gas phase. Carbon obviously forms a chemical compound with NiO on the catalyst surface (NiC•

735 The methane conversion rate increases by a factor of 4 or 5 with dcclilic in tile alumillulll content of pentasil zeolites (Fig.2). The highest methane collvelsioll (75 wl%) ,,vas obtained at a reaction te~nperature o1"843 K and a catalyst/feed ratio o1" 10. l-ligh-temperature treatment of nickel-incorporating pct~tasil zcolitcs by Ilyd~t~gcl~ deactivates them in methane conversion into carbon and hydrogen; in this case tl~e quantity o1 the absorbed oxygen in a sample decreases by 0.35 mmol/g at 77 K as compared with the oxidized catalyst.

0.8 0.7 .,lt-

9- 0.6 ---....

o

E

0.5

Pentasil (Ni 8 wt%) treated by Oxygen

0.4

Pentasil (Ni 8 wt%) treated by llydrogcn

I-.,

= o

.,..~ k,i,

to >.

- a - I n i t i a l pcntasil treated by Oxygen

= 0.3

o r

.= 0.2 to

0.1

1 0.75

i m

i m

1 . 7 5 2.75

I m

3.75

.

I

4.75

Aluminum content of pentasil framework (wt%) Figure 2. Methane conversion rate vs. the aluminum content of zeolite and kind o1' its preliminary treatment.

4. C O N C L U S I O N

The introduction of Ni-containing pentasii zeolites into TCC and FCC (Fluid Catalytic Cracking) catalyst matrices makcs it possible to use (without considerable reconstruction)the technology of the production of these catalysts and the reactor-regenerator vessels of crackling units for tile process of hydrogen generation from natural gas. For example, a regenerator with a coke burning capacity of 1400 kg/h possesses a calculated hydrogen yield o1"460 kg/l~ (the vacuum gas-oil capacity of the initial cracking unit is 50000 kg/h). A TCC unit with a vacuum

736 gas oil capacity of 300000 t/year can be partially reconstructed for the process of hydrogen obtaining from natural gas with 4000 t/year output of the target product. For comparison, 4500 to 5000 t of hydrogen a year can be produced at a reforming unit of 300000 t/year capacity.

REFERENCES

1. O.V. Krylov, Catalysis Today, 18 (1993) 209. 2. G.E Keller and M. Bhasin, J.Catal., 73 (1982) 9. 3. M.I. Levinbuk and V.M. Melnikov, 211 th National Meeting, American Chemical Society, New Orleans, (1996) 410. 4. U.V. Shumovski, React. Kinet. Calal. Lett., 21 (1983) 3.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

T h e Effect of t h e P b O L o a d i n g in Methane over PbO/SiO2 Catalysts.

the

737

Oxidative

Coupling

of

H. J. Lugo, N. Teran, L. Villasmil, G. Castillo and D. M. Finol Centro de Superficies y Cat~lisis, Facultad de Ingenieria, Universidad del Zulia, Apartado 15251, Maracaibo 4003A, Venezuela

The effects of the PbO loading and the C H 4 / O 2 ratio on the oxidative coupling of methane (OCM) over PbO/SiO2 catalysts were studied. Special emphasis was made in the interpretation of the product distribution in the gas phase and its relationship with the nature of the catalyst surface. At CH4/O2 > 2 ratio, the catalytic behavior of the 2 and 6 % PbO/SiO2 catalysts was very similar. The activity for the deep oxidation of CH4 to CO2 and the activity for the formation of C2 hydrocarbons were almost the same in both catalysts. At CH4/O2 = 1 ratio, the behavior of the 2 and 6 % PbO/SiO2 catalysts differed. There were a lower activity for the oxidation of CH4 to CO2 and a higher activity for the formation of C2 hydrocarbons in the 2 % PbO/SiO2 catalyst than in the 6 % PbO/SiO2 catalyst. At any C H 4 / O 2 ratio, the 10 % PbO/SiO2 catalyst was very different from the 2 and 6 % PbO/SiO2 catalysts. It had a very low activity for the oxidation of CH4 to CO2 and was selective for the generation of C2 hydrocarbons. By doubling the a m o u n t of catalyst, the amount of reacted methane doubled, while the selectivity remained almost constant. Temperature programmed reduction experiments showed almost the same behavior in all the catalysts. However, the 10 % PbO/SiO2 catalyst showed clearly reducible species at about 670 K, which were practically absent in the 2 and 6 % PbO/SiO2 catalysts. Probably, these species were responsible for the low activity for the oxidation of CH4 to CO2 in the 10 % PbO/SiO2 catalyst.

1. INTRODUCTION

The last decade has witnessed great efforts by scientists from many countries to convert methane to value added products. The pioneering work of Keller and Bhasin [1] stimulated great interest in the oxidative coupling of methane. Redox-types oxides constitute a category of catalysts extensively studied for this purpose. Within this category, lead oxide and supported lead oxide

738 have been reported to be suitable for the OCM reaction [2-7]. These studies have demonstrated that supported lead oxide catalysts are very active and that their catalytic behavior depends strongly on the lead oxide loading; however, a clear u n d e r s t a n d i n g has not been reached yet. With respect to the mechanism, Bytin and Baerns [8] distinguished two adsorption steps for m e t h a n e on lead oxide: (a) dissociative adsorption with the s u b s e q u e n t recombination of the adsorbed fragments to yield ethane, and (b} adsorption as methylcarbonium species on acid sites, which then undergo attack by surface O"- ions, yielding methoxide species, which then undergo deep oxidation. In this study, we tried to explain the effect of the PbO loading and the CH4/O2 ratio on the oxidative coupling of m e t h a n e (OCM) over PbO/SiO2 catalysts. Special emphasis was made in the interpretation of the product distribution in the gas phase and its relationship with the n a t u r e of the catalyst surface. The methodology of the investigation involves incorporation of variable a m o u n t s of PbO to the SiO2 to examine the product distribution for both low and high lead oxide loadings. Temperature-programmed reduction (TPR) reveals the surface changes of PbO.

2. EXPERIMENTAL 2.1. Catalyst preparation Pb(NO3)2 (99.101%) and SiO2 (Davisil, grade 646, SB~=245 m2/g) were p u r c h a s e d from Riedel and Fisher, respectively. Before preparation of the catalysts, the silica was calcined at 1200 K for 4 h. PbO/SiO2 catalysts were prepared by impregnating a m o r p h o u s SiO2 (60-80 mesh) with aqueous solutions of Pb(NO3)2 of appropriate concentration to yield Pb0 loadings of 2, 6 and 10 wt. %. Excess water was removed in a rotary evaporator. The catalysts were then dried at 393 K in an oven for 12 h and subsequently calcined at 1073 K for 4 h. 2.2. Reaction s y s t e m Methane conversion was performed in a conventional ruxed-bed continuous flow reactor operated u n d e r atmospheric pressure. The reactor consisted of a quartz, U type tube of 9 m m internal diameter. The a m o u n t of catalyst used for a test r u n was about 0.25 g, which was held in place by quartz wool plugs. The reactor was placed in an electric furnace with approximately 20 cm of the quartz-filled tube serving as a preheater. Before the reaction the catalysts were pretreated in an oxygen flow at 1048 K for 1 h. The reactant mixture of CH4 and 5% 02 in He was adjusted to meet several CH4/O2 ratios and a total flow rate of about 1.2 dm3/h, keeping constant the oxygen

739 partial p r e s s u r e (4.5 kPa). The coupling reaction w a s carried o u t at 1048 K for at least three h o u r s to establish steady state conditions. Catalyst deactivation w a s not observed over this time period. The r e a c t a n t s a n d p r o d u c t s were analyzed with a n o n - s t r e a m gas c h r o m a t o g r a p h equipped with a TCD. Two c o l u m n s , a C h r o m o s o r b 102 (3 m), a n d a Molecular Sieve 5A (2.5 m) were employed in the analyses. Care w a s t a k e n to avoid c o n d e n s a t i o n of the p r o d u c t s at the outlet of the reactor. The conversion a n d selectivities were calculated from the a m o u n t s of reaction p r o d u c t s formed (carbon a t o m basis) as d e t e r m i n e d by the GC analysis. The error in c a r b o n balance w a s found to be below 5% in all cases. Total conversion of r e a c t a n t (XT) a n d selectivity to p r o d u c t i (Si) are defined as;

XT --'--

Si =

moles of reactant transformed X 100 moles of reactant in the feed

C atoms of i 9moles formed of i x 100 moles of CH4 transformed

2.3. Experimental

techniques

The surface a r e a s of the catalysts were m e a s u r e d by the conventional BET nitrogen a d s o r p t i o n method. Values of 26, 7 a n d 4 m 2 / g were o b t a i n e d for the 2, 6 a n d 10 % PbO/SiO2 catalysts, respectively. T e m p e r a t u r e p r o g r a m m e d reduction e x p e r i m e n t s were performed u s i n g an a p p a r a t u s described by Robertson et al. [9]. The reduction w a s carried o u t with a purified h y d r o g e n - a r g o n mixture (10 vol.% hydrogen) at a h e a t i n g rate = 10 K min -I u p to 1048 K. The TPR reactor w a s charged with 0.25 g of c a l c i n e d (fresh) catalyst. Before the reduction, the catalyst w a s oxidized in a 02 flow to 1048 K for 1 h, a n d t h e n cooled down to 300 K in Ar. This reduction-oxidation cycle w a s repeated several times.

3. RESULTS AND DISCUSSION

Table 1 s h o w s the catalytic properties for the oxidative coupling of m e t h a n e over 2, 6 a n d 10 % PbO/SiO2 catalysts at different CH4/O2 ratios at 1048 K, u s i n g 250 mg of catalyst. It w a s also r u n u s i n g 500 mg of 10 % PbO/SiO2 catalyst at a CH4/O2 = 2 ratio. As it is well k n o w n on OCM catalysts, the CH4 conversion decreases a n d the C2 selectivity increases w h e n the CH4/O2 ratio increases. For the 2 a n d 6 % PbO/SiO2 catalysts,

740 the 02 conversion is less t h a n 100 % only for a CH4/O2 = 1 ratio, while for the 10% PbO/SiO2 catalyst, it is always well below 100 %. Besides, by doubling the a m o u n t of the 10% PbO/SiO2 catalyst, the conversion of m e t h a n e almost doubles, while the selectivity remains nearly constant. However, Table 1 does not provide sufficient information a b o u t the behavior of these catalysts. Therefore it is necessary to look further into the feed and p ro du c t gas composition.

Table 1 C a t a l ~ i c Properties of PbO/SiO2 catalysts PbO load P(CH4) Conv CH4 Conv 02 /P{O2)

(%1

(O/o)

(o/o)

C2H6

Selectivity

(%)

C2H4

CO

CO=

2 (250 mg)

1 2 5

39.3 29.0 15.6

65.4 100 100

7.9 9.9 13.3

11.1 15.8 25.0

5.9 4.8 4.2

75.1 69.5 57.5

6 (250 mg)

1 2 5

48.8 30.4 16.4

86.4 100 100

5.8 13.2 14.3

7.8 17.5 25.4

2.3 4.2 5.3

84.1 65.1 55.0

10 (250 mg)

1 2 5

8.2 6.4 3.6

6.5 7.4 12.0

46.7 53.9 60.7

19.1 22.5 19.7

8.7 4.7 3.6

25.5 19.0 16.0

10 (500 mg)

2

11.3

17.3

44.1

25.6

5.2

25.1

Reaction conditions: F = 1.2 din3/h, P(O2)ffi 4.5 kPa, Ptot~ffi 101 kPa, T = 1048 K, Inert = helium.

Figure 1 shows the c o n s u m p t i o n of CH4 and 02, and the a m o u n t of the different p r o d u c t s formed for several CH4/O2 ratios over a 6 % PbO/SiO2 c a t a l y s t . For a CH4/O2 = 2 ratio, the c o n s u m p t i o n of CH4 reaches a value n e a r or equal to the m a x i m u m a m o u n t allowed by the a m o u n t of oxygen available for the reaction. At this point all of the available oxygen is c o n s u m e d (100% O2 conversionl. For a CH4/O2 = 1 ratio, the c o n s u m p t i o n of CH4 does not reach the m a x i m u m allowed value (86.4 % 02 conversion). This is because with a decreased a m o u n t of m et hane, and a c o n s t a n t

741 a m o u n t of oxygen, competition for the active sites o c c u r s where the m e t h a n e is disfavored, so t h a t the m e t h a n e c o n s u m p t i o n decreases. This competition between m e t h a n e a n d oxygen h a s been noted in the literature [10]. For a CH4/O2 --- 5 ratio, the excess m e t h a n e displaces the oxygen on the active sites. It i n c r e a s e s the m e t h a n e c o n s u m p t i o n until there is no more oxygen available (100% 02 conversion) to regenerate the active sites, resulting in c o n s t a n t a m o u n t of c o n s u m e d m e t h a n e at the steady state. In this situation, the n u m b e r of active sites utilized can be less or equal to the total n u m b e r of active sites. Figure 1 also shows t h a t w h e n m e t h a n e is in excess with respect to oxygen, t h a t is, at higher CH4/O2 ratios, the oxidative coupling of m e t h a n e is favored over total c o m b u s t i o n . It is evidenced by the increasing p r o d u c t i o n of C2 h y d r o c a r b o n s a n d the decreasing a m o u n t of CO2. At these conditions, the m e t h a n e m u s t occupy a higher proportion of the surface, limiting the a m o u n t of oxygen t h a t h a s access to it. This situation favors a methyl radical p r o d u c t i o n a n d their coupling in the gas p h a s e a n d disfavors the oxidation of CH4 to CO2 on the surface. When m e t h a n e is in deficit, t h a t is, CH4/O2 < 1 ratio, the excess surface oxygen restricts the formation of m e t h y l radicals. Rather, m e t h a n e is deeply oxidized, favoring the p r o d u c t i o n of CO2.

o

3

D

2,5

uJ O am

--o--02

o

-r -o-- C02

n,,o O(9 1,5

a

o----o.-._...._

~

~-. C2H4 (*)

I--

n,,l~ w > =-0,5 z

o0

C2H6 (*) --

0 0

1

2

3

4

5

CO

6

CH4/O2 RATIO

Figure i. Effect of the CH4/O2 ratio on the gas p h a s e composition in a 6 % PbO/SiO2 catalyst. (*) Equivalent ~mol of CH4 = C atoms of i product. #Jmol formed of i product

742 Figure 2 shows the consumption of CH+ and 02, and the a m o u n t of several p r o d u c t s versus the CH4/O2 ratio for a 2 % PbO/SiO2 catalyst. The behavior of this catalyst is the same as the 6 % PbO/SiO2 catalyst for CH4/O2 > 2 ratio. The activity for the deep oxidation of CH+ to CO2 and the activity for the formation of C2 hydrocarbons are almost the same in both catalysts.

uJ

3

o

2,5

o

2

.--0--02 --r 9

CH4

--o--- CO2 x

Z

o

(3

C2H4 (*)

--b-- C2H6 (*)

0,s

--

0 0

1

2

3

4

5

CO

6

CH4/O2 RATIO

Figure 2. Effect of the CH4/O2 ratio on the gas p h a s e composition in a 2 % PbO/SiO2 catalyst. (*) Equivalent pmol of CH, = C atoms of i product, pmol formed of i product For a CH+/O2 = 1 ratio, the behavior in both catalysts differs slightly. There are a lower activity for the oxidation of CH4 to CO2 and a higher activity for the formation of C2 hydrocarbons in the 2 % PbO/SiO2 catalyst t h a n in the 6 % PbO/SiO2 catalyst. For this ratio, the adsorption of oxygen competes with that of methane, and results in a lower m e t h a n e c o n s u m p t i o n t h a n in the case of a CH4/O2 = 2 ratio. The lower consumption of m e t h a n e in the 2 % t h a n in the 6 % PbO/SiO2 catalyst at CH4/O2 = 1 ratio, corresponds to a decrease in the conversion of m e t h a n e to CO2 in the 2 % PbO/SiO2 catalyst, and an increase in the conversion of m e t h a n e to C~ hydrocarbons. The decrease in the production of CO2 a n d the increase in the production of C2 h y d r o c a r b o n s would be related to a lower n u m b e r of active sites for the deep oxidation of m e t h a n e a n d to a higher n u m b e r of active sites for the formation of methyl radicals in the 2 % PbO/Si02 catalyst. Figure 3 shows the consumption of CH4 and 02, and the a m o u n t of p r o d u c t s versus the CH4/02 ratio for the 10 % PbO/SiO2 catalyst. It is

743 observed that the oxygen consumption is very low, so there is excess of oxygen in the gas phase at any time, and oxygen suppy is not limiting the rate. The consumption of methane increases with the a m o u n t of methane in the gas phase. This consumption is strongly related to the production of C2 hydrocarbons and to the active sites able to generate methyl radicals, while the low production of CO2 suggests a shortage of active sites for the oxidation of methane to CO2.

0,5 o LU o :::) am O r,r n,'o 0(.9

uJ~

0,4

---o-- CH4 ---0.-02

0,3

C2H6 (*) 7. C2H4 (*)

0,2

IuJ :~. > 0,1 z O o

-o--CO2 #, CO

0 0

1

2

3

4

5

6

CH4/O2 RATIO

Figure 3. Effect of the CH4/O2 ratio on the gas phase composition in a 10% PbO/SiO2 catalyst. (*) Equivalent #Jmol of CH, = C atoms of i product, pmol formed of i product Evidently the I0 % PbO/SiO2 catalyst has a different nature than the two lower loading s a m p l e s , in the sense that it has something that inhibits the deep oxidation of methane, increasing the generation of methyl radicals and so the production of C2 hydrocarbons. Figure 4(a} compares the production of ethane, ethylene and carbon monoxide from the various catalysts. Figure 4(b) compares the generation of carbon dioxide from these catalysts, and shows that only the 2 % and 6 % PbO/SiO2 catalysts have high production of CO2. Clearly the 10 % PbO/SiO2 catalyst has undergone a modification in its structure, which allows lower formation of CO2.

744

,2

a uJ

o

1,2

......

1

a

o ~ . 0,8 n~

--o--- 6 % PbO

0.6

---a-- 10% PbO

~.~ +

o.j

0,4

0 o "0,2

+"0,2

0

0,8

- . - o - 2% PbO

0,4

"1r

DU)

|

0

1 2 3 4 5 6 CH4102 RATIO

0

,

|

1 2 3 4 5 CH4/O2 RATIO

(b)

(a)

Figure 4. Effect of the PbO loading in the OCM reaction over PbO/SiO2 catalysts. (*) Equivalent lJmol of CH4 = C atoms of i product, lJmol formed of i product

Ill 1 I'--

2 O E ia,.

PbO 10%

=E

6%

z

o

2%

O N

:z: 0

400

I

|

I

I

600

800

1000

1200

TEMPERATURE (K)

Figure 5. TPR profiles for PbO/SiO2 catalysts.

|

6

745 Figure 5 compares the results of the TPR experiments and shows that all the catalysts have similar behavior. However, the 10 % PbO/SiO2 catalyst shows clearly a reduction feature at about 670 K, which is absent in the 2 and 6 % PbO/SiO2 catalysts. The presence of some reducible species may be responsible for the very low activity for the deep oxidation of CH4 to CO2 in the 10 % PbO/SiO2 catalyst.

4. CONCLUSIONS

At CH4/O2 -> 2 ratio, the behavior of the 2 and 6 % PbO/SiO2 catalysts is very similar. The activity for the deep oxidation of CH4 to CO2 and the activity for the formation of methyl radicals ( C2H6, C2H4 and CO) are almost the same in both catalysts. At CH4/O2 = 1 ratio, the behavior of the 2 and 6 % PbO/SiO2 catalysts differs slightly. There is a lower activity for the oxidation of CH4 to CO2 and a higher activity for the formation of C2 hydrocarbons in the 2 % PbO/SiO2 catalyst than in the 6 % PbO/SiO2 catalyst. At any CH4/O2 ratio, the 10 % PbO/SiO2 catalyst has a very low activity for the deep oxidation of CH4 to CO2 and its catalytic activity is for the generation of methyl radicals. By doubling the a m o u n t of catalyst, the a m o u n t of reacted methane doubled, while the selectivity remained almost constant. Temperature programmed reduction experiments show almost the same behavior in all the catalysts. However, the 10 % PbO/SiO2 catalyst shows a reducible species at about 670 K, which is absent in the 2 and 6 % PbO/SiO2 catalysts. Probably, these species are responsible of the very low activity for the oxidation of CH4 to CO2 in the 10 % PbO/SiO2 catalyst.

REFERENCES

I. G.E. Keller and M.M. Bhasin, J. Catal., 73 (1982) 9. 2. W.Hinsen, W. Bytyn and M. Baerns, in Proceedings 8 th International Congress on Catalysis, Berlin, 1984, Vol. 3, Verlag Chemie, Weinheim, 1984, p.581. 3. K. Asami, S. Hashimoto, T. Shikada, K. Fujimoto and H. Tominaga, Ind. Eng. Chem. Res., 26 (1987) 7. 4. G. Wendt, C.D. Meinecke and W. Schmitz, Appl. Catal., 45 (1988) 209. 5. A. Machocki, A. Denis, T. Boroniecki and J. Barcicki, Appl. Catal., 72 ( 1991) 283. 6. S.E. Park and J. S. Chang, Appl. Catal. A 85 (1992] 117,

746 7. R. Mariscal, J. Soria, M. A. Pefia and J. L. G. Fierro, Appl. Catal., A: General 111 (1994) 79. 8. W. Bytyn and M. Baems, Appl. Catal., 28 (1986) 199. 9. S. D. Robertson, B.D. McNicol, J.H. De Bass, S. C. Kloel and J. W. Jenkins, J. Catal., 37 (1975) 424. 10. G.A. Martin and C. Mirodatos, Fuel Processing Technology, 42 (1995) 179.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

747

Catalytic Combustion of Ethane over High Surface Area Lnl.xKxMnO3 (Ln La, Nd) Perovskites: The effect of Potassium Substitution. =

Y. Ng Lee a*, F. Sapifiab, E. Martinezb, J. V. Folgado b and V. Cort6s Corberfina+. a Instituto de Cat~lisis y Petroleoquimica, CSIC, Campus U.A.M. Cantoblanco, 28049 Madrid,

Spain. b Instituto de Ciencia de los Materiales de la Universitat de Valencia, Dr. Moliner 50, 46100 Burjassot (Valencia), Spain.

Monophasic Lnl.xKxMnO3 (Ln = La, Nd) perovskites with high surface areas (8-27 m2/g) were synthesised at mild conditions by the ~eeze-drying method, and were found to be active for the catalytic combustion of ethane at low temperatures (5 73 to 648 K). As a general trend, the substitution of the rare earth cation by potassium decreased the intrinsic activity, reduced the reaction order in oxygen and, for the more substituted samples (x>0.10), it increased the selectivity to ethene. It was found that the rare earth cation also influenced the catalytic activity of the substituted perovskites. These effects were analysed in terms of structural modifications induced by the introduction of potassium in the perovskites.

1. INTRODUCTION The advantages of the catalytic (flameless) combustion of light hydrocarbons are numerous [1-3]. In the presence of active catalysts complete combustion may be achieved at temperatures several hundred degrees lower than in the flame (thermal noncatalytic) combustion [3]. This low temperature is the basic advantage of catalytic combustion. For example, noxious nitrogen oxides formation is almost eliminated [3] and energy savings are achieved. Supported platinum or palladium catalysts have so far been used almost exclusively because of their excellent activity [2-3]. However, their high price, tendency to be poisoned, and volatility in some applications, have estimulated the search for suitable substitutes. In 1971 Ll"bby [4] proposed LaCoO3 as a potential auto exhaust catalyst, starting a wide interest in the catalytic properties of perovskites. However, their specific application as catalysts in flameless combustion has been investigated only since 1980 [6-15], mainly in Japan. Nakamura et aL [6] began these studies reporting the catalytic oxidation of propane, methane and CO on several LaBO3 (B = Co, Mn, Ni, Fe), and lanthanum substituted (Ca, Sr, Ba, Ce) perovskites. The order in catalytic activity was found to be Co > Mn > Ni > Fe > Cr, On leave from: Universidad Central de Venezuela, Facultad de Ciencias, A.P. 47586, Los Chaguaramos, 1041 A, Venezuela. + Correspondingauthor. E-mail: [email protected] Carar

748 and Sr-"doping" was the most effective in enhancing the activity of these compounds. Seiyama et al. [7] reported the activity for the catalytic oxidation of propylene of several perovskite oxides (ABO3) and their corresponding B oxides (BO~), showing that their activity was determined mainly by the nature of the metal in B positions. The most active catalysts were Co and Mn perovskites. Arai et aL [8] investigated the catalytic activity of various transition-metal perovskites for methane combustion. In every case, carbon monoxide or partial oxidation products were scarcely detected during oxidation. The activities of LaCoO3, LaMnO3 and LaFeO3 were quite close to that ofa Pt/ahxmina catalyst. An interesting feature of rare-earth perovskites is the possl~oility to vary the unit cell dimensions by controlling the nature of the A ion, and thereby the covalence of the B-O bond. A thorough study of the role of the A- and B-site ions on the catalytic properties of ABO3 perovskites for propane and methanol oxidation has been reported by Nitadori et aL [9]. They concluded that the influence of the rare-earth ions in the A-site on the oxidation properties of these compounds was secondary. Zhang et aL [10] studied the oxygen sorption and catalytic properties for methane and n-butane combustion, of Lal.xSrxCo~.yFeyO3. Whil~ the catalytic activity for n-butane oxidation was affected both by the transition-metal substitution and rareearth substitution, the catalytic activity for methane oxidation was only influenced by the rareearth ion substitution. Marti and Balker [ 16] have found that the activity of ACoO3 oxides for methane oxidation was, with exception of the PrCoO3 perovskite, only slightly influenced by the A-site cations. Perovskite-like compounds derived from laflVlnO3 by partial substitution of La3+ by divalent ions (Sr, Ca, Ba, Ce) are well known. However, the effect of partial substitution of the A site by alkaline ions on the catalytic activity in combustion reactions has been very little studied. Voorhoeve et aL [ 17] reported for the first time that the substitution of La 3+ by alkali ions in LaMnO3 increases its catalytic activity for NO reduction. Apart from their intrinsic activity, the performance of perovskite catalysts depends on their particle morphology, and, in particular on their specific surface area (SSA) [5, 12, 18]. To obtain homogeneous perovskite ~mples, the conventional ceramic route requires repetitive grinding and heating procedures as well as high temperatures, which results in large grain sizes and low SSA (< 2 m2/g), unsuitable for their application as catalysts. The use of alternative synthetic pathways produces high purity, homogeneous powders, and requires low temperatures to obtain the phases, thus leading to products with smaller particles and high SSA [19]. Besides these advantages, these alternative methods play an essential role in the preparation of potassium containing materials, due to the volat'flity of potassium oxide at high temperatures. In the alternative method involving freeze-drying all the cations are mixed at the atomic scale in the solution and this facilitates the incorporation of potassium to the perovskite lattice since it occurs at the initial stage of the preparative procedure [20]. Freeze-drying has proved to be a powerful and versatile technique for mild preparation of complex oxides, such as high temperature superconducting material and other perovskites (lanthanide, nickelates and cobaltates) [21, 22]. We report in this work the low temperature synthesis and catalytic properties of two series ofLal.xKxMnO3 (Ln = La, Nd; 0 _<x _<0.25) perovskites. We have investigated the effects of lanthanide and potassium substitution of the A site-cations, and of the oxygen partial pressure (Po2), on their catalytic activity for ethane combustion at low temperatures (573 to 648 K). The use of a precursor-based synthetic

749 method, namely, the freeze-drying of acetate solutions, allows the preparation of monophasic perovskites at temperatures as low as 873 K (973 K for Nd samples).

2. EXPERIMENTAL 2.1. Catalysts p r e p a r a t i o n

Aqueous solutions of metal acetates with molar nominal compositions: Ln : K : M n = 1-x : x : 1 (Ln = La, Nd) with x = 0.00, 0.05, 0.10, 0.15, 0.20 and 0.25, were prepared as follows. KHCO3 was dissolved in 100 ml of glacial acetic acid. Addition of Ln203 lead to a suspension, which was gently heated while stirring for 15 minutes. After addition of 20 ml of H20, a transparent solution resulted. After cooling, Mn(CH3COO)2.4 H20 was added and dissolved while stirring. Droplets of the resulting pink (or lilac with Nd) pale acetic solutions were flash frozen by pouring them into liquid nitrogen and, then, freeze-dried at a pressure of 10-3 Pa. In this way solid precursors were obtained as pink (or lilac)-coloured, dried and loose powders. Perovskites were obtained from these powders by calcination at 873 K (or 973 K for samples with Nd) for 12 h (heating rate: 5 K/rain, cooling rate 80 K/rain) under flowing oxygen. These calcination temperatures allowed the formation of monophasic perovskite samples for each series. Samples will be denoted hereinafter as AKx (A = L, N) where A is the rare earth cation (L for La, N for Nd) in the A site and x the degree of substitution by potassiun~ 2.2. P h y s i c o c h e m i c a l c h a r a c t e r i s a t i o n

Phase identification of the catalysts was carded out by powder X-ray ditfractometry using a Siemens DS01 computer-controlled diifractometer, with monochromatized Cu Kcx radiation. Patterns were registered in steps of 0.08 ~, in the angular range 20 to 65 ~ 20, with integration times of 5 seconds per step. The samples were dusted through a sieve onto the holder surface to reduce preferred orientation. The obtained patterns were compared with JCPDS data files. Physisorption measurements were performed with a Micromeritics ASAP 2000 instrument. The BET surface areas were determined by nitrogen adsorption at 77 K assuming a crosssectional area of 0.162 nm 2 for the nitrogen molecule. Before the adsorption measurements the samples were outgassed in vacuum at 423 K for 18 h. The morphology evolution was followed by Scanning Electron Microscopy (Field Emission) in a I-fitachi 4100 instrument. 2.3. Catalytic tests

The catalytic oxidation of ethane at 573-648 K was carded out at atmospheric pressure in a fixed bed flow reactor. Mixtures of ethane (4 mol%), oxygen (4-12 mol%), and helimn (balance) were fed to the reactor with a residence time of 38 g . h/mol C2H6, using a catalyst load of 0.36 g, (particle size 0.25-0.42 ram) mixed with SiC bits (dilution 1:4 v/v) to reduce the heat release per unit volume. Reactants and products were analysed by gas chromatography on a Varian 3400, equipped with a thermal conductivity detector, using Porapak QS (3 m) and molecular sieve 13X (1 m) cohnnns. In all reaction conditions, the mass and carbon balances were within 100+_2 %.

750 3. RESULTS AND DISCUSSION 3.1. X-ray Diffraction (XRD) The formation of the perovskite phase was confirmed by XRD. The patterns of the Lal.~ KxlVinO3 and Nd~.~KxMnO3 samples are shown in Fig. 1. The reflections of the unsubstituted LK0.0 sample agreed with those of the non-stoichiometric phase LaMnO3.~5 (JCPDS data file 32-0484). Wold and Amott [23] reported an oxidative non-stoichiometry in LaMnO3.8 varying between 6 = 0 and 6 = 0.15. They also found a change in the symmetry of LaMnO3+8 from orthorhombic to rhombohedral when the non-stoichiometry (5) exceeded 0.105. The XRD pattern of the unsubstituted NK0.0 sample agrees with those ofNdMnO3 (JCPDS data file (250565). The progressive substitution of lanthanum or neodymium by potassium did not modify the patterns: no peaks due to the component single oxides or any other secondary known phase were detected. In fact, very weak peaks corresponding to mixed oxides of Mn and K could be detected only for samples with x > 0.15 when calcined at much higher temperatures (1273 K) than those used for catalyst preparation. 3.2. Surface areas

BET specific surfaces areas of the catalysts are given in Table 1. The lower values for the Ndl.xKxMnO3 series are due to the higher calcination temperature used (973 K), needed to obtain monophasic samples. In general, these values are relatively high. Thus, SSA values obtained for the series Lal.xK~MnO3are similar for those reported for La~.xSrxMnO3 prepared by the sol-gel method (17.5-23.5 m2/g) [24] and LaMno.sCt~503 (27 m2/g) prepared by t~eeze-drying [25], but much higher than that ofLao 95K0.osM O3 prepared by the ceramic

NKx

.......

z

20

3'o

4'0

s'o

2e (degrees)

6'o

20

3'o

4'0

20 (degrees)

Figure 1. XRD patterns of LKx (left) and NKx (right) catalysts.

s'o

6'o

751 Table 1 BET Specific Surface Area# ofLn~.xK~MnO3 (Ln = La, Nd). Catalysts x Specific Surface Area (m2/g) Lal.xKxMnO3 Ndl_xKxMnO3 0 19.7 19.0 0.05 24.3 14.8 0.10 26.5 10.0 0.15 24.4 9.6 0.20 21.7 8.0 0.25 20.8 10.6 From N2 adsorption at 77 K. Sm= 0.162 nm2. method (0.62 m2/g) [17]. The SSA of NdMnO3 (NK0.0) is also higher than that of NdCoO3 (12.3 m2/g) obtained by fleeze-drying of amorphous nitrate precursors [21], which was found to be the highest reported in the literature for the perovskite.

3.3. Scanning Electron Microscopy (Fidd Emission) The morphology evolution has been followed by Scanning Electron Microscopy Field Emission (SEM-FE). Micrographs shown in Fig. 2 correspond to two representative samples: sample Lao.ss~.tsMnO3 prepared at 873 K (catalyst LK0.15) and sample Ndo.ssKo.~sMnO3 prepared at 973 K (catalyst NK0.15). No segregation of secondary phases can be observed in the micrographs. No variation of morphology and particle size due to composition was observed within the homologous series for this temperature.

Fig.

2.

SEM-FE

micrographs

of

samples

LK0.15

(A)

and

NK0.15

(B).

752 SEM-FE images show the existence of a noticeable difference in the dism'bution of particle size between the LKx and NKx samples. The lanthanum sample (LK0.15, monophasic by XRD), is composed of homogeneous nanometer particles, with a narrow particle size distEoution (30-50 nm), while the Nd sample shows a broader distn'bution (20-100 nm). Nevertheless, the average particle size of both series was similar, about 40 nm In both La and Nd materials the particles were rounded, and non-faceted. Only when ~mples were treated at high temperature (1173 K) did they exls an increase in particle size, accompanied by the development of crystalline facets. The particle size range of La samples remained narrower than that of Nd samples (50-200 nm~ 60-300 nm~ respectively). 3.4. Catalytic Activity In the conditions used, all the catalysts were active for the combustion of ethane, producing conversions ~om 1-5 % at 573 K up to 40 % at 648 IL The only detected products were CO2, H20 and small amounts of ethene, that is consistent with a triangular reaction scheme involving parallel formation of CO2 and ethene and the consecutive combustion of ethene. In separate measurements of CO oxidation on these catalysts, made in a FTIR cell [26], CO was transformed completely into CO2 at temperatures as low as 573 K. This would explain the absence of CO among the reaction products of ethane oxidation, carried out here at higher temperatures. Fig. 3 shows the variation of areal rates of ethane oxidation over the LKx and NKx catalysts with temperature. In all cases, ethane conversion increased with increasing

x I

t

E450 /

O=! 0".I

INKx[

x 0 0.05 O 0.00

[] o.ool

0.10

(3//

0.15

;" ".~ o 2o

0 575

600 625 650 675 Temperature (K)

0.20

5"/5 600 6~ 6~ Temperature (K)

675

Figure 3. Intrinsic rates of ethane oxidation on Lai.xKxMnO3 (leit) and Ndl.xK~MnO3 (right) catalysts as a function of temperature and degree of substitution (x). Feed composition: 4% C2I~, 12% 02, He balance; W/F: 38 g. h/mol C2I-I6.

753 2O

INI 15

/

-I-

ID

.r. ,t

o

-o 0

0

10

20

30

C2H6 Conversion (%)

40

0

/

10

I

20

i

I

30

C2H 6 Conversion

40

(%)

Figure 4. Variation of ethene selectivity with ethane conversion on Lal.xKxMnO3 (left) and Ndl.xKxMnO3 (right) catalysts. Experimental conditions and symbols as in Fig. 3. temperature, but in both series the intrinsic activity decreased with increase in potassium substitution (x), with the only exception of catalyst NK0.05, where the areal rate was higher than that of the unsubstituted NdMnO3. Within experimental error, the calculated values of the apparent activation energy for CO2 formation were the same, 18.4_+0.2 kcal/mol, for all the LKx catalysts, while they were lower, 15.8_+0.9 kcal/mol, for the NKx catalysts. Another interesting effect of potassium substitution can be seen in Fig. 4 which shows the evolution of ethene selectivity as a function of ethane conversion in both series. In general, as potassium substitution increased, the ethene selectivity for a given conversion also increased. For the unsubstituted samples (x = 0.0) and for x = 0.05, CzI-I4selectivity remained constant c a . 2%, in the whole temperature range. However, in the case of catalysts with higher substitution degrees (x > 0.10), this selectivity tends to increase with increasing x and also with increasing conversion, with the later effect much more marked in the Nd-containing samples. This indicates that the presence of potassium favours either the formation and/or the desorption of ethene formed via oxidehydrogenation (OXD) of ethane. The effect of the cation in the A-position can be more clearly seen in Fig. 5, in which the areal rates of ethane transformation as a function of x are compared. The activity of the unsubstituted samples was very similar, being slightly higher for the La catalyst. By contrast, the ethane combustion rate was higher on the substituted Nd~.xKxMnO3 perovskites as compared with the homologous La~.~K~MnO3 samples with the same substitution level. This reveals that the nature of the rare-earth cation also has a si~ificant effect in the catalytic performance, especially in the presence of potassium The effect of oxygen partial pressure (Po2) on the areal rates of ethane transformation at 648 K is shown in Figure 6. In the samples without potassium or with low substitution, the activity increased as Po2 increased. However, in the samples with x = 0.20=0.25, only little changes in the intrinsic rate of ethane conversion could be observed when Po2 increased. Reaction orders

754

600

- - [ 3 - - LKx

N"

- - O - - NKx 450

r| O

E

300

ID

150 < 0

I

I

0.00

I

I

I

I

0 . 0 5 0 . 1 0 0 . 1 5 0 . 2 0 0.25 Substitution x

Figure 5. Effect ofpotassium on the areal rate. Temperature: 648 K, Po2" 1.1 kPa.

E '7" e"

jo O

45O

=L v

m

3OO V

150

+__._+_____+_ I

3

6

~

V

+

;

O2 Pressure 0oal

12

3

~

~

12

0 2 Pressure (kPa)

Figure 6. Ethane conversion areal rates at 648 K as a fimction of Po2. W/F: 38 g . h/mol C2I-I6.. Symbols as in Fig. 3. in oxygen for catalysts with x = 0-0.20 are summarised in Table 2. It can be seen that, within experimental error, the reaction order in oxygen decreased monotonically with increasing K substitution, while being practically equal for the homologous LKx and NKx samples with the same x. This may be attn~buted either to a change in the mob'dity of the oxygen species, or to a change in oxygen species that participate in the reaction, specifically, from molecularly

755 Table 2. Dependence of ethane oxidation rate on Llll.xKxlVInO3 ( L n = La, Nd). catalysts on the partial pressure of oxygen x Reaction order n Lal.xKxMnO3 Ndl.xKxMnO3 0 0.44 + 0.11 0.42 + 0.02 0.05 0.38 + 0.06 0.31 + 0.02 0.10 0.35 + 0.07 0.25 + 0.05 0.15 0.23 + 0.03 0.17 + 0.03 0.20 0.15 + 0.04 0.29 + 0.16 a

Oxidation rate, r = k (Pc2n6) ~ (Po2)"

adsorbed to lattice oxygen [8, 12]. Neutron di~action studies of LaMnO3.15 [27] are consistent with a structure consisting of a compact packing of oxygen ions with cationic vacancies at both La and Mn positions. The substitution of one trivalent rare earth ion by K+ will reduce these vacancies, that become filled for x = 0.15, thus reducing the mobility of oxygen species.. On the other hand, the promoting effect of potassium doping in selective oxidation reactions is well known. So, addition to metal oxide catalysts, such as V2Os/TiO2 and MoO3/TiO2 increases the selectivity to oxidehydrogenation (OXD) of propane [28]. It has been shown that potassium addition brings about a decrease in acidity, lowers surface potential (work function) and hinders the formation of electrophilic O species, which are the responsible for total combustion. Thus, the reduction of total conversion as well as the increase in the selectivity to OXD products when the potassium content increases could be interpreted as due to the modifications of these properties induced by the presence ofpotassiunt

CONCLUSIONS The freeze-drying of acetates solutions is a suitable method for preparing high SSA monophasic Lnl.xKxMnO3 (In = La, Nd) perovskites. The presence of potassium strongly affects the catalytic properties, probably by modifying the nature and concentration of surface species on the catalyst surface. The increase of K content would decrease the ability of the perovskite to activate 02 and to store active oxygen species, and would reduce the electrophilicity of these species, thus increasing the C2I-I6selectivity.

Acknowledgements The authors thank Dr. Rochel M. Lago for useful discussions. Y. Ng Lee was supported by a Fellowship from CDCH de la UCV, Venezuela. Financial support from the Spanish CICYT under project MAT96-0688-C02-02 is acknowledged.

756 REFERENCES

[ 1] G. E. Voecks, 3rd Workshop on Catalytic Combustion (1977). [2] R. Prasad, L. A. Kennedy and E. Ruckenstein, Catal. Rev.-Sci. Eng., 26 (1984) 1. [3] D. L. Trimm~ Appl.Catak, 7 (1984) 249. [4] W. F. L~by, Science, 171 (1971) 499. [5] N. Yamazoe and Y. Teraoka, Catal. Today, 8 (1990) 175. [6] T. Nakamura, M. Misono, T. Uchijima and Y. Yoneda, Nippon Kagaku Kaishi (1980) 679. [7] T. Seiyama, N. Yamazoe and K. Eguchi, Ind. Eng. Chem_, Prod. Res. Dev., 24 (1985) 19. [8] I-1.Arai, T. Yamada, IC Eguchi and T. Seiyama, AppL Catal., 26 (1986) 265. [9] T. Nitadori, T. Ichiki and M. Misono, Bull Chem_ Soc. Japan, 61 (1988) 621. [10] H. M. Zhang, Y. Shimizu, Y. Teraoka, N. M ~ a and N. Yamazoe, J. Catal., 121 (1990) 432. [11] Z. Kaiji, L. Jian, and B. Yingli, Catal. Lett., 1 (1988) 299. [ 12] L. G. Tejuca, J. L. G. Fierro and J. M. D. Tasc6n, in: Advances in Catalysis, Vol. 36, eds. D. D. Eley, H. Pines and P. B. Weisz (Academic Press, New York, 1989) p. 237. [13] J. G. McCarty and H. Wise, Catal. Today, 8 (1990) 213. [14] B. de Collongue, E. Garbowsky and H. Primet, J. Chem_ Sot., Faraday Trans., 87 (1991) 2493. [ 15] C. B. Alcock and J. J. Carberry, Solid State Ionics, 50 (1992) 197. [16] P. E. Marti and A. Baiker, Catal. Lea., 26 (1994) 71. [ 17] R. J. H. Voorhoeve, J. P. Remeika, L. E. Trimble, A. S. Cooper, F. J. Disalvo and P. K. CraHagher, J. Solid State Chem_, 14 (1975), 395. [18] R. J. I-1. Voorhoeve, D. W. Johnson Jr., J. P. Remeika and P. K. Gallagher, Science, 195 (1977) 4281. [19] J. Kirchnerova, D. Klvana, J. Vaillancourt and J. Chaouki, Catal. Lett., 21 (1993) 77. [20] Y. Ng Lee, F. Sapifia, E. Martinez, J. V. Folgado, R. Ibafiez, D. Beltrfin, F. Lloret and A. Segura, J. Mater. Chem_, submitted. [21] A. Gonzhlez, E. Martinez Tamayo, A. Beltrhn Porter, V. Cortes Corberhn, CataL Today, 33 (1997) 361. [22] V. Primo, F. Sapifia, M. J. Sanchis, R. lbafiez, A. Beltrhn, D. Beltrhn, Solid State Ionics, 63 (1993) 872. [23] A. Wold and R. J. Amott, J. Phys. Chem_ Solids, 9 (1959) 176. [24] I-I. Taguchi, D. Matsuda, M. Nagao, K. Tanihata and Y. Miyamoto, J. Ant Ceram. Sot., 75 (1992) 201. [25] D. W. Johnson, P. K. Gallagher, F. Schrey and W. W. Rhodes, Ant Ceram. Soc. Bull., 55 (1976) 520. [26] Y. Ng Lee, F. Sapifia, E. Martlnez, J. V. Folgado, V. Cort6s Corberfin, in preparation [27] J. A_ N. van Roosmalen, E. I-1. P. Cordfunke, R. B. Helmholdt and H.W. Zandbergen, J. Solid State Chem_, 110 (1994) 100. [28] B. Grzybowska, P. Mekss, R. Grabowski, K. Wcislo, Y. Barbaux, and L. Gengembre, in "New Developments in Selective Oxidation IF' (V. Cortes Corberfin and S. Vie Bell6n, Eds.). Elsevier, Am~erdam 1994, Studies in Surf. Sci. Catal., 82 (1994) 151.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

757

Effect o f R e d o x T r e a t m e n t o n M e t h a n e O x i d a t i o n o v e r B i n a r y C a t a l y s t . Yu. P. Tulenin, M. Yu. Sinev, V. V. Savkin, and V. N. Korchak Semenov Institute of Chemical Physics, Russian Academy of Sciences 4 Kosygin street, Moscow 117334, Russia The analysis of critical phenomena, such as hysteresis and self-oscillations, gives valuable information about the intrinsic mechanism of catalytic reactions [1,2]. Recently we have observed a synergistic behavior and kinetic oscillations during methane oxidation in a binary catalytic bed containing oxide and metal components [3]. Whereas the oxide component (10% Nd/MgO) itself is very efficient as a catalyst for oxidative coupling of methane (OCM) to higher hydrocarbons, in the presence of an inactive low-surface area metal filament (Ni-based alloy) a sharp increase in the rate of reaction accompanied by a selectivity shift towards CO and H2 takes place and the oscillatory behavior arises. In the present work the following aspects of these phenomena have been studied: (i) the effect of the state of the surface on the kinetic behavior; (ii) the localization of the oscillatory process. In this work the oxidative transformations of methane were studied with a catalyst system that combines an oxide and a metal component. The presence of both components gave rise to complex oscillation phenomena. The influence of pretreatment and reaction conditions over a wide range of parameters (temperature, total pressure, and oxygen concentration) on the oscillatory process was studied. The possible role of mass transfer and the balance of heat in the reactor were analyzed, and a model for the role of the components in the binary catalyst system is suggested. 1. EXPERIMENTAL The scheme of the experimental setup used for the studies of methane oxidation is given in Fig.l. The oxide catalyst (10% Nd/MgO, specific surface area 15 m2g-1, obtained by impregnation of magnesium hydroxide with neodymium nitrate solution, drying and final calcination at 850~ was placed into a quartz reactor (inner diameter 3 mm). The heating device ensured a constant temperature along the whole length of the catalyst bed (-- 1 cm) and a sharp decrease of temperature outside it, as indicated in Fig. 1. Concentrations of components of the reaction mixture were measured by online gas chromatography (GC). The initial CH4- 02 mixtures containing 5 - 15 vol. % oxygen were prepared in a cylinder and supplied to the reactor. Total pressure

This study was carried out under the financial support of the Russian Foundation for Basic Research (research grant No. 96-03-32440)

758

~

3

5

11

10 2 ///.// /7//,"

1.1

8

T/\

reactor axis

9

>

Figure 1. Experimental setup and temperature profile along the axis o f reactor 1 - quartz reactor; 2 - oxide catalyst bed; 3 - cylinder with gas mixture; 4,5 pressure/flow adjusting valves; 6 - thermocouple (metal filament); 7 temperature recorder; 8 - GC line; 9 - cartier gas inlet line; 10 - GC-sampling valve; 11 - exhaust line

Table 1. Methane and methane/ethane mixtures oxidation over oxide, metal, and combined catalysts (650oC, moxide = 15 mg, W = 9 ml/min, P = 20 kPa) Catalyst

Mixture

0 2 Conversion,

CH4:O2:C2H6

Concentration, vol.%

%

C2H6

CO

C02

H2

Oxide

90: 10:0

37

0.25

1.1

1.2

1.3

Oxide+metal*

90: 10:0

82-88

0.035

3.2-4.3

3.6

6.3-8.3

Metal

90: 10:0

-~0

.

Oxide+metal*

88.5:10:1.5

Metal

88.5:10 : 1.5

* - temperature oscillates near 650oC

89.5-93 -4)

.

.

1.2

3.2-4.1

1.5

-

. 3.5 3.8-5.7 -

-

759 in the reactor ranged from 20 to 100 kPa, and reaction temperature was varied from 500 to 750~ The details of preparation and catalytic performance of the mixed MgNd oxides, as well as the experimental procedure are described elsewhere [5]. A chromel-alumel thermocouple (diameter 0.3 mm, sheathed in a quartz cover or bare) was placed co-axially into the reactor filled with oxide catalyst making it possible to detect temperature oscillations accompanying concentration oscillations. This thermocouple in bare form also acts as the metal component in the oxide-metal binary system. The surface area of the thermocouple filament is ~ 7.5x10 -5 m -2. The effect of the state of the surface on the kinetic behavior was studied using various feeder including the methane-oxygen mixture alternating with inert (He), oxidizing (02), and reducing (HE) gases. 2. RESULTS 2.1 Regularities of Oscillatory Process When the metal filament (thermocouple) is placed inside the quartz cover, methane oxidation proceeds in a steady-state regime with high selectivity to C2 hydrocarbons (Table 1). However, if the cover is removed, i.e. the reaction mixture is in direct contact with the metal surface in the hot zone of the reactor, a sharp increase of conversion accompanied by a dramatic change of selectivity from O C M products to CO and H2 takes place. If the oxide component is removed from the reactor, no conversion of reactants is observed indicating a very low activity of the metal filament in methane oxidation. No reaction occurs also when ethane is added to the methane-oxygen mixture. However, if both oxide and metal components are present in the reactor, some part of ethane undergoes conversion causing additional consumption of oxygen (Table 1). Oxidation of methane in the presence of such a binary oxide-metal catalyst proceeds in an oscillatory regime, and both temperature and concentration oscillations take place. Oscillations arise at the temperature at which the rate of reaction over the oxide component becomes noticeable (-500~ As temperature increases, the oscillation amplitude passes through a maximum. The oscillatory behavior disappears when complete conversion of oxygen is reached. In other words, the range of temperatures in which the oscillations are observed covers the range of oxygen conversions from -- 0 to -~ 100%. Variations of total pressure and of oxygen concentrations in the initial gas mixture change significantly the parameters of oscillation. At reduced pressures in oxygen-rich mixtures complex, but regular temperature oscillations are observed. An example of such a behavior is given in Fig. 2: a simple cyclic temperature oscillation observed at low oxygen levels (5 %) changes to a more complex repetitive pattern if oxygen concentration in the initial mixture is increased up to 15 %. At a total pressure 40 kPa and "bare" temperature Tn~n = 630-640~ the number of peaks in one group increases from l to 7 when oxygen concentration is increased from 5 to 15 %. However, the intervals between maxima in one group remain nearly constant, whereas the magnitudes of the main maxima are very sensitive to oxygen concentration (Fig. 3).

760

T,~

Conc. 02 ,o o ,,

720 -

700 -

680

660

640

-

-

JJL

500

1.000

time, sec.

Figure 2. Variations of the shape of oscillations vs. oxygen concentration

A Tmax, ~

A tl,

sec.

1500

(Ptot =

40 kPa)

t~T, ~

~ t, sec.

50

80 -

450

60

200

40

150

30 - 300

40

100

20 150

10

20

O

I I I I I I I 1 11

0

5

I I I I I I I I I I I I I I I I I I I I I I I I

10

0

15

02 Conc., % Figure 3. Magnitudes of the main maxima and intervals between peaks in one group (At1) vs. oxygen concentration (40 kPa, Tmi~= 630~

0

50

IIIIIIIIIlllllllllllllllllllllllllllllllllllllllllllllllll

0

20

40

60

80

100

P, kPa Figure 4. Oscillation amplitude and period vs. total pressure (10 % 02, Train = 640~

761

T

10 min. in He flow

) L time ~/1 ~ 0 rain. in air flow

| l

n

hydrogen flow

time

T CH4 +02

~H4+02

1

time

time

Figure 5. Effect of reduction / oxidation on the oscillatory behavior (Ptot = 100 kPa, 13 % Train= 550~

02,

their amplitude (AT - Tmax - Tmm) increases at increasing total pressure (Fig. 4). However, the range of variation in the parameters of the oscillations in these experiments was not as wide as that caused by the variations of oxygen concentration at constant total pressure. Since the variations of oxygen pressure have such a strong effect on the oscillatory behavior, some additional information about the mechanism of the observed phenomena can be derived from experiments with preliminary reduction and oxidation treatments of the catalyst(s).

2.2 Effect of Redox Treatment of Catalyst(s) Effects on oscillatory behavior of the treatment of the binary oxide-metal catalyst bed in different gases are presented in Fig.5. If the binary catalyst was treated in inert gas, the sharp increase of temperature begins immediately after the reactants are supplied to the reactor, and then the process proceeds in the regular oscillatory manner, despite a phase of oscillation in which the flow of reactants was switched to the inert gas flow. The oxidative treatment leads to an initial disordered set of oscillations at increased activity. After some period of time the regular oscillations are restored. The reductive treatment resulted in a low-rate steady-state regime. The oscillatory behavior can be restored only by oxidative treatment.

762 3. D I S C U S S I O N 3.1 Localization of Oscillatory Process Synergistic effects and oscillatory behavior are observed only when oxide and metal components are present in the reactor simultaneously, i.e. there is a cooperative effect. Their roles can be better understood, if the space where the oscillatory process is localized is determined. This can be done by an analysis of heat transfer in the reactor. The balance of heat in the reactor may be described as follows:

(1)

W •(AXi AHi) = (dAT/dt) Z(cj mj) + k AT where W AT k q mj AXi AHi

- total flow rate; = (T - Tmm), Tmm - temperature in the reactor in the absence o f reaction or in a low-activity phase of oscillations; heat dissipation coefficient; heat capacities of the substances under heating; masses of the substances under heating; - fraction of reactants converted into different products; - AHf (CH4) - ~AHt (prod.), AHf - enthalpies of formation of methane and of different products of its oxidation via the following reactions: -

-

-

- AHi, kJ/mol a. CH4 + 2 02

--> CO2 -[- 2 H20

801.5

b. CH4 + 02

=> CO2 "[- 2 H2

318.3

c. CH4 + 1.5 0 2

~--> CO 4" 2 H20

518.7

d. CH4 + 0.5 02

=> CO + 2 H2

35.5

e. CH4 + 0.25 02 => 0.5 C2H6 + 0.5 H20

469.0

The value of heat dissipation coefficient k may be estimated from the steady-state data (at dAT/dt = 0), for example from the experiment in which a mixture of reactants is switched to a non-reactive gas flow (Fig.5). Estimates of heat transfer in the reactor based on the data from Table 2 give k -- 6.7x10 -4 J K -1 s-1. Taking into account the additional heat produced in the peak of oscillations, their amplitude and the rate of temperature rise, we estimated the value of Z(q mj) for the substance which undergoes heating as -~ 2xl 0 .3 J K -~. In our experiments presented in Table 2 the masses of oxide and metal components were 0.023 and 0.005 g respectively, giving the (q mj) values equal to 2.15x10 -2 and 2.35x10 -3 J K -l respectively. This estimate indicates that the additional heat evolution in the peak of the oscillatory process is not enough to heat the mass of the oxide component, and the stages of the overall reaction responsible for such a non-steady state behavior and for the formation of final products are likely localized on the metal surface.

763 Table 2. Temperatures and concentrations of products in different phases of oscillation and in the non-reactive mixture (moxide = 23 mg, W = 20 ml/min, P = 100 kPa) Mixture CH4:O2:N2

Phase

T, ~

90: 1 0 : 0 90: 1 0 : 0 90 : 0 : 1 0

minimum maximum no reaction

655 675 585

Concentration, vol.% C2H6 CO CO2 H2 --0 -0 .

3.0 3.4 . .

0.35 0.45 .

6.3 6.5

Effects of the treatment of the binary oxide-metal catalyst in oxidizing and reducing gases on the oscillatory behavior provide further evidence for this conclusion. As is shown elsewhere [5], the treatment of oxide O C M catalyst in oxygen and hydrogen leads to a sharp increase and decrease in their activity, respectively. However, this effect is of short duration: after few minutes the rate of reaction undergoes a relaxation to a steady-state level which is the same in all cases, i.e., if the treatment in hydrogen has an irreversible effect on oscillatory behavior, this can be explained by its influence on the metal component. 3.2 Mechanism of Synergy Since the metal filament is inert in both methane and ethane activation, but active in the binary catalyst, this effect is likely due to reactions involving some intermediates. In the absence of the metal filament, the oxide component is a very efficient catalyst for the O C M process, which is well-known to proceed via the formation and recombination of free methyl radicals [6]:

[O1

+ CH4

=> [OH] + CH3

CH3

+ CH3 (+ M) => C2H6 (+ M)

(2) (3)

where [O] and [OH] represent active sites on the oxide catalyst surface in oxidized and reduced states, and M represents a third body. Kinetic simulations based on the model described elsewhere [7] show that, if the probability of repeated collisions with the oxide catalyst surface is high, up to 90% of CH3 radicals formed by reaction (2) can undergo the reverse transformation into methane [OH] + CH3

=>

[O]

+ CH4

(4)

The apparent conversion rate measured as the rate of the formation of the final products is much lower than the rate of reaction (2) due to competition by reaction (4). A complete analysis of mass transfer should include consideration of diffusion (in the pores of the oxide catalyst and in the space between the grains) and of the accompanying reactions of radicals with different species in the gas phase and with surface active sites. This work is presently in progress, and here only a brief discussion is providedThe diffusion coefficient for methyl radicals at the conditions of our

764 experiments (500-600~ and 20-100 kPa) is 0.5-5 cm 2 s-l. If the maximum diffusion path is equal to the radius of reactor, the diffusion time (5x10 -2 - 5x10 -3 S.) is substantially less than the residence time in the reactor (-0.1 s.). This estimate indicates that CH3 radicals escaping from the grains of the oxide catalyst can reach the metal surface and undergo secondary transformations, i.e. additional processes of CH3 removal can occur that decrease the fraction of radicals transformed back into methane and increase the apparent rate of conversion. Since O C M products do not form over the binary catalyst, this suggests that CH3 radical capture by the metal surface is highly probable. As is mentioned above, the parameters of the oscillatory process are very sensitive both to the variations in oxygen pressure and to the pretreatment of the catalyst. In general, the increasing degree of oxidation causes a larger oscillation amplitude. On the contrary, after the preliminary treatment in reducing gas (hydrogen) the process proceeds in a steady-state regime and no major cyclic changes in activity are observed. Since the oscillatory process is localized on the metal component, it is likely that the effects of reduction and oxidation treatments are caused by the differences in the degree of oxidation of the metal surface. According to existing notions [1], rate oscillations in oxidative catalytic reactions may be caused by the existence of some "buffers", or "reservoirs" of oxygen which are able to supply it to the zone of reaction. In particular, a kinetic model based on such ideas is in satisfactory agreement with the experimental data on self-oscillations observed during CO oxidation on a Pt single crystal [8]. In the case of our experiments, the strong effect of oxygen concentration on oscillatory behavior and the sharp increase of activity after the treatment in oxygen is likely due to the variations of surface concentration of reactive oxygen species participating in transformations of methyl radicals:

I01

+ CHa

=> CO

+ 3/2H2

(5)

w h e r e / O / - oxygen species able to oxidize adsorbed CH3 radicals on the metal surface. The data presented above indicate that, although the metal component is not able to activate saturated hydrocarbon molecules, it is very active as a trap for free radicals generated by the oxide catalyst. As a result, reaction (5) competing with (4) leads to the apparent increase in activity of the binary catalyst, and the selectivity of the overall process is determined by the competition between reactions (3) and (5). On the other hand, the treatment in hydrogen flow causes exhaustion of this oxygen "buffer", which cannot be restored in the presence of both reactants in the reaction mixture due to the high reducing activity of methyl radicals. The positive effect of the treatment in helium is likely due to the removal of some species which are present in the reaction mixture and compete with CH3 radicals for the surface active sites. Their removal from the surface in the inert gas flow leads to the immediate restoration of activity when the reaction mixture is switched back. The combination of reactions (2) and (5) may be considered as a scheme for direct methane oxidation to synthesis gas (CO + H2). Similar reactions may determine the high efficiency of mixed catalysts containing Ni and rare-earth oxides for the partial oxidation of methane to synthesis gas [9]. This mechanism does not require a preliminary total oxidation of methane followed by its reforming with CO2 and/or water which was considered as the main route for synthesis gas formation [10,11]

765 CH4

+ 2 02

=> CO2

CH4

+ CO2

=> 2 CO + 2 H2

CH4

+ H20

=> CO2 + 3 H2

+

2 H20

(6) (7) (8)

and, consequently, has no thermodynamical limitations resulting from the positive values of AG in reactions (7) and (8). CONCLUSIONS 1. The cooperative effects observed during methane oxidation over a binary oxidemetal system are due to the formation of active intermediates (free methyl radicals) over the oxide component, their escape from the grains of oxide, and transformation into the final products (including CO and HE) over the metal component, which proceeds in a non-steady-state oscillatory regime. 2. The strong effects of the variations of oxygen concentration and of the preliminary reduction and oxidation treatments of the catalyst indicate that active oxygen species formed on the metal component participate in the transformations of methyl radicals into the final products. REFERENCES 1. V. I. Bykov, Modeling of Critical Phenomena in Chemical Kinetics, Nauka, Moscow, 1988. 2. M.M.Slin'ko and N. I. Eaeger, Oscillating Heterogeneous Catalytic Systems. Studies in Surface Science and Catalysis, v.86, Elsevier, 1994. 3. Yu. P. Tulenin, M. Yu. Sinev, and V. N. Korchak, I l th Int. Congress on Catalysis, June 30 - July 5, 1996, Baltimore, ML, USA, Programme and Book of Abstracts, P-275. 4. D.G. Filkova, L.A. Petrov, M.Yu. Sinev, and Yu.P. Tyulenin, Catal. Lett., 13 (1992) 323. 5. M. Yu. Sinev, G. A. Vorob'eva, and V. N. Korchak, Russ. Kinetics and Catalysis, 27 (1986) I 164. 6. D.J. DriscoU and J. H. Lunsford, J. Phys. Chem., 89 (1985) 4415. 7. M. Yu. Sinev, Catal. Today, 24 (1995) 389. 8. A. L. Vishnevskii, V. I. Elokhin, and M. L. Kutsovskaya, React. Kinet. Catal.Lett., 51 (1993) 211. 9. V.R. Choudhary, V. N. Rane, A. M. Rajput, Catal. Today, 22 (1993) 289. 10. D. Dissanayake, M. P. Rosynek, K. C. C. Kharash, and J. H. Lunsford, J. Catalysis, 132 (199 l) I 17. I I. O. V. Krylov, Russ. Chem. Rev., 61 (1992) 2040.

This Page Intentionally Left Blank

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

767

C a t a l y t i c Combustion of Methane: Activation and Characterization of Pd/AI203 Maria GraCa Carneiro da Rochaa and Roger Fretyb "Chemical Engineering Department, Polytechnic School, Federal University of Bahia, Brazil bLACE, Laboratoire des Applications de la Chimie/t l'Environnement, CNRS, France

Pd/ml203 catalysts for methane oxidation were observed to show an increase in activity after treatment under a reaction mixture. In order to better understand this phenomenon, 2% Pd supported on 6 and a-alumina were pretreated under 1-12or the reaction mixture (CI-h, 02, N2, CO2,1-120) at 500, 700 and 900~ then tested at 320~ (OJCI-I4=2/1) and I atm. The catalysts were characterized by temperature programmed reduction, X-ray diffraction, transmission electron microscopy and X-ray photoelectron spectroscopy. It was observed that the activity was related to the presence of both Pd/PdO species in the bulk and at the surface of the catalyst. In the active state the PdO in the biggest particles appeared more reactive towards hydrogen than in smaller particles, suggesting a better mobility of the existing oxygen species. The relative proportion between PdO and Pd and the increase of the extent of their interface is probably the origin of the activity of the catalysts. 1. Introduction Methane oxidation at mild or low temperatures can be catalyzed by platinum group metals. Palladium is one of the most efficient metals (1) and has been studied over many supports (26). This particular metal, when supported on alumina, begins to show an increase in its activity between 350 and 420~ At these conditions a general increase in the active species particle size is observed. Primet and Briot (7,8) defined two states for the Pd/AI203 supported catalyst: a state I, obtained after simple reduction and a state II after the catalyst had reacted at 600~ for 14 h under O2/CI-h=4/1. State II was more active than state I and showed a lower binding energy of oxygen with palladium. However, the state of the active phase was not clear. The differences in activity, also observed by others, have also been related to the formation/decomposition of PdO (9), to the oxygen adsorbed on metallic Pd (2), to the modification of Pd surface species (3), and to the reconstruction ofPdO crystallites (4, 10). One of the hypotheses for the activation of the Pd catalysts was the establishment of an epitaxy between the metal and the support (8, 11). In the present work, palladium catalysts supported on a well-crystallized alumina, the aalumina, and a transition alumina were studied. The state of the catalysts was examined after a reproducible standard test procedure and the catalysts samples were characterized by different physicochemical techniques.

2. Experimental Pd catalysts supported on aluminas (Puralox transition alumina from Condea and SPH 512 a-alumina from Rhone~Poulenc) were prepared by wet impregnation with PdCI2 to give approximately 2 wt.% Pd, and were then dried overnight in air at 120~ treated in N2 at 550 oC, reduced in H2 at 600 ~C, cooled in N2 to room temperature and stored in a dry box. The catalysts supported on ~-alumina and on a-alumina will be referred to DA and AA respectively. DA-SR means a simply reduced catalyst, DA-500 a catalyst treated under reaction mixture at 500~ C and so on.

768 The powdered catalysts were evaluated in an integral continuous fixed bed reactor (1 g cat, total gas flow rate 6 l/h), with on line gas analysis by gas chromatograhy using both thermal conductivity detector and flame ionization detector. A standardized procedure was adopted to get reproducible samples for characterization. The stored catalysts were pretreated "in situ" in H2 at 400 ~C and heated under N2 flow to 500, 700 or 900 ~C. At those temperatures the catalysts were treated under reaction mixture (molar ratio O2/CI-h=2/1, N2 balance ) for 3 h and cooled in N2 to 320~ to be isothermally tested for 4 hours consecutively at 320, 340, 350 and 360~ (standard reaction). After reaction, the system was cooled to room temperature in N2 and the catalysts were stored in a dry box after air exposure. The catalysts were characterized in the fresh state and at the end of the standard reaction. Temperature programmed reduction was performed in a dynamic system with TCD and quadrupole mass spectrometer analysis (Balzers QM1420) with a 1% H2-Ar, total flow rate of 1.1 l/h and a heating rate of 10 ~C/min. TPR analysis was performed in three steps between -78 and 500~ in order to determine differences in reactivity of the oxygen present in the samples. This analysis allowed also the quantification of the total amount of oxygen present in the catalyst after the standard reaction and to determine the behavior of palladium hydrides. At the end of the TPR at 500 ~C, the catalyst was cooled under Ar to room temperature and oxygen was introduced. This oxygen was then titrated with hydrogen and the Pd dispersion was determined from the amount of irreversible consumption of hydrogen. Powder X-ray diffTaction patterns were recorded using Cu Ktx (1.5418 A) radiation on a Philips 1050/81 vertical goniometer, fired with a diffracted beam graphite monochromator. X-ray photoelectron experiments were carried out in a Escalab 200R (Fisons Instruments) using monochromatic AI Kot radiation and operating at 10.9 Tort base pressure. Transmission electron microscopy analysis was performed in a JEOL 100CX either by direct observation or after extractive replication of samples (support dissolution). 3. R e s u l t s

The Pd coment determined by atomic absorption was 1.77 % wt for the DA catalyst and 2.13 % wt for the AA. Table 1 Initial activities and Pd dispersion Catalyst

Dispersion(%)

Initial activities 320~ (molCHa/h.molPd)

A c t ~ tr~t/Actn~

DA-SR

26.8

0.9

DA-500

20.7

2.6

3

DA-700

5.0

19.6

22

DA-900

2.4

37.3

42

AA-SR

16.6

1.1

-

AA-500

7.9

1.3

1

AA-700

3.5

15.3

13

AA-900

-

8.7

8

769 Table 1 summarizes the results of activity tests at 320~ and the values of dispersion, after reaction. The dispersion of the fresh catalysts was 29% for DA and 23% for AA. Despite a decrease in the dispersion of the active phase after the high temperature pretreatment in the reaction mixture, there was an increase in activity with the DA catalyst showing a higher degree of activation (DA-900/DA-SR=42). Examination of the values of initial and final activity of the AA catalyst also indicates unstable behavior for the AA-SR and AA-500 that could be called an extensive induction period. These results indicate that the support has a strong influence on the catalyst stability (Figs. 1 and 2).

~oo

Conversion(%)

.

Conversion(%) 100 - - - -

~

....

8o~

*'-~

i

80

~ ....

80 . 99

"-~."

60

i:::S

.~.~

..,.-

40i

%

40

~-~ ..:

........ .~

20

..... ., ............

.."

6

8

0

0

2

4

~176

,~

20

,4~.~. e * ' t t'

I

i

I

10

12

14

0:~

16

0

'

2

Time (h) Figure

1

~

Conversion x time for DA

4

6

8

i

i

I

10

12

14

16

Time (h)

series:(-)DA-SR, (+)DA-500, (~)DA-700,

Figure 2: Conversion x time for AA series:(-)AA-SR, (+)AA-500, (~-)AA-700,

(~)DA-900.

(D)AA-900

Figure 3 shows a typical TPR analysis profile between -78 and 500~ Three different temperature domains were revealed in the hydrogen consumption profile when contacting the catalysts with the reducing mixture: a first peak ~ - 7 8 ) at -78~ a second peak when the temperature rose from -78~ to room temperature (HT25) and the fourth one, only present for the AA catalysts, at temperatures higher than 120 o C (I-IT120). In the case of HT-78 two peaks were obtained, their difference giving the amount of irreversibly consumed hydrogen at these conditions. The hydrogen consumption features HT-78 and HT25 are both assigned to reaction with oxygen and to Pd hydride formation. The third peak observed between 75 and 120~ during the heating period, corresponds to an increase in the gas phase hydrogen concentration, associated with the decomposition of the previously formed hydrides. Finally, a fourth peak is observed, only for the AA catalysts, indicating a new hydrogen consumption feature at temperatures close to 120 ~C, that is likely due to the reduction of particular PdOx species. At higher temperature, a small hydrogen consumption is observed and is ascribed to the reduction to methane of the COx species adsorbed on the catalyst surface. The oxygen content of the catalysts in their fresh state and at the end of the standard reaction were estimated, after correction for the hydrogen consumption due to Pd hydride formation. It was thus possible to calculate the average oxidation degree of the catalysts and the number of

770 oxygen layers refered to the exposed metallic surface area determined by hydrogen titration. These results are summarized in Table 2.

Hydrogen Pressure (arbitrary units) HT(-78) ...........................

500 ~

................ . T ( t a m b i

0

.......................................................................

200

400

...............

600

Time (arbitrary unit) Figure 3" Typical TPR profile

Table 2: Hydrogen balance and Catalyst HT-78 HT25 . . . . Fresh DA 0.04 1.1 DA-SR 0.01 2.5 DA-500 0.06 2.4 DA-700 0.2 1.2 DA-900 0.3 0.9 Fresh AA 0.02 0.7 AA-SR 0.03 0.4 AA-500 0 0.4 AA-700 0.05 0.9 AA-900 0.09 0.5

oxygen estimation during TPR (H atoms/Pd atoms) Hyd X HT120 HT (1) O/Pd PdO(%) O . . . . . layers 0.15 1.0 0.9 0.5 50 1.1 0.2 2.3 0.8 1.1 100 3 0.4 2.1 0.6 1.0 100 3 0.4 1.0 0.2 0.5 54 6 0.4 0.7 0.06 0.4 36 11 0.1 0.4 0.2 0.7 0.2 22 0.6 0.01 1.1 0.7 0.5 0.5 54 2 0.03 1.2 0.8 0.2 0.6 60 6 0.2 0.8 0.1 0.1 0.5 46 9 0.3 0.9 0.2 0.01 1 -

(1) Hydrogentitration For the AA and DA catalysts the "HT-78" peak increases with the severity of the treatment in the reaction mixture. The "HT25" decreases for DA catalysts but increases for AA catalysts, indicating that the reactivity of the oxygen of the active phase is dependent on the support. The amount of hydride formed is lower for the AA catalysts because of the existence of a quantity of PdO which is not reducible in the temperature range where hydrides can be formed. These results indicate that the low initial activity of DA-SR and AA- SR and AA-500 catalysts does not have the same origin. The fresh catalysts possess a total amount of oxygen which is close to that of an oxygen monolayer on the Pd surface. This monolayer is likely formed upon contact with air. The CH4 oxidation reaction changes this figure. Difference in the ratio O/Pd is observed. If it is assumed that PdO is the stable oxide in the catalysts, all the available Pd is oxidized in the catalysts DA-SR and 500. It appears that a fully oxidized Pd is not an active state for the

771

catalyst (Table 1). On the other hand, in the highly active catalysts, as are those after treatment at 700 and 900 ~C, the Pd phase is partially in a metallic state and partly an oxidic state. The AA-SR and A-500 catalysts showed a high activation which can be related to the presence of a metal-oxide mixture. In all cases, the metallic character increases after treatment in the reaction mixture at 900 oC. The preceding results indicate that the small particles (D=15-25%) generated by the low temperature treatment do not display a good initial catalytic activity. On the contrary, the large particles obtained alter the high temperature pretreatment (D=2-5%), consist of a mixture of a reduced and an oxidized phase, and exhibit higher specific activities. Furthermore, the amount of hydrogen consumed for hydride formation relative to all the hydrogen consumed increases with the pretreatment temperature. In agreement with the literature, hydrides are more easily formed in big particles (12). XRD diagrams (Fig. 4) show the evolution of the structure in the DA series catalysts. The only peaks that appears in the DA-SR catalyst and DA-500 are due to PdO. The metallic Pd content increases for the catalysts pretreated at 700 and 900 ~C under the reactant mixture. Calculation of the crystallites sizes indicates that the PdO crystallites are smaller than the metallic ones, suggesting a breakup of the latter upon oxidation (Table 3). The more active catalysts are composed of both oxidized and reduced phases, in agreemem with the results of the TPR experiments. The relative amounts of each phase was calculated alter standardization (Table 3). The results are consistent with the complete accounting of Pd in the more active catalysts. In the lower activity catalysts, there is a fraction of palladium that is not visible by XRD, probably because of the small size of the particles, below the XRD detection limit. Pd [11t1 PdO[lOI]

oo-

oC

'

' I 30

"'

i

1 35

'

1~ 40

"

1 2/)"

Figure 4: XRD diagram for the DA series catalyst: a) pure A1203, b) DA-SR, c) DA-500, d) DA-700, e) DA-900. Table 3 Particle size calculated by XRD by the _S.cherrer equation and from H T Catalyst ....DRX (/~) . Phase (_%) . . . . . VdO(10!),, P d ( l l l ) ,PdO Vd Total SR 42 70 70 500 56 70 70 700 78 130 56 36 92 900 89 180 44 56 100

HT(/~) 40 60 220 560

772 As shown in Table 3, the titration values also give an idea of the size of the metallic particles. Unfortunately, the technique does not allow to distinguish between what was PdO and Pd before the reduction. Another AA catalyst with 2.7 wt % Pd content showed the presence of both phases by XRD, after simple reduction and after treatment at 900 ~C, the later with a higher metallic content. In general, the spectra do not show significant crystal modifications, neither for the active phase, nor for the support. TEM micrographs (Figs.5-7) show the sintering and evolution of the Pd particle morphology in going from the reduced sample to the catalysts pretreated under severe conditions of reaction. The results agree with those of Chert and Ruckenstein for oxidized Pd (13). The DA-SR catalyst are composed of regular, spherical small particles (20-50A). The DA-700 catalyst has a heterogeneous texture with small spherical particles (20-80 A) decorating larger irregular particles (>280 A). On DA-900 the small particles disappear and a particle size between 210 A and 700 A is observed. Both, large particles decorated by small ones and smooth particles, are observed. Many particles seem to be polycrystalline. The AA series shows the same particle evolution with a slightly different size distribution. The fresh catalyst has particles of size between 20 and 300 ~ most of them around 20 K The AA-SR has basically the same size distribution but large particles begin to appear (250-500 A). In the AA-700 catalyst, particles with size lower than 40/~ disappear and sintering is observed with particles up to 900 A. Finally in the AA-900 only big particles are observed. Micrographs show the size and morphology evolution of the DA and AA series of catalysts (Fig. 5 and 6 ). The morphology changes in these catalysts with the high temperature pretreatment is evident: initial spherical particles change to irregular ones decorated by small ones. Perforated particles are also observed (Fig. 7).

Figure 5: TEM micrographs for DA series catalysts: (a)DA-700, (b)DA-900.

773

Figure 6: TEM micrographs for AA series catalysts: (a)AA-fresh, (b)AA-SR

80 nm Figure 7 TEM micrographs details of DA and AA catalysts: (a) DA-700, (b) AA-SR.

....

!ili!ii:i(ii!il

774 XPS results are presented in Table 4 . The decrease in Pd/AI ratios is consistent with the catalyst sintering. Higher values for the AA catalysts are related to the low specific area of the support compared to that of the DA catalysts. The CI/AI ratio is reported because the literature suggests that dechlorination could be related to catalyst activation during methane oxidation (3, 10, 14). This could be true for the DA catalyst, but not for the AA catalyst which shows a lower activity for a lower CI/Pd ratio. Table 4: Binding enerBies for Pd 3d 5/z and atomic ratio by XPS Catalyst BE (eV) Atomic ratio Pd/Al Cl/Al DA-SR 336.0 337.9 0.007 0.009 DA-900 334.6 336.5 0.002 0.002 AA-SR 335.7 336.9 0.044 0.01 AA-900 334.3 336.1 0.011 0.004 After treatment at 900~ both catalysts series present similar BE values (334.6 and 334.3 eV) which are assigned to metallic Pd, even though these values are lower than that reported in the literature, 335.2 eV (15). Such negative deviations have been observed in the literature and related to a partial reduction of the support (15). The 336.5 and 336.1 eV values may be assigned to PdO, whose reported B.E. is 336.5 eV (16). These results suggest that the active phase contains both metallic Pd and PdO. The DA-SR presents BE values of 336.0 and 337.9 eV close to those reported by Otto et al. (16) assigned to two species of PdO: the first one, normal, and the second one, with BE of 338.3 eV possibly belonging to small clusters ofPd § in strong interaction with the support or PdO2. The AA-SR presents a mixture of oxidized and reduced palladium, different from the species obtained after the reactive high temperature pretreatment. 4. Discussion

An increase in the activity of alumina supported Pd catalysts under conditions of methane complete oxidation was once again observed in the present work. Even catalysts which were treated under the reaction mixture at 900 ~C, an elevated temperature even for a combustion reaction, were activated. The samples can only be activated under reaction conditions and at high temperatures. Furthermore, the increase in activity was accompanied by an increase in the particle size. However, the catalytic activity could not be correlated directly with the particle size as already observed by others (4, 5). The absence of a correlation is not surprising since it was observed that the active catalyst contained both metallic and oxidized palladium. By quantification of the oxygen content it was possible to estimate the 'thickness' of the oxide layer, related to the exposed surface of Pd (whose size was estimated after complete reduction of the sample). It was observed that the more active catalysts had a larger number of oxygen "layers" which suggested the participation of bulk Pd in the reaction. Therefore, for methane oxidation, the Pd/AI203 catalytic properties seem to be controlled by a redox mechanism, where the oxidation and reduction rates of the active phase depend on the size and stability of the Pd and PdO species. The small PdO particles seemed to be poorly reactive, due to their higher stability or to a greater interaction with the support. In fact, analysis of TPO and TPD data revealed that the small particles of PdO decompose at

775 temperatures 40 to 50 ~C higher than those present on the particles after reactive treatment at 700 and 900 ~C. These results, associated with the heat of adsorption of oxygen on different sizeA Pd particles (17) or to hydrogen reactivity on PdO (7,18) show that oxygen mobility is higher in the larger particles than in the smaller ones. In summary, for the complete oxidation of methane, Pd/A1203 catalysts with small particles do not constitute the active phase. A metallic phase alone is not effective. On the metallic phase, methane undergoes dissociative adsorption, probably favored by a high electronic density as observed by XPS. The high electronic density of Pd could be due to a partial reduction of the support, to the presence of carbon donor species around the particles, or to adsorbed oxygen species on PdO localized very close to the metallic state. QMS analysis during TPD of the used catalysts indicated the evolution of CO2 from the sample at the temperature of PdO decomposition. The presence of carbon filaments after exposure to CI-h at high temperature was also observed by SEM. It is known that PdO decomposes around 850~ so the treatment at 900~ under the reaction mixture might be expected to form only metallic Pd. However the DA-900 catalyst, was more active than DA-SR and DA-500 and at the end of reaction a mixture of Pd-PdO was observed. This means that the reaction itself allowed a partial oxidation of the catalyst. Hence, it can be concluded that catalysts which are fully reduced or fully oxidized are not active (DA-S1L AA-SR and AA-900). The determination of the active sites remains a challenging subject because we are not able to quantify directly the active fraction of the catalyst constituted by PdO. The overall metallic surface area is not directly related to the active sites working during the reaction. The TEM observations in this study are similar to those obtained by Chen and Ruckenstein (13), although in a different range of temperature. In this case the atmosphere and its interaction with palladium was more complex. An initial strong sintering was observed, and, subsequently in some particular sites there was an apparent oxidation which was responsible for the formation of cavities. In a following step, the particles started to fragment and their shape became irregular (cauliflower shaped particles). Further differential oxidation was observed which seemed to depend on certain active sites present on the catalyst. At 700 ~C, a fraction of the Pd could reduce and apparently this led to the agglomeration of particles, while at the same time a spreading of particles occurred due to their oxidation. This is probably the reason for the decorated cauliflower shaped particles; these particles are likely metallic inside with small PdO particles outside. At 900~ PdO is not stable and the particles are more crystalline, even if some decorated particles still exist. The disappearance of those small particles was also observed after reduction of the used catalyst (17) and was the possible cause for the decrease of activity, even if some particular sites were still oxidized. EDX analysis of some particles indicated the presence of both Pd and PdO in the very same particle in different proportions depending on the position. In this particular system, well crystallized particles are not active, but particles rich in defects, capable of inducing modifications in the Pd state are. In this situation, the catalyst activity is probably related to the existence of a Pd-PdO interface, not to an active phasesupport interface From the differences observed by TPR analysis, the species present in the fresh catalyst underwent changes after reaction. All results indicated that the active catalyst contained a PdPdO mixture, coexisting mainly in the same particles, which was formed during the reaction. These phases were of a size and morphology that allowed oxygen mobility. The small

776 particles alone, initially in complete oxidized or reduced state, were not active, due probably to a strong interaction with the cartier. Depending on the support, the activation can proceed in a different way. It is probable that during the reaction there is oxidation and reduction of the active phase, with the catalytic reaction occurring at the oxide-metal interface which is continuously undergoing reconstruction. Acknowledgements: M. G. C. da Rocha acknowledges the financial support of CNPq, Conselho Nacional de Desenvolvimento Cientilico e Tecnol6gico of Brazil and the contributions of M. Bran, P. D61ich6re in the XPS experiments, G. Bergeret in XRD and C. Leclercq, F. Beauchesne in the electron microscopy analysis. References 1. R. B. Anderson, K. Stein, J. J. Feenan and L. J. E. Hofer, Ind. Eng.Chem., 53 (1961) 809. 2. C. F. Curtis and B. M. Willatt, J. Catal., 83 (1983) 267. 3. R~ F. Hicks, H. Qi, M. L. Young. and R. G. Lee, J. Catal., 122 (1990) 295. 4. T. R. Baldwin and R. Butch, Appl. Catal. 66 (1990) 359. 5. F. H. Ribeiro, M. Chow and R. A. Dalla Betta, J. Catal., 146 (1994) 537. 6. J. G. McCarty, Catal. Lett., 26 (1995) 283. 7. P. Bdot and M. Primet, Appl. Catal., 68 (1991) 301. 8. P. Bdot, Thesis, Universit6 Claude Bernard, Lyon I, France, 1991. 9. R. J. Farrauto, M. C. Hobson, T. Kennelly and E. M. Waterman, Appl. Catal. A: Gen., 81 (1992) 227. 10. T. R. Baldwin and R. Butch, Appl. Catal., 66 (1990) 33 7. 11. P. Briot, P. G~ezot, C. Lederc and M. Pdmet, Microsc. Microanal. Micr0strut., 1 (1990) 149. 12. M. Boudart and H. S. Hwang, J. Catal., 39 (1975) 44. 13. J. J. Chert and E. Ruckenstein, J. Phys. Chem., 85 (198 l) 1606. 14. D. O. Simone, T. Kennelly, N. L. Bmngard and R. J. Farrauto, Appl. Catal., 70 (1991) 87. 15. T. H. Fleisch, R. F. Hicks and A .T. Bell, J. Catal., 87 (1984) 398. 16. K. Otto, L. P. Haagk and J. E. de Vales, Appl. Catal. B: Environ., 1 (1992) 1. 17. P. Chou and M. A. Vannice, J. Catal., 105 (1987) 342. 18. M. G. Almeida, F. Beauchesne, M. Pfimet and R. Frety, Book of Abstracts, vol 2, Europacat, p 917, Montpellier, France, 1993.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

777

Activity o f manganese dioxides towards V O C total oxidation in relation with their crystallographic characteristics C. LAHOUSSE, A. BERNIER, E. GAIGNEAUX, P. RUIZ, P. GRANGE, B. DELMON Unit6 de Catalyse et Chimie des Mat6riaux Divis6s, Universit6 catholique de Louvain, 2/17 P1. Croix du Sud, B 1348 Louvain-la-Neuve, Belgium This paper describes the catalytic activity of various forms of manganese dioxides towards volatile organic compounds deep oxidation. Important differences in activity are evidenced for very closely related structures ; the most active sample is a high surface area nsutite. A parallel is drawn between the findings in the literature of battery applications and the catalytic activity results. The superior activity of nsutite is attributed to a different oxygen coordination and to the clustering of cationic vacancies in the bulk which improves electronic and protonic conductivity.

1. INTRODUCTION Whatever the application, mild or deep oxidation, a limited number of parameters are usually thought to influence the catalytic properties of oxides. These include the nature of the metal cation(s), their oxidation state, their first coordination sphere and the corresponding density of each type of site at the surface. In this work, the catalytic activity towards Volatile Organic Compound (VOC) deep oxidation of 3 crystallographic forms of manganese dioxides: pyrolusite (13), nsutite (~,)and ramsdellite was studied. These 3 forms, respectively named in this article: py-MnO2, nsu-MnO2 and rams-MnO2, are very closely related in structure [ 1]. In each structure the Mn IV cation is in an octahedral environment with each oxygen comer of the MnO6 octahedron shared by 3 octahedra, the only difference being in the arrangement of the O coordination which can be either planar or tetragonal truncated. So, all of the above characteristics namely cation nature, oxidation state of manganese and first coordination sphere are identical in these 3 phases. Moreover, a SEM study did not detect any anisotropy and preferential exposure of specific faces which could give rise to a different catalytic behavior. The aim of this paper is first to report on the catalytic activity of these very closely related structures of MnO2, then to show that characteristics like oxygen coordination, which are not usually considered for explaining the catalytic behaviour of oxides in catalysis, do affect their catalytic properties. One sample of py-MnO2 and rams-MnO2 and two samples of nsu-MnO2 of different specific surface areas were used. 2. EXPERIMENTAL

2.1. Catalysts origin The catalysts were provided by SEDEMA-SADACEM, Belgium. Table 1 gathers the

---

Table 1 : Comparison of the characteristics of the different forms of MnOz

Catalysts studied y-Mnoz (1 00) y-Mnoz ( 100) y-Mnoz ( 100) y-Mnoz ( 100) Y-MnOz (35) Y-MnOz (35) Y-MnOz (35) Y-MnOz (35) (3-Mn02 P-MnOz P-MnOz P-MnOz

MnOzrams MnOzrams MnOzrams MnOzrams

data p Crystalline form nsutite nsutite nsutite nsutite nsutite nsutite nsutite nsutite pyrolusite pyrolusite pyrolusite pyrolusite ramsdellite ramsdellite ramsdellite ramsdellite

EDEk )article

surf.

mVg ratio ize/l.tn 1.95 35 100 1.95 1.95 1.95 1.95 1.95 1.95 1.95 1.96 1.96 1.96 1.96 1.95 1.95 1.95 1.95

35 35 35 37 37 37 37 34 34 34 34 20 20 20 20

100 100 100 35 35 35 35 50 50 50 50 18 18 18 18

voc

7

Ea Hex Ea Hex Ea Hex Ea Hex Ea Hex Ea Hex Ea Hex Ea Hex

----

Conversion % (W=O. Ig ; [VOC]=2SO feed composition 100 120 140 150 160 21 60 88 100 25 Ea+Hex 0 0 0 0 51 Ea+Hex 16 39 48 77 Ea+Hex+HzO 10 4 4 5 5 Ea+Hex+HzO 7 84 7 6 24 48 Ea+Hex 1 5 0 0 0 Ea+Hex 10 32 47 20 Ea+Hex+H20 12 7 9 10 9 7 Ea+Hex+H,O 41 7 10 21 2 Ea+Hex 0 0 0 1 0 Ea+Hex 20 23 32 12 Ea+Hex+HzO 10 7 3 7 7 Ea+Hex+HzO 8 7 40 13 20 2 Ea+Hex 1 1 0 0 0 Ea+Hex 40 9 16 24 Ea+Hex+H20 10 7 6 5 5 Ea+Hex+HzO 6

10 d m i n air)

TI --- - - -180 200 220 240 --- -- 260 - 280 O(

93

100

100

50

80

98

100

87

98

100

57

84

98

7

18

38

90

100

100

70 99 20 95 0 63 6 90 10 61 4

99 60 100 22 98 8

61

100

40 95 16

--- ---- --43

64

81

90

779 characteristics of these products. Thanks to careful preparation by SEDEMA, each of these samples also presents a reasonable surface area together with a "good" crystallinity. According to the data provided by SEDEMA, the amount of CO2, S, alkaline, alkaline-earth and transition metal cation content is extremely small on the nsu-MnO2 samples ; 0,3 % CO2 and 0,02% S can be found in the rams-MnO2 sample. In addition, the XPS analysis shows a trace amount of S (SO4) on the surface of the py-MnO2 sample.

2.2. Catalyst characterization and testing XPS enables the determination of the oxidation state of the Mn ions. The most precise evaluation rests on the measurement of the binding energy difference between the Mn 3s main peak and its shake-up satellite [2-4].The XPS analysis of the samples was performed on a SSI X-probe (SSX 100/206) spectrometer of FISONS equipped with a monochromatised microfocused AI I ~ X-ray source (1486.6 eV). The angle between the sample surface and the electron detection axis was 55 ~ The analyser pass energy was set at 50 eV and the analyzed area was 1.4 mm 2. At these conditions, the energy resolution determined by the Au 4f7/2 full width at half maximum (FWHM) of gold was 1.1 eV. The XPS results are reported in Table 1. The BET surface area of the samples was measured by nitrogen physisorption at 77 K (196~ using a Micromeretics ASAP 2000 apparatus (Table 1). The samples were examined using a Hitachi S-70 scanning electron microscope (accelerating voltage : 15 kV). XRD patterns of the samples were obtained on a Siemens D5000 apparatus working with the copper K~ line. Ethylacetate (Ea) and n-hexane (Hex) were chosen as representative VOC's reactants [5]. 250 ppm of the two molecules in air were used with a contact time of 60 kg s/m3 (which corresponds to a space velocity of 72000 h-1 (NTP) with nsu-MnO2). For experiments with water vapor, the stream was saturated at 25~ with 20000 ppm of water. The contact time, the concentration, the nature of the VOCS and the presence of water vapour are representative of the conditions of VOC removal in printing industries [5]. Before measuring the conversion as a function of temperature, the catalysts were stabilized by letting the reaction proceed for 16 h at 150~ The conversion was measured as a function of time during this period, at which end it reached a constant value. Then, the temperature was decreased to 100~ and increased by steps of 20 or 10~ until complete combustion was obtained. The conversion was then again measured at 150~ to verify that the conversion at the end of the test was the same as after the preliminary stabilisation period. Table 1 presents the conversions for the same catalyst weight, and thus allows a comparison of the specific conversions (conversion per gram of catalyst). The intrinsic conversions (conversion per square meter of surface area) will be compared in the figures presented in this paper. 3. RESULTS The conversion of the two reactants on the different catalysts in the presence or in the absence of extra water vapour is given in Table 1. The existence of a competition for adsorption between the different VOCs has been demonstrated in an earlier communication [6,7]. As far as Ea and Hex are concerned, the presence of Hex in the flow has no influence on Ea conversion, but Ea, strongly inhibits Hex conversion as long as some of it remains present. Nevertheless, for a same catalyst weight, the conversions measured with nsu-MnO2 are about five time higher than those measured with the other forms of manganese dioxide (see Table 1). Figure 1 provides a comparison between the different catalysts corrected from the variation of surface area. The differences in activity appear when both specific conversion

780

and conversion corrected for surface area are considered. According to the Ea and Hex conversion, the MnO2's efficiency for the VOC removal can be ranked as follows: py-MnO2< rams-MnO2< nsu-MnO2 of 35 m2/g< nsu-MnO2 of 100 m2/g. A mechanical mixture of pyMnO2 and rams-MnO2 shows an activity which corresponds well to the average value between those of rams-MnO2 and py-MnO2. Figure 1: Comparison of the intrinsic activity (towards Ea)of different form of MnO2 30 25

, 2o ~ 5 ~10 5 0 py-MnO 2 + rams-MnO 2 35 m2/g

rams-MnO 2 18 m2/g

nsu-MnO 2 35 mVg

nsu-MnO 2 100 m2/g

The effect of water will be described more thoroughly in a future paper [8]. As shown by Table 1 data, water affects the activity of manganese dioxides. It also dramatically reduces the time needed to obtain a stable conversion. The two effects can be explained by a hydration of the oxide surface. As far as the specific activity of the different catalysts is concerned, the order of catalyst activity is not modified, but the differences are very much attenuated. Figure 2 presents the activity of the different MnO2 samples in the presence of water vapor corrected for the variation in surface area. The order of activity is different from Figure 2: Comparison of the intrinsic activity (towards Ea) of different form of MnO2, in presence of water 25 20 [] 100"C, [] 120"C [] 140"C

~9 lO

9150"C i

I! 160"C, ~

5 0 py-MnO 2 50 m 2/ g

nsu-MnO 2 35 m 2/ g

nsu-MnO 2 100 m 2/ g

rams-MnO 2 18 m 2/ g

781 that attained at the "dry" conditions. Indeed, the conversion measured with 7-MnO2 of 35 m2/g is in between those measured with py-MnO2 and rams-MnO2. The activity of the nsu-MnO2 of 100 m2/g is still above any of the others. 4. DISCUSSION Considering that the metal cation (s), their oxidation state, their first coordination sphere and the crystallite morphology are identical, the differences in activity (figs. 1 and 2) between these different forms of MnO2 are unexpected. The impurity level of all these samples is very low. Miscellaneous data collected throughout our work suggests that the wellknow acid-base modifiers (e.g." K, C1), as well as probably sulphur containing compounds, have negligible effect on activity. The differences of activity between the different samples cannot therefore be attributed to the effect of impurities. Furthermore, the O/Mn ratio in the rams-MnO2 and the nsu-MnO2 samples are identical" this implies that they present exactly the same oxidation state (O/Mn=l.95). The py-MnO2 sample initially contains a slightly more oxidised Mn (O/Mn=l.96). But XPS shows that the surface oxidation state becomes identical to that of the other forms under the reaction conditions (XPS Mn 3s shake up satellite = 4,9 +_ 0.1 eV). The variation of activity reported above cannot be linked to differences in the Mn mean oxidation state. In addition, as mentioned in the introduction, Py-MnO2, nsu-MnO2 and rams-MnO2 are so closely related structurally that the first coordination sphere of the Mn TMions is the same in these 3 phases. A preferential exposure of a given type of site could in principle explain the differences in catalytic behavior. But SEM did not reveal any such differences. Py-MnO2 was constituted of particles in the form of "columns" with an average size of 180xlS0x500 nm; rams-MnO2 particles were more like thick "platelets" of 200x5 nm; nsu-MnO2 particles did not exhibit angles ("pebble form") and the particles had a diameter of 160 nm. The elemental crystallites of the nsu-MnO2 samples were smaller than the apparatus resolution. However, SEM micrographs seemed to indicate that the smallest observable particle was constituted of smaller elements in the case of the 100 m2/g nsu-MnOz than in the 35 m:/g sample. But, none of these differences corresponded to important changes in the development of different type of faces. In summary, the parameters which are usually considered to affect the performances of oxidation catalysts are almost identical. Accordingly, our results point to a new important effect. The only meaningful difference known between these phases concern the crystalline structure and the oxide texture. This is discussed herafler. As far as texture is concerned, the particles size (Table 1) and the crystallite size shown by SEM pictures are comparable for py-MnO2, rams-MnO2 and the 35 m2/g nsu-MnO2. Considering surface area, the exposed surface of rams-MnO2 (18 m2/g), py-MnO2 (50 m2/g) and of the sample of nsu-MnO2 (35 m2/g) could be considered as reasonably similar. But the nsu-MnO2 with 100 m2/g could be considered as substantially different. Its high surface area could be explained by or linked to a higher density of point defects on this oxide surface. This conclusion is consistent with the SEM study results, where the elementary crystallites appeared smaller for this sample. As for crystalline structure, we already mentioned that Mn coordination is the same in each of the samples. But there is a difference in the arrangement of the MnO6 strings. 3D views provided in the figs 3,4,5 attempt to reproduce the ones found in literature [ 1]. In py-MnO2, MnO6 octahedra form single strings linked to each other by shared O octahedron summits. Rams-MnO2 has a chain structure similar to that of py-MnO2. But in rams-MnO2 the chains are sorts of ribbons constituted of 2 strings, which involves the sharing of edges between two adjacent chains. Nsu-MnO2 is a structural intergrowth of py-MnO2 and rams-MnO2 in which

782 layers of single strings and ribbons alternate in a random fashion. The shift from rams-MnO2 to py-MnO2 is actually linked to a change in oxygen coordination. Normal oxygen bulk coordination number is identical, namely 3 for rams-MnO2 and py-MnO2, but the form of the coordination sphere differs. Indeed, in py-MnO2, the 3 Mn linked to one oxygen are placed in a single plane, whereas in rams-MnO=, 2 types of oxygen coordination are encountered in equal amounts: the py-MnO2 planar type (Opl) and a tetrahedral truncated type (Ott) (see fig. 6 and 7). The Opl type is the O linking two ribbons. The Ott occurs at the junction between the two strings constituting the ribbons. The nsutite being a random intergrowth, it presents also the two types of oxygen ions but in variable amounts 0 < Ott/(Ott+Opl) <0,5. From a structural point of view, all the properties of nsu-MnO: would be expected to be in between the ones of py-MnO2 and rams-MnO2. Figure 3" 3D view of the pyrolustite structure Figure 4 " 3D view of the ramsdellite structure (reproduced from [1 ])

(reproduced from [ 1])

9 ~~="~

"~1

o2~r '~'-'..~."11":-.~,; :~

".-'-':25 ~'-.., ~

":'~:'~-

.

~~

~

L'!:"

~

pyrolusnte

ramsdellite

Figure 5" 3D view of the nsutite structure (reproduced from [1 ]) ...... ~,-; ..,-...,. . . . . . . . . . . . . . .

":t,"~'": 9 "~--r ~ ~ : ' ~ . n , z " 9

,r.,,

"

~

"-'~-;'~.i:, .r ~.~'t..}" .~~ ' ~ ;

~'

:

't~

-.

-. " .

.

.~,~-~ :

,,..

-.-~,~'.~,~ -"~.~. ',-'~~."~ " ~ ,

~,

":~

9

-~ ',5".

..

:~

:';

9

.......

"~':':~i~

~.,...l~

"":'~;'~~

~:t

~'-~--~,.'4"

.

" " ","~"~,t'~';

~7,.'.%~"-" "'-'~ -'.'" . . . . . . . . . . . . . . . . . ."-~: ,..'~-.~ ~:~.-

.

.'-.'--'~ " " ;

;

~:.... .K': 9

9

"'

9:"

.3:',g ....

g;.~;. . ..:~ .. ,~ :..."" ~'. '. ':'.'~.

9

"" " -" .... '-" ........

representotive nsutite pyrolusite

i!'"l

...~.....~...~;~,

ramsdellite L..._l

The forms which present the Ott type of sites, namely rams-MnO2 and nsu-MnO2, are more active than the forms which, like py-MnO2, do not. The behaviour of rams-MnO2 and nsuMnO2 in the presence of steam suggest that they are more sensitive to water adsorption. A possible explanation is, thus, that the O in tetrahedral truncated coordination may confer better adsorption properties to the oxide, explaining some of the activity increase. This hypothesis may correlate with the literature findings concerning battery applications, namely that the Ott ions are the preferential sites for proton adsorption [9,10].

783 However, although rams-MnO2 presents more of these O than nsu-MnO2, it is nevertheless less active. All the variations of activity between the different forms of MnO2 cannot therefore be attributed to those O. MnO2 is an essential component of batteries. So, the insertion of electrons into the framework of the different forms of MnO2 has been extensively studied. In batteries, electron insertion is electrically compensated by proton insertion. According to the battery application literature and as evidenced by picnometry [11], non negligible amounts of cation vacancies leading to a decrease of density can be found in manganese dioxides. The presence of some amount of vacancies is considered to be useful for battery applications. Indeed, it causes the Mn 4+ ions neighbouring these vacancies to become more reducible and it creates localized adsorption sites for protons in the bulk which favor their mobility. As a consequence, Ruetschi and Giovanoli[ 11 ] predicted that : "In the presence of cation vacancies, the reaction zone will be moved towards the interior of the oxide phase, allowing higher discharge rates" in batteries. Figure 6 :The Opl type

Figure 7 :The Ott type

In nsutite, some ordering of the anion vacancies can occur. As in the scheme reproduced [ 11 ] in fig. 8, the vacancies are no more randomly distributed in the bulk, but, as confirmed by HRTEM [12], they preferentially cluster in a pyrolusite type single layer. As far as electrode applications are concerned [ 11 ], the ordering of these vacancies is thought to create proton mobility channels and to favour protons and electron mobility. Because of these proton mobility channels, nsu-MnO2 is more appropriate for battery applications than the other forms of MnO2. As far as deep oxidation catalysis is concerned, although nsu-MnO2 structural properties are in between those of the py-MnO2 and rams-MnO2, it is nevertheless more attractive than the other forms of MnO2. This better performances cannot be linked to long range cooperation (such as remote control [13]) between py-MnO2 and rams-MnO2-1ike domains in nsu-MnO2 since the mechanical mixture of these two phases did not show any improved performances. This explanation probably rests on the fact that, of the 2 nsu-MnO2 samples, the best for deep oxidation catalysis is the one for which textural and SEM characterizations suggest a higher density of cationic vacancies. Our results thus suggest that a strong parallel exists between battery applications and the catalytic activity of the MnO2 phases. The clustering of the cationic vacancies seem to considerably improve the catalytic properties of MnO2, as it does for the properties measured for batteries. The higher the number of vacancies, the better the performance. This improvement of the catalytic activity is certainly partly linked to an increase of the

784 M n O 2 oxygen

lability. Indeed, vacancies locally affect the O lability: the coordination number of the 6 0 surrounding the Mn vacancy is decreased. As admitted for battery electrode applications, Mn ions neighbours become thus more reducible. The complete oxidation of VOC is a reaction demanding many O. So, Mn vacancies (if they are situated close to the surface) offering 6 labile O are likely to take an active part in the catalytic process. However, the removal of O has to be compensated by electrons movement. So, what favours the electrons mobility should also certainly confers to the MnO2 better catalytic properties. This is undoubtedly the reason why the clustering of the vacancies in the bulk~ which favours the electronic conductivity by increasing the proton mobility, so dramatically improves the catalytic performances of the nsu-MnO2. Figure 8: schematic view of the proton channel formed by the clustering of anionic vacancies. (reproduced from [ 11])

l_ = --

~

9.3s

Z~

I VACANCIES

t

Schematic view of o cation-deficient intergrowth structure of ramsdellite and pyrolusite, with vacancy clustering in the pyro|usite zone. Projection onto the ( 0 0 | ) plane.

5. CONCLUSIONS To conclude, the comparison of the activity of manganese dioxides of different structures shows that very similar oxides can nevertheless exhibit quite different catalytic properties. As far as manganese dioxides are concerned, the changes in the catalytic activity from one oxide to another must be attributed to the difference in coordination for the oxygen anion and to the presence of (clustered) anionic vacancies which modifies the electronic and protonic conductivity of these oxides. REFERENCES 1. A.M. Potter, G.R. Rossman; Am. Mineralogist, 64 (1979) 1199. 2. D.A. Shirley, Physica Scripta, 11 (1975) 117 3. B.W. Veal, A.P. Paulikas, 51-21 (1983) 1995

785 4. F. Kapteijn, L. Singoredjo, A. Andreini and J. A. Moulijn, Appl. Catal., B3 (1994) 173. 5. Catalytica environmental report E4 (1993) 6. C. Lahousse, A. Bernier, P. Ruiz, P. Grange, B. Delmon and P. Papaefthimiou, T. Ioannides, X. Verykios, 5thworld Congress of Chemical Engineering, July 14-18 1996, San Diego, 3 (1996) 493. 7. C. Lahousse, A. Bernier, P. Ruiz, P. Grange, B. Delmon and P. Papaefthimiou, T. Ioannides, X. Verykios, European Workshop on Environmental Technologies, November 13-15 1996, Copenhagen 8. C. Lahousse, A. Bernier, P. Ruiz, P. Grange, B. Delmon, to be published. 9. W.C. Maskell; J.A.E. Shaw; F.L. Tye; J. Appl. Electrochem., 12 (1982) 101; Electrochim. Acta, 28 (1983), 225. 10.J.J. Coleman; Trans. Electrochem. Soc., 90 (1946) 545. 11.P. Ruetschi; R. Giovanoli; J. Electrochem. Soc., 135-11 (1988) 2663. 12.S.Tuner; P.R. Busek, Nature 304 (1983) 143. 13.L-T Weng, B. Delmon, Applied Catal. A, 81 (1992) 141.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

787

U n d e r s t a n d i n g the surface c h e m i s t r y for suptx~rted v a n a d i u m oxide systems m o d i f i e d w i t h p h o s p h o r u s oxide at h y d r o c a r b o n s o x i d a t i o n V.&Zazl~alov, L.V.Bogmskaya, L.V.Lyashcnko and I.V.Bachcrikova Ul~amian-Polish Laboratory of Catalysis, Institute of Physical Chemistry, National Academy of Sciences of Ukraine Pr. Nauki 31, Kyjiv-22, 252022 Ukraine

1. INTRODUCTION Despite widespread commercial use a n d considerable research efforts, the vanadium-phosphorus oxide multifunctional catalytic system rcmains poorly understood_ It is 8cncrany accepted that (VO)2P207 bulk phase is the most selective one to produce maleic anhydride (MA) from n-butane [1]. However, in spite of the large number of studies on the VPO catalysts thcrc arc still many unclear things and disagreements m scientific literature. Thc catalyst surface composition and structure under reaction conditions arc difficult to study and still remains unrcsolvcd. Due to numerous compounds possfoly synthesized in VPO system depcndin8 on thc preparation mcthod and their easy changes under reaction conditions it is hardly rcalizcd desircd active centers structure reproduction ovcr the supports surface. In the present paper wc attempted to undeTstand the surface chemistry for supported on acrosil vanadium oxide systems modified with phosphorus oxide at benzene and n-butane oxidation. As for the latter reactions there arc some investigations on usin8 vanadiumphosphorus oxide supported on silica [2-4], titania [5, 6], alumina [6-8], A1PO4 [9]. The use of supported catalysts gave not so promising rcsults. It was found that the prcscncc of V4~ probably increased the selectivity to malcic anhydride. The titania-supported catalysts showcd a hish dispersion of VPO compound over the surface but they were very selective towards COx. The use of SiO2, Al~O3 and AIPO4 as supports increased thc activc surface arca without remarkable cffcct on thc catalytic pcfformancc. All thc invcstisatcd supported VPO systems did not give much information to make their surface chemistry clear yet. As for the reaction of benzene partial oxidation, no one supported VPO system was tested bcforc. Ncvcrthclcss, such catalysts could bc quite suitablc to cstimatc active sites nature and role of thc catalyst composition, to study surface chemistry and hydrocarbons convcrsion mcchanism.

788 2. EXPERIMENTAL 2,1, Preparation of supported VPO catalysts Two series of grafted VPO samples wcrc prepared usm8 aerosil-200 (Degussa) preheated at 250 ~ (PVD catalysts) and 600 ~ (SVP catalysts). The surface OH-sroulz concentration was 0.78 and 0.48 retool/8, respectively. Solutions of POC13 and VOC13 m cc]4 were added m a desired ratio to CC14 s u s ~ o n of aerosil in time difference of 15 ram. After the solvent removal a hydrolization with wet air was carried out at 200 ~ for 3 h (samples PVD-4, -5, -6). After their calcination at 200 ~ for 2 h no more chloride was found in the desorption products. Sample PVD-3 was prepared by mcam of consecutive graflin8 of phosphorus and vanadium salts with hydrolizatiom at the mtmmcdiate and final stages. The catalysts of PVD-senes had atomic ratio fp = P/(P+V) close to 0.6. SVP samples were prepaxcd by the similar to PVD samples' way except for hydrolization procedure: it was caxlicd out with wet air first at a room tcml~ratmc directly in the mother's sohrdon and then one more time at~r removal of the solvent at 200 ~ For these catalysts, havin8 different phosphorus content, vanadium concentration has bccn maintained at 0.8 wt.~. Imprcsuatcd catalysts (IPV) wcrc prcpaxcd by ~ of ammomum vanadatc aqueous solution and orthophosphonc acid with acrosil suspension at the presence of oxalic acid followed by hcatm8 under stimn8 and further calcination of the solids at 600 "C for 2-3 h (fp=0.6). All the catalysts prepared were pelleted, cnmhcd and sieved_ The catalytic tests wcrc pcffonncd on thc 0.6-1.0 m m flaction. 2.2. Catalysts ~armterization ESR spectra of the catalysts were recorded on Vanan E-9 radiospectrometer in the X-ray region. The degree of reduction of the vanadium has bccn determined by potenhomctnc titration with 0.01 N solution of (znum (IV) sulphate and calculated as V44"/V~t in pcrccnt. The surface of the supported VPO systems wcrc studied by SEM with MPA (Jr ffSM-35 SF and Ormk EEDS-II) and XPS (VG ESCA-3, A1 Kcq~, standard C ls B.E. = 284.8 eV) m much the same manncr as it was described in [10]. Study of the samples by EXAVS was carried out analogously to that published m [11]. Some of thc catalysts, namely imprcsnatcd ones, wcrc analysed by XRD method (DRON-3M). 2.3. Catalytk performance testing The catalytic oxidation of b e n ~ n e (0.5 % m air) and n-butane (1.5 % m air) was earned out in flow reactor operated at atm~phcnc pressure. The tests were done on a constant catalyst volume basis (2 cm ) m a quartz reactor of mtcmal diameter 1.0 cm. The reactions were performed m the temperature ransc of 300550 ~ at thc flow rate of 50-200 cm3/mm. Thc analysis of the substratcs and products was pcffoImcd by gas chromatography.

789

3, RESULTS AND DISCUSSION

3.1. Impre~s~ed smnples Bulk vanadium pcntaoxidc is quite activc but low sclcctivc catalyst of hydrocarbons partial oxidation. It was established by XRD that the hi8hcr content of the phosphorus additive in it [12] the weaker pcal~ attributed to V205 in bulk catalysts and, simultaneously, p-VOPO4 phase rcflcctiom appearccL The latter became the major component of VPO catalyst at fp > 0.67. The constituents of the prepared sample wcrc found to bc also (VO)2P2OT, VO(PO3)2 and some amorphous compounds. At this takes place, an the cations wcrc considered by authors [12] to bc bonded in vanaclyl groups V - O and phosphorus atoms form Bromtcd acidic center each. It has bccn found an increased concentration of phosphorus over the surface as compared to the bulk and the hishcr phosphorus content m the sample the 8town bulk concentration of the reduced vanadium ions were observed. As for V-P containin8 samples prepared by impregnation method, one can expect predominated mamtmnin8 of the supported compounds structure over the surface which would be very similar to those synthesized in bulk samples. On the other hand, it is well known that amon8 oxide supports SiO2 is the weakest dispersant of vanadium oxide over the surface [13] and always there is some V2Os phase-like formations despite the very low surface occupation de8rcc [11,14,15]. Quite limited information on the mixed vanadium-phosphorus catalysts supported on silica arc represented m open scientific literature. For cxamplc, in [4] it was proposed the method to prepare well dispersed VPO composition over SiO2 but, unfortunately, the authors did not ttun out successfully to determine the chemical composition of the surface phase formed (supposcly VOPO4 to8cthcr with VOx). In [3] r 4 as a major phase of the similar samples was d e s c ~ d and r toscthcr with umdenfilicd compound were found to be constituents of the surface layer in [6]. In the present paper, impregnated IPV samples on acrosil, containin8 5, 10 and 15 wt.% vanadium at fF of about 0.6 were prepared. All the XRD patterns for the samples arc proved to be very similar and contain the most characteristic reflections of [~-VOPO4, cc-VOPO4 and (VO)2P~O7 phases a8~in.~t the considerable halo-effect. By means of scannin8 electron microscopy it has been established (Figure 1), that there are crystal resions over the surface of IVP imprc8natcd samples, winch resemble the structure of ~-VOPO4 phase. These regions size arc defined by supported compounds amount and the more active components supported the lar8cr domains formed. Microprobe analysis in 10 sites on the surface showed the ratio V/P m these domains was close to that observed for bulk VOPO4 phase (see Table 1). It can be asstmacd that thermal treatment of the catalysts as prepared causes the phosphorus migration over the surface to produce I~-VOPO4 phase and the rest phosphorus and vanadium form the prcerystal compounds or remain unbonded toscthcr. The microprobe analysis of the surface between the mentioned above VP oxide domains confirmed this asstmaption (Table 1). The

790

~,~

Figure 1. Microphotographs of the impregnated and graBcd VPO/SiO2 catalysts surface: a -bulk p-VOPO4, b - IPV-3 (15 wt.% V), c - IPV-4 (5 wt.% V), d - PVD-5 (4 wt. % V), c - PVD-3 (2.72 wt.9~ V).

791 binding cn~gy of V 2p-clcctrom lacing 517.2 cV (B.E. P 2Pv4~ 133.7, O ls 532.4, Si 2p = 102.6 r can br indicative of a presence of ions on the samples surface d ~ to its lower value as compared to that for ~-VOP04 bulk phase (518.6 r It is conceivable that these vanadium ions do not r into the composition of crystal formations and they arc part of prcc~stal structures. Tablc 1 El cnts ratio on the surface of imprcgnatcd and graftcd VPO/SiO~_ samples a I ~ microprobe analysis Preparation Concenmethod tration of vanadium, wt.% Iml~csnahon Grafting

[3-VOPO4

Bulk

Ratio V:P:Si on the surface

P/V

ratio

15.0 5.0 4.0 2.7 1.8 -

MPA

1.2 1.2 1.4 1.5 I.I 1.0

XPS

C~taIs

Support

1:1.10:0.52 1:1.05:0.70 I: 1.02:0_90 I:I.II:0.96 1:1.02:1.08

1:1.24:3.02 1:1.16:5.96 1:1.34:6.04 I: 1.38:6.45 I:I. I0:7.32

I:I.03:0

-

1:1.36:4.07

1:1.32:5.25 -

I:I.08:0

Table 2 Properties of impregnated IPV and bulk VPO catalysts m the reactions of benzene and n - b ~ c oxidation*

Catalyst

IPV- 1 IPV-2 IPV-3 I3-VOPO4 tt-VOPO4 V20s

(VO) 2P207

Vanadium on support, wt.%

SSA,

5 10 15 -

120.5 71.4 25.3 2.0 -

-

4.5

-

m2/g

4.2

Benzene oxidahon

Tm~, *C 450 420 460 550 500 -

410

Sin, % 65 68 58 60 56 0

55

n-Butane oxidation Sst~, 17 20 16 17 35 I

54

* - Tmx - tcml)~aturc of the maximmn of malcic anhydride yield; selectivity to malcic anhydride

Sst~

792 A correlation between bulk and supported vanadium phosphorus catalysts properties (Table 2) shows much higher activity of the latter samples as related to former ones. The most active catalyst is IPV-2 where VI'O microerystals over the surface are of an average size. The f m ~ e r growth of these formations size leads to decrease of the catalytic activity and selectivity (see sample IPV-3 in Table 2). From the catalysts efficiency comparison it can be concluded that no mdividtml phase tested is responsible for the supported samples catalytic performance. However, when the crystal formations on the surface arc large c n o u ~ the properties of the samples draw near the bulk compounds activity and selectivity. 3,2, Properties of gralted VPO catalysts, At simultaneous graltin8 of vanadium and phosphorus compounds followed by hydrolysis of the surface grOUl~ (samples PVD-4, -5, -6) the components can be fixed on the surface accordin8 to the reac~om:

H20

>Si-OH + EOC1,--> >Si-O-EOC12 + HCI . . . . > >Si-O-EO(OH)2,

)Si-OH

+ EOCI~ -->

)si-o\

2

)si-5

si_o/Roci + 2I-ICa .... >

"

si-o

o(oI-I),

~

Si-OH )Si-Ox Si-OH + EOC1,--> ) S i - O - - ) E = O + 3HC1, Si-OH ~Si-O/

where E = V or P. Table 3 Catalytic properties of the grafted PVD samples m benzene and n-butane oxidation*. Catalyst

Benzene oxidation Vanadium fp = content, wt.% P/(P+V)

Xuc, ~; PVD-6 PVD-5 PVD-3 PVD-4

2.78 4.00 2.72 2.87

0.50 0.58 0.60 0.67

90.0 89.6 89.0 90.5

S~, ~ 60.0 62.0 72.0 74.6

n-Butane oxidation

S~, % 19 20 24 22

W.10v 9.4 8.8 9.1 9.3

* X H C - hydrocarbon conversion, SMA - selectivity to malcic anhydride, W. 10~ rate of n-butane oxidation, mole/s.atom V

793 From the Table 3 it can be seen that the total m o u n t of grafted compounds exceed that of OH 8roulm on aerosfl surface. It can be explained by additional centers folvning dl~ to either partial dissociation of the surface siloxanc groups or the mixed (V-O-P) structmca formin8 during the grafting. At the consecutive g t a f l i ~ of the phosphorus and vanadium compounds the latter can bc fixed to not only fzcc OH grOUl~ but also to nascent P-OH groups to produce mixed P-O-V structures. Due to close catalytic properties of the samples prepared by simultaneous and consecutive grafting one can conclude that the obtained in both cases P-O-V structmes arc very similar despite the order of grafting. An investigation of their surface by microscopy supports this idea (scc Figure 1). Over the surface of the grafted samples island crystal structures, similar to those for bulk 13-VOPO4 and impregnated catalysts, arc observe& The mcasurcd by XPS clcctrom bmdin8 cncrgics for all the clcmcnts of PVD-5 sample dillcr ncgli81bly from those found for impregnated catalyst. At the vanadium concentration reduced to 0.8 wt.% (samples series SVP) such crystal structmcs arc not found_ In the last case, an influence of phosphorus atomic fraction fp at the constant concentration of vanadium over the surface on the prepared samples p r o p o s e s were studied (Table 4). Table 4. The suffacc state for SVP catalysts and thc sclcctivity towards malcic anhydridv at benzene and ~ - b ~ a n c o~daho~ Catalyst

fp

Degree of Content V~ wt.9~ SiOH $roulm covcra$c before after rcaction rcaction

Oxidation* bcn2~nc, SMA, %

n-butane SuA,

SVP-0 SVP-3 SVP-4 SVP-1 SVP-8 SVP-7 SVP-9

0 0.27 0.49 0.62 0.71 0.82 0.89

0.33 0.48 0.66 0.85 1.21 1.72 2.77

0.09 0.09 0.12 0.15 0.20 0.38 0.55

0.11 0.13 0.43 0.69 0.63 0.45 0.55

36 57 52 0 0

4 6 7 33 36 18 19

E 205 202 170 120 126 182 176

* - SMA - selectivity to malcic anhydridv, E - activation cncrsy of n-butane oxidation, kl/molc

An introduction of phosphorus to vanadium-containing aerosfl leads to significant changes m its catalytic properties. In so doing, an increase of the phosphorus concentration (up to certain value) provokes growth m both activity and sclcctivity of the obtained samples at hydroca~ons oxidation (scc Table 3 and 4). Low occupation dcsrce as compared to OH g r o u ~ concentration favours

794 vanadium fixation sepazately from phosphorus or/and little amount of mixed PO-V structures formins. At the essential excess of phosphorus over the optimum amount the catalytic performance changes for the worse. An activation cncrsy for n-butane oxidation rcachcs its minimum value over thc most active samplc and rises asain with increase of the phosphorus concentration. F_~SRspectra (Figure 2) recorded for the samples with relatively low phosphorus concentration shows high rcsolulion structure attributed to V4t ions at their low concentration and an ~ c c of cxchansc interaction between them. With increase of the phosphorus concentration m the samples V4t ions amount growth and cxchansc interaction between them or in V4~-V~+ pairs appeaxs. Further increase (over optimum) phosphorus content leads to rcappcaxancc of high resolution structure of the spectra.

/

Figure 2. ESR spectra of the grafted V-P/SiO2 catalysts: a - SVI'-4 (fp = 0.49), b - SVP-1 (fp = 0.61), c - SVP-9 (fF = 0.89), d - SVP-8 (fp = 0.71).

795 Compann8 V4+ ions concentration in the samples with cor~cspondm8 specific reaction rates and selectivity towards malcic anhydride one can notice their symbatic growth (except samples having fp > 0.7). T h o u ~ V4+ ions concentration arc found to play a decisive role m selective oxidation, it is important but not sufficient requ~ement for the VPO/SiO2 catalysts to be selective (Table 4). As for the participation of intermediate O~- and O species m selective run of the hydrocarbon oxidation, there arc some doubt on such mechanism. The mason for this conclusion can bc the observed increase of the proccascs selectivity with the decrease of 0 7" and O- concentration [16] and reaclmtg the maximum selectivity at surface oxygen radicals content of about zero. So, it should be assumed mechamsm for the hydrocatbom selective oxidation revolving oxysen from surface grou~ V=O. The same conclusion has been drawn for the bulk V2Os-P205 catalysts in butene-1 [7] and b e ~ e oxidation [17]. The presented results on grafted VPO composition allows to assume that at the low phosphorus content oxide clusters of V20~ type predominate over the surface. As for the samples containing only supported vanadium, it should be mentioned that the catalyst prepared using aerosil preheated at 250 ~ and having 0.78 mmol/8 OH groups displays catalytic properties very similar to those of the bulk V205. At the same time, another preparation using aerosil preheated at 800 ~ (0.25 mmol/g OH groulm) catalizcs the benzene selective oxidation to produce about 19 % tool. yield of M& The best catalytic performance was found to be the property of grafted/ impregnated systems having fp = 0.65-0.72 for all the tested reactions. When fp > 0.7 the activity and selectivity sharply decreased due to isolation of vanadium ions and fmlhcr covering them with phosphorus-oxide groups. Study of an the grafted samples by EXAFS [18] showed the presence of peaks attiibutcd to V-O, P-O, V-O and O-O bonds length m p - v o P 0 4 phase. Moreover, some of the lines (0.400, 0.517, 0.650, 0.695 nm) were indicative for the presence of several ordered layers of this phase. Conclusively, it should bc emphasized that at the vanadium and phosphorus g r ~ on the surface of acrosil mainly p-VOPO 4 phase flasmcnts arc folmm8 for the V-P/SiO~. efficient catalysts. At the same time grafted (VO)2P207 phase foxtmng was failed, probably due to its space structmc which cannot bc created on the surface. REFERENCES 1. Vanadyl pyrophosphatr catalysts (Edit. O.Ccnti), Catal.Today, 16 (1993) 1. 2. V.A Zazhigalov, YtLP. Zajtscv, V.M. Bclousov, B. Parlitz, W. Hankc and O. Ohlmann, React. Kinet. Catal. Lett., 32 (1986) 209. 3. KE. Birkcland, H.H. Kung, S.R. Barl, O.W. Coulston, R. Harlow and P.L. L ~ , Hctcrog.Hydrocarbons Oxidation, 21 lth Na~onal Mo~t. Amcr. Chem. Soc., New Orleans, 41 (1996) 197. 4. RA. Overbeek, A.RC.L PekelhaxinS, KJ. van Dillen and ff.W. Geus, Appl. Catal. A, 135 (1996) 231. 5. R & Ovexbeek, P.A Wamnsa, M.J.D. Crombas, L.M. Vmser, AJ. van DBlen and J.W. Gem, Appl. Catal. A, 135 (1996) 209.

796 6. M. Maxtmcz-Laxa, L. Momno-Rr R Pozas-Torm0, A Jmaencz-Lopez, S. Bruque, P. Ruiz and O.Pocelet, Can. L Chem., 70 (1992) 5. 7. M. Nakamura, K. Kawai and Y.Ftowara, L Catal., 34 (1974) 345. 8. RamstcRcr and M. Batrrns, J. Catal., 109 (1988) 303. 9. P.S. Kuo and B.L. Yan8, L Catal., 117 (1989) 301. 10.V.A_ Zazl~alov, J. H ~ , J. Stoch, h~I. Pyatmtskaya, O.A_ Komashko and V.M Br Appl. Catal. A, 96 (1993) 135. 11.V.& Zazt~alov, V.M. Br R Kozaowski and YuLP.Zaitscv, Tcorct. and Expenm Chem., 23 (1987) 650. 12.A. Satsuma, K Hatton, A_ Furuta, A_ Miyamoto, T. Hattori and Y. Murakami, I. Phys. Chem., 92 (1988) 2275. 13.M. Niwa, Y. Matsuoka and Y. Murakami, I. Phys. Chem., 93 (1989) 3660. 14.Z. Roozeboom, M.C. Mittelmeijer-I-Iazaleser, J. Medema, V.H.L de Beer and P.J. Gellings, I. Phys. Chem., 84 (1980) 2783. 15.M. Dercwinsld, I. Haber, R Kozlowski, W.A_ Zazhigalow, J.P. Zajcev, I.W. Bachcrikowa and W.M. Bclousow, Bull. Polish Acad., Sci. Chem., 39 (1991) 403. 16.R. Frickc, H.-G. lcrschkcwitz, O. Lischkc and O.Ohlmatm, Z. anor8, a118. Chem., 448 (1979) 23. 17.& Satsuma, F. Okada, & Hatton, & Miyamoto, T. Hattori and Y. Murakami, Appl. Catal., 72 (1991) 295. 18.V.A_ Zazl~alov, V.M. Belousov, I.V. Bacherikova and Yu.P. Zaitsev, Teoret. and Exl~rim. Chom., 27 (1991) 370.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

797

Effects of cesium doping on the kinetics and mechanism of the n-butane oxidative dehydrogenation over nickel molybdate catalysts L.M. Madeira* and M.F. Portela GRECAT. Instituto Superior T6cnico - Technical University of Lisbon Av. Rovisco Pais, 1096 Lisboa Codex 0~ortugal)

The kinetics of the ODH of n-butane has been investigated for unpromoted and cesium

promoted tx-Ni]~oO4catalysts. The reaction rates of dehydrogenation products as functions of the butane and oxygen partial pressures are described by a kinetic model based on the Mars and van Krevelen mechanism. The effects of Cs on the kinetic parameters can be interpreted on the basis of recently published results concerning the properties of those catalyst3. 1. INTRODUCTION Nickel molybdates proved to be promising catalysts for n-butane oxidative dehydrogenation (ODH) [ 1,2]. A detailed study about the effects of alkali metal promoters (Li, Na, K and Cs) showed that selectivity for dehydrogenation products (C4's) can be increased by suitable modification of the nickel molybdate catalysts with those promoters, specially with cesium [3]. In this way, the effects of doping them with increasing Cs loadings (1, 3 and 6 %) were carefully investigated [4,5] and a maximum of selectivity for dehydrogenation products was found with 3 % promoter loading. Several kinetic studies have been carried out with lower alkanes but they were usually performed with vanadium and phosphorus containing catalysts. According with Mamedov and Cort6s Corber/m [6], the ODH of lower alkanes on V205 - based catalysts appears to occur by redox mechanism but it is evident that no clear interpretation regarding the alkane activation can be derived from the literature. Other studies with the V-P-O system showed that the oxidation of butane to maleic anhydride occurs by redox reaction kinetically consistent with the Mars and van Krevelen model [7,8]. For this system, the participation of lattice oxygen was demonstrated [9] and when oxygen is consumed in the reaction, the vacancy is filled by gaseous oxygen. Centi et al. [ 10] also found that a redox process is involved in the ODH of nbutane on a V-P mixed oxide. In what concerns to nickel molybdate catalysts, a kinetic study showed that the propane ODH could proceed by redox mechanism [11 ]. For improving our knowledge of the factors determining the kinetics and mechanism of nbutane ODH, and due to the small amount of available information for the Ni-Mo-O system, a systematic kinetic study with unpromoted and 3% Cs promoted NiMoO4 was undertaken. "L.M. Madeira thanks PRAXIS XXI Program from JNICT (Junta Nacional de Investiga~o Cientifica e Tecnol6gica) for financial sup~rt.

798 2. EXPERIMENTAL

2.1. Catalysts preparation and characterization The ct-NiMoO4 catalyst was prepared by eoprecipitation [2] and afterwards doped by wet impregnation with a solution of cesium nitrate The impregnated sample was filtered, dried and finally calcined in air for 2 h at 550 ~ The catalysts were carefully characterized by several techniques such as BET, ICP (inductively coupled plasma spectroscopy), AA (atomic absorption), HTXRD, FTIR, XPS, CO2-TPD, TPR and electric conductivity. Experimental details and results can be found elsewhere [3-5,12].

2.2. Oxygen temperature programmed desorption experiments To investigate the interaction of oxygen with the catalyst surface, several TPD experiments were undertaken. A pretreatment was first performed by heating the sample in He flow up to the desired temperature but avoiding the transition to the metastable 13-phase, which is already present at 595 ~ [13]. After 02 adsorption at different temperatures the sample was cooled to room temperature, purged in He and finally 02 was desorbed by heating (10 ~ in a He flow (1 ml/s) with on line gas analysis performed with the TC detector of a gas chromatograph.

2.3. Catalytic tests Catalytic runs were performed in a fixed bed, continuous flow tubular quartz reactor with a coaxially centered thermocouple. Reactants and products (1-butene, trans-2-butene, cis-2butene, butadiene and carbon oxides) were analyzed with an on line gas chromatograph with two columns [3,4]. The catalyst was mixed with inert quartz (180-254 ttm) in a volume ratio of 1:2 catalyst to quartz. The catalyst charge was chosen to yield differential conversions, in order to minimize the effects of products and secondary reactions. With unpromoted NiMoO4 the catalyst charge was 0.150 g and the contact time 4.45 g.h/(mol butane). With the cesium promoted sample W = 0.300 g and W/F = 5.80 g.h/(mol butane). The feed was a mixture of nbutane, 02 and N2 and for both catalysts, the influence of the reactants partial pressure was studied. The total pressure in the reactor was 1.10 bar and the temperature range studied was 500 - 560 ~ Blank runs proved that under the experimental conditions used the homogeneous reactions can be neglected. The stability of the NiMoO4 activity was previously checked for 50 h and a good stability in butane conversion and products distribution was found [ 14]. 3. RESULTS

3.1. Catalysts characterization Characterization data evidenced that the prepared NiMoO4 is stoichiometric and that Cs is deposited only on the catalyst surface (atomic ratio Cs/Mo = 0.03) not affecting the molybdate structure. However, Cs doping causes a decrease of the catalyst surface area: SBET(NiMoO4) = 44.1 m2/g and SBET(3% Cs-NiMoO4) = 28.7 m2/g. Moreover, the promoted sample exhibits a higher surface basicity, electrical conductivity and also a larger resistance to reduction [4,5,12].

3.2. Oxygen temperature programmed desorption experiments Adsorbed oxygen on a stoiehiometric surface of a oxidized oxide is readily identifiable because it is usually desorbed at temperatures lower than the sublimation temperature of

799 surface lattice oxygen [9]. One method to detect the presence of adsorbed oxygen is by temperature programmed desorption (TPD) after exposing the oxide to 02. While MoO3 does not adsorb oxygen, NiO exhibits an O2-TPD profile with several peaks, corresponding to different forms of adsorbed oxygen, either molecular or atomic [15]. However, we have not found in the literature any reference with respect to nickel molybdates. In this way, several O2-TPD experiments were performed with unpromoted and promoted NiMoO4 with different experimental conditions such as the kind of pretreatment, the way of 02 adsorption or even changing the adsorption temperature because it is well known that for certain species the adsorption is activated [9]. Nevertheless, no adsorption of 02 was detected. Only a slight increase in the TCD response at temperatures higher than 600 ~ was observed which can be due to the sublimation of lattice oxygen. 3.3. Catalytic tests To investigate some aspects of the kinetics of the reaction as well the effects of Cs doping, a systematic study over the unpromoted and 3% Cs doped NiMoO4 catalysts was undertaken for n-butane ODH by changing the temperature and the reactants partial pressures (Pi). However with the unpromoted NiMoO4 it was not possible to operate with butane pressures higher than 0. l0 bar, when Po2 was fixed at 0.05 bar, due to catalyst reduction and/or coke formation. Therefore, some tests were performed with Po2 = 0.15 bar but again, catalyst reduction and/or coke formation was observed at butane pressures higher than 0.15 bar. With the Cs promoted sample we worked with Pb~t~ up to 0.25 bar (with Po2 = 0.05 bar) without significant reduction of the catalyst or coke formation. The influence of the reactants partial pressure in the dehydrogenation products formation rates (based on surface area) are shown in Figure 1 for NiMoO4 and in Figure 2 for 3% Cs - NiMoO4. It is noteworthy that a zero order in O2 is apparent for all C4 products with both catalysts. Moreover, a positive order (smaller than one) in butane is visible with NiMoO4 while an almost linearity is found with the Cs doped catalyst, specially with 2-butenes. In what concerns to products distribution, the effect of temperature is similar with both catalysts and corresponds to an increase in butane conversion: the butenes selectivity decreases as the butadiene selectivity increases and the carbon oxides formation also increases, specially CO. These effects are more pronounced with the Cs doped sample. Butane partial pressure does not affect the products distribution with NiMoO4 but increases the C4's selectivity (specially butenes) decreasing mainly the CO2 formation with the 3% Cs doped catalyst. The effects of increasing Po2 are the same for both catalysts. A decrease of C4"s selectivities and an increase of COx formation at low pressures is mainly observed. It is noteworthy that the main effects of Cs doping in the selectivities are increase in 1-butene selectivity and decrease in CO formation. 4. DISCUSSION It is well known that in oxidation reactions oxygen acts by adsorption on the oxide catalyst as O, Os etc. or incorporation as lattice O 2 species. The solid is oxidized in this step, and the electrons received by adsorbed oxygen could come from reduced surface cations or anionic vacancies with trapped electrons. If oxygen is incorporated as lattice oxygen ion, the sites for adsorption of oxygen and for oxygen attack in the catalytic reaction may be different, and migration of oxide ions in the solid between the two sites would occur [9].

800

~ - -

L ~ L

1.0E~4

0.05 bar

x _~

x

pc~= O05 bet

1.2D04

":

A

pc~=Q15bar

,= ~0E~5

~ S,~~ 8 4

0,0E~ 0,00 7,~E-05

0~1~C10

o,05

,

~

*

t

,

~

,

,

ao5

I

,

~

0,0E~

a~o

P t ~ , (bar)

&0E-0~

F~u~, = (I05 mr

E

~

am

0,10 0,15 F~ (t~)

aoo ~0E-~

P=-0.05bar

ao5

a~o

P ~ , , (bar)

a~5

am

Pcz - 0.15 I:=r

x

o

o

,,,,,,

....

(~oo 7,~-,.05 E ,'-: 5,0E,.05

o,o5

~1o q15 Pc= ~a0

.._

0,~,

q2o

pbuw== gO5 i:=r

x=_..__..___~~

: : : : I : : : : : ~ : o,oo o.o5 O.lO 8m,,(t=0

~0~05

P o ~ " 0.05 bar

9,0E~5 -

~

~

0,05

1,6E-04 --'-o 8,0E-05

0,0E+E)

~ A ~

I

0,05

.

I

t

0,10 0,15 Pc~ ~a~

: : ; : ; : : : ; ; : : 0,00 0,05 0,10

0,0E-Kg) 0,00

P = m (ti)

..-.

.

Pc~ = O 15 bar

~0E~4-

P~-O.05bar

x

~0~04 T

/

9

c~ 1,(E-04 "O

I

0,0EH~,

0,20

Pmana" G05 bm

o

=-.

0,0E*(I),

0,20

pha~=0.05ber

0,00

~0E~04

0,10 0,15 Pc~ 0~a0

0,00

"O i,. ::

~5E~04-

qo5

O,lO

I---r-=---1----~

=.. O,0E+(]0

0,00

I

0,05

I

0,10 Paz~

I

0,15

I

0,20

0,0EN:}0/

0,00

po==0.05ber

6,QE-04

E

E

J=..

J=_ (~0E-t{]~

0,00

~

I

t

I

I

I

m

i

~

I

o,o5 qlO Pt,am (bar)

t

I

o,2o

Po~-(1151:ar

t

I

0,(IS+{]0

o,oo

t

0,05 0,10 0,15 Ra=m ~a~

x

I

0,05

I

I

A - 540 ~

x - 560 ~

I

qlO o,15 o,20 Pu=~ Oa~r)

Figure 1. Influence of reactants partial pressure on the dehydrogenation products formation rates over NiMoO4 (e - 500 ~ 9 520 ~ cq. (7) with the data presented in Figure 3.

I

0,21)

po~=Q15b ~

E ~ 4,0E-04

....

o,15

Pt,,-,,(bar)

+

::;~: : : ; : '. 0,05 0,10 Raam~

.=

o= ZOE.O4

o.2o

x

r~~6'0E~5 ~ 3,~--05 ~ -

=.. 0,00

o.o5 O.lO o . 1 5 P=m (t~

E

x ,,E~ ~ 40,E 0-5,

~OE+~

0,0E+G) o.oo

Full lines obtained from

801

8,0E-05

Pbutane

T

/

N

x

x

~ 4,0E-05

Po2 = 0.05 bar

= 0.05 bar .

~"

___JK

o E

m

9

2,0E-04

1,0E-04

T=...

:

1

0,0E+00 0,00

! i 0,05 0,10 Po2 (bar)

4,0E-05 -1-

P~tane = 0.05 bar

2,0E-05

~

I

9

~,

:

I

5,0E-05

~~

t I 0,05 0,10 Po2 (bar)

T

Pbut=~ = 0.05 bar

t,

x

x

1,0E-04

T

[

0,0E+00 0,00

I 0,15

2,0E-04 x

9

I 0,05 0,10 P02 (bar)

x

=j9

,

T

"S

.

.-. 1,0E-04

! ~

t

t 0,0E+00 l 0,00

I

0,05 0,10 Po2 (bar)

X

9

9

9

0,15

0,20

0,25

0,20

0,25

0,20

0,25

0,20

0,25

Po2 = 0.05 bar

T

0,0E+00 0,00

0,05

3,0E-04

Po2 = 0.05 bar

0,10 0,15 Pbut=~ (bar)

g to 1,0E-04

.~

O,OE§

i

0,15

1,0E-03 T x

9

9

9

I 0,05

0,00

&

_ -

t 0,10 P 0 2 (bar)

0,05

0,10

0,15

Pbun~ (bar)

Pbut=~ = 0.05 bar x

0,10

Pb.t~,, (bar)

~ ~,o~-0,

9

0,00

0,05

~

K"

_=

,,-,, 2,0E'--04 ~

0,15

x

O,OF+I~

0,25

9

Pbutane= 0.05 bar

~

E 5,0E-05 o9

0,20

Po2 = 0.05 bar

T

, O,IXI

0,15

.,~~,OE-O~ = l

~

0,0E+O0 i

0,10

~ 1,0E-04 J

9

i

0,05

Pbutane (bar)

.... 1,5E-04

_=

0,0E+00 0,00

0,0E+00 0,00

I

0,15

X

9

9

9

4 0,15

i|

~'~ Z

Po2 = 0.1)5 bar

o

E 5,0l::-04

~"

" d

I i

0,0E+O0 0,00

0,05

0,10 0,15 P~t=~ (bar)

Figure 2. Influence of reactants partial pressure on the dehydrogenation products formation rates over 3% Cs - NiMoO4 (symbols and legend as in Figure 1). Approaching from a different point of view, Bielanski and Haber analyzed the question of selectivity by considering the types of oxygen available [ 16]. They propose that the reactivity of oxygen depends on whether it is electrophilic or nucleophilic. Electrophilie oxygen species,

802 such as adsorbed O2 and O" are very active and attack hydrocarbon molecules at the regions of high electron densities. Nucleophilic oxygen species, such as lattice oxygen are less reactive and are suitable for partial oxidations. In a reaction with nucleophilie oxygen, activation of the hydrocarbon molecule is the rate-determining step. In a reaction with electrophilic oxygen, adsorption of oxygen to form the electrophilie species is the rate-limiting step. Therefore, a selective oxidation catalyst should be able to adsorb and activate a hydrocarbon molecule for nucleophilic attack by oxygen. It should not generate electrophilic oxygen species. For the NiO-MoO3 system it was found that lattice oxygen is responsible for ODH, while other products of selective and complete oxidation are formed by other species of surface oxygen [ 17]. Moreover, Mazzoeehia et al. [13] found by electrical conductivity measurements the presence of anionic vacancies at the surface lattice of nickel molybdates. They suggest that in propane ODH, propene would be produced via formation of such vacancies resulting from reaction of 0 2- surface anions. The role of gaseous oxygen would be to replenish rapidly the lattice oxygen consumed, reoxidizing the solid. Finally, it must be remarked that gaseous oxygen may have another role: the formation of carbon oxides by attacking the weakened C-C bonds of the adsorbed butane molecules. The kinetic studies found in literature (of. Introduction) suggest that this reaction proceeds via a Mars and van Krevelen mechanism [ 18] and then the following steps are involved: (i) reaction of the hydrocarbon with the oxide to give products and a partially reduced catalyst; (ii) reoxidation of the catalyst by gaseous oxygen to restore the catalyst in its original state. In this mechanism the agents responsible for oxidation are the oxygen ions of the oxide lattice. Therefore, the catalytic behavior of oxide catalysts towards hydrocarbon oxidation should depend on the strength of the metal-oxygen bond of the catalyst [19]. In a previous paper it was found that the catalytic activity for butane ODH over various Cs doped NiMoO4 samples (which decreases when increasing Cs contents) is related to the catalyst reducibility [12]. In fact, Cs promotion increases resistance to reduction, as evidenced by the increase in the temperature of onset of reduction with the Cs loading. This relationship between catalytic activity and reducibility of the catalysts (which gives an idea of the metal-oxygen bond strength) suggests therefore also the existence of a redox mechanism The involvement of lattice oxygen in the reaction is also supported by the fact that no oxygen desorption was observed by O2-TPD and by experiments without oxygen in the reactor feed [20]. In fact, even in the absence of gas phase oxygen, butane can be converted to C4's with a high selectivity, in spite of the low conversion values. The decrease of activity would be due to the decrease of the fraction of active oxidized surface. Also the apparent activation energies for butane consumption for both processes (with and without oxygen in the gas phase) are similar, being the differences due to the different products distribution [21 ]. Taking into account all the above mentioned results, a generalized Mars and van Krevelen model was considered for butane ODH over both catalysts. Consequently, the two steps mentioned above can be generally described by the following redox scheme: butane + OC RC + 02

r, products + RC > OC

where OC and RC are oxidized and reduced active sites, respectively. The reduction and reoxidation rates of catalyst surface earl therefore be expressed as: rr= k ~ P ~ e o

(1)

803

ro = ko Po2"1 p ~ , . 2 Or

(2)

where k~ and ko are kinetic constants, Oo and O~ the fractions of oxidized and reduced sites, respectively, Pi the partial pressure of reactants and ni the partial orders. In steady-state, rr= ro/Ot

(3)

ct is the stoichiometric coefficient of oxygen, i.e., the number of 02 moles reacting per mole of butane. For butenes a = 0.5; for butadiene a = 1; and for the total of C4"s, a is a function of the products distribution. From eqs. (1), (2) and (3) 0o and the global rate can be obtained:

Oo = ko Po2 nl / ( ko P02 nl + o~ kr Pb,aane(n'n2))

(4)

r = ko kr Pbutanen POE"1 / ( ko Po2"1 + ot l~ Pb~.r ("-"2))

(5)

For estimation of the parameters of eq. (5) including the partial orders, the equation was fitted to the experimental rates of each dehydrogenation product by optimization of a multiple correlation coefficient W [20] using non-linear regression analysis. It is noteworthy that for both catalysts, the equation rates that best describe the formation of all dehydrogenation products are all of the same type with n = 1 and nl = n2 = 0. The influence of butane partial pressure in eq. (2) was investigated because a negative effect was previously found for 1butene ODH over bismuth molybdates due to a competition of oxygen and 1-butene over the reduced site [22]. Nevertheless, such order was found to be null in this case. An interesting feature is the fact that Po2 does not affect the oxidation rate of the reduced active sites, i.e., nl = 0. This may evidence that the diffusion of lattice oxygen controls the global process of the catalyst reoxidation. A similar behavior was found for 1-butene ODH over bismuth molybdate at lower temperatures [22]. Therefore, eqs. (4) and (5) can be simplified to 0o = ko / (ko + ct k~ P ~ t ~ )

(6)

r = ko kr Pbutane / ( ko + o~ kr Pbutane )

(7)

The other computed parameters (ko and IQ are presented in Figure 3 as Arrhenius plots. The fitting curves obtained with the computed parameters are included in Figures 1 and 2. The computed values of 1%and k, agree with the Arrhenius law: k i - ki'x exp (- Ea~ / (RT))

(8)

which is additional evidence in favor of eq. (7). From the Arrhenius plots (Figure 3) the preexponential factors k( and the activation energies Eai were computed (Table 1). For the unpromoted NiMoO4 the activation energies for reduction are always higher than Eao. The negative activation energies computed for butenes evidence complexities not taken into account. For Cs doped catalyst activation energies are larger for the reoxidation process, except for 2-butenes. In this case, the difference in the kinetic constants is so high that ko >> ct k~ P ~ and eq. (7) is simplified to a linear variation of the reaction rate with the butane partial

804 1- b u t e n e

e-

1-butene

-7

-9

-11 1,18E-03

-5

~ 1,22E-03 1,26E-03 1 /T (K1)

-11 1,18E-03

I

1,30E-03

trans -2-butene

-7

10

_-

-5

-5

B

-2-butene

tra n s

_-

t ~ 1,22E-03 1,26E-03 1 / T (K "1)

0

J i 1,30E-03

c/s -2-butene

-10 1,18E-03

A

t

I

1,22E-03 1,26E-03 1 / T (K"1)

cis

A

B

-2-butene

t

-7

~

~

I

I

1,22E-03

1,26E-03

1,30E-03

B

~

1 / T (K "~) -3 -5

C4"s

A

I

~

1,22E-03 1,26E-03 1 / T (K "~)

-4

i

1,30E-03

B

C4"s

~ " " ~ 6 _ ~ . . . . . ~ c

-9 1,18E-03

~

-11

1,18E-03

1,30E-03

butadlene

-9 -9

i

1,22E-03 1,26E-03 1 / T (K "1)

-5

.to _r -7

1,18E-03

1,30E-03

0

-10 1,18E-03

I

1,30E-03

butadlene

J 1,22E-03 1,26E-03 1 / T (K 1)

10

_c

-11 1,18E-03

i 1,30E-03

~

,-

-11 1,18E-03

t p 1,22E-03 1,26E-03 1 I T (K "~)

~ I 1,22E-03 1,26E-03 1 I T (K~)

I 1,30E-03

-6

-8

1,18E-03

I

t

1,22E-03 1,26E-03 1 / T (K"I)

~

1,30E-03

Figure 3. Arrhenius plots of the kinetic constants ko (B) and kr (0) for dehydrogenation products over NiMoO4 (A) and 3% Cs - NiMoO4 (B). pressure (r = k~ P ~ ) what is in very good agreement with the experimental data, specially at higher temperatures. In what concerns the other products with the 3% Cs promoted catalyst,

805 Eao > Ear what means that in eq. (7) ko increases stronger with temperature than ~t kr Pb~t~c. In this way, at a temperature of 560 ~ it is necessary a higher butane partial pressure than at 500 ~ in order to run away from the linear variation of r vs. Pbm,o (of. Figure 2). Table 1 Pre-exponential factors and activation energies for formation of dehydrogenation products Catalyst Const? 1-butene t-2-butene c-2-butene butadiene C4"s ko' 2.5• -5 6.1• -7 3.6x10 ~ 6.0 9.6• .3 NiMoO4 kr' 8.8• 5 7.9• 2 1.0• 4 2.4x10 z 3.0• 3 Eao -4.1• 3 -3.0• 4 -1.7• 4 6.7• 4 2.2• 4 Ear 1.3• 5 9.0• 4 1.1• 5 7.8• 4 8.7• 4 ko" 1.8• 1.1• 1.2• 6.0x10 s 5.4x107 3% Cs - NiMoO4 kr' 1.4 8.9 12.8 33.3 22.8 Eao 1.4• 3.4x104 3.4• 1.9x105 1.6xl05 Ear 4.9x104 6.7x104 6.8x104 7.0x104 6.0• a ko" in mol / (h.m2); k~" in mol / (h.m2.bar) and activation energies in J/mol. 5. CONCLUSIONS The negative activation energies obtained with fittings for butenes with the unpromoted ctNiMoO4 catalyst (Table 1) indicate that their formation involves more complex processes than the admitted ones. In fact, in a previous study with the 3% Cs promoted sample [20], when feeding the reactor with 1-butene, selectivities for butadiene around 40% were found while the selectivities of formation of 2-butenes do not exceed 7%. When cis-2-butene is fed, the butadiene selectivity reaches 60% and the selectivity to isomers is around 8-9%. These results point out that isomerization reactions can be neglected and that the subsequent dehydrogenation of each butene to butadiene proceeds in a considerable extension. For undoped catalyst due to its higher acidity, both subsequent dehydrogenation to butadiene and isomerization reactions play a more significant role. However, it must also be taken into account the possible direct formation of butadiene from butane. On the other hand, it is well known that in this ODH reaction n-butane first reacts by cleavage of a secondary C-H bond to form an adsorbed alkyl species. This step is generally accepted to be the rate-determining step. After a second H-abstraction, olefinic intermediates are formed and desorbed as olefins or undergo another dehydrogenation to form butadiene [1,6,7]. Therefore, it is convenient to consider the global process and to use the rate of C4"s formation as a comparative term. It was found that for the dehydrogenation products, and in contrast to expectations, cesium doping decreases the activation energy for the reduction step increasing the corresponding to oxidation (Table 1). Nevertheless, the Cs effect is more pronounced in the pre-exponential factors. In fact the doped catalyst exhibits smaller k~' values and larger ko" than the undoped one. The increase in ko" is so strong that higher kinetic constants for reoxidation are obtained with 3% Cs - NiMoO4 and higher values of kr with NiMoO4 (Figure 3). The increase of the nickel molybdate resistance to reduction after Cs addition explains the smaller k~ values for the doped catalyst. On the other hand, Cs promoted catalysts are much more conducting than unpromoted NiMoO4 due to the contribution of a surface ionic conductivity by mobile Cs + and 02- ions to the overall conductivity [5]. Moreover, solids with high oxygen-

806 ion conductivity have a high capability of transforming any surface oxygen species into lattice oxygen [5]. Consequently, Cs doping favors the reoxidation of the solid because it favors both the oxygen incorporation into the lattice and the diffusion of those species through the solid. Finally it must be pointed out that by using eq. (6), higher 0o values for the promoted catalyst are obtained what means that under any reaction conditions Cs maintains the catalyst in a higher oxidation state allowing the use of higher butane concentrations. REFERENCES

1. H.H. Kung, in D.D. Eley, H. Pines and W.O. Haag (eds.), Advances in Catalysis, Vol. 40, Academic Press, New York, 1994, pp. 1-38. 2. C. Mazzocchia, R. Del Rosso and P. Centola, An. Quim., 79 (1983) 108. 3. R.M. Martin-Aranda, M.F. Portela, L.M. Madeira, F. Freire and M.M. Oliveira, Appl. Catal. A, 127 (1995) 201. 4. F.J Maldonado-H6dar, L.M. Madeira, M.F. Portela, R.M. Martin-Aranda and F. Freire, J. Mol. Catal. A, 111 (1996) 313. 5. L.M. Madeira, J.M. Herrmmm, F.G. Freire, M.F. Portela and F.J. Maldonado, Appl. Catal. A, in press. 6. E.A. Mamedov and V.C. Corberfin, Appl. Catal. A, 127 (1995) 1. 7. M.A. Pepera, J.L. Callahan, M.J. Desmond, E.C. Milberger, P.R. Blum and N.J. Bremer, J. Am. Chem. Sot., 107 (1985) 4883. 8. A. Escardino, C. Sol/Land F. Ruiz, An. Quim., 69 (1973) 385. 9. H.H. Kung, in "Transition Metal Oxides: Surface Chemistry and Catalysis", Studies in Surface Science and Catalysis, Vol. 45, Elsevier, Amsterdam, 1989. 10. G. Centi, G. Fornasari and F. Trifir6, J. Catal., 89 (1984) 44. 11. A. Kaddouri, R. Anouchinsky, C. Mazzocchia, M. Madeira and M.F. Portela, Catal. Today, submitted for publication. 12. L.M. Madeira, M.F. Portela, C. Mazzocchia, A. Kaddouri and R. Anouchinsky, Catal. Today, submitted for publication. 13. C. Mazzocchia, C. Aboumrad, C. Diagne, E. Tempesti, J.M. Herrmann and G. Thomas, Catal. Lett., 10 (1991) 181. 14. F.J.M. H6dar, L.M.P. Madeira and M.F. Portela, J. Catal., 164 (1996) 399. 15. M. Iwamoto, Y. Yoda, N. Yamaoe and T. Seiyama, J. Phys. Chem., 82 (1978) 2564. 16. A. Bielanski and J. Haber, in "Oxygen in Catalysis", Marcel Dekker, Inc., New York, 1991. 17. C. Mazzocchia, R. Del Rosso and P. Centola, Proc. 5th Ibero-American Symp. Catal., 1976, Rev. Port. Quim., 19(1-4) (1977) 61. 18. P. Mars and D.W. van Krevelen, Chem. Eng. Sci. (Spec. Suppl.), 3 (1954) 41. 19. D.J. Hucknall, in "Selective Oxidation of Hydrocarbons", Academic Press, London, 1974. 20. L.M. Madeira, F.J.M. H6dar, M.F. Portela, F. Freire, R.M.M. Aranda and M. Oliveira, Appl. Catal. A, 135 (1996) 137. 21. F.J.M. H6dar, L.M. Madeira, M.F. Portela and R.M.M. Aranda, Proc. 15th IberoAmerican Symp. Catal., (1996) 251. 22. M.F. Portela, M.M. Oliveira, M.J. Pires, F.M.S. Lemos and L. Ferreira, Proc. 8th International Congress on Catalalysis, (1984) 533.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

807

A COMPARISON OF IRON MOLYBDATE CATALYSTS FOR METHANOL OXIDATION PREPARED BY COPREClPTATION AND NEW SOL-GEL METHOD A. P.

Vieira Soaresa, M. Farinha Portela~9and A.Kiennemannb

aC~ECAT- Gn~ de ~ de Cmtdise~ UniversidadeT~-'nicade Lidx~ ~ Rm4.so3Pais, 1096Lidx~ Cxxt~ Portugal,Fax:351-1-847"/695

SulxriorT6cmco,Avenida

%CPM - D6partement Catalyse, LERCSI, Universite Louis Pasteur, 1 Rue Blaise Pascal, BP296, F 67008 Strasbourg Cedex, France Iron molybdates, with atomic ratio Mo/Fe=3, were prepared by coprecipitation and sol-gel techniques in acid medium. Characterisation results show that sol-gel catalysts have much higher surface areas than coprecipitated catalysts. Study of catalytic activity shows that Fedefective catalysts, prepared by sol-gel technique, perform better than an industrial catalyst in the same conditions. In fact these catalysts achieve higher performances at 50K lower than industrial catalysts. 1. I N T R O D U C T I O N Formaldehyde is nowadays one of the major produced chemicals due to its uses in many fields of chemical industry [ 1]. The commercial production started in 1890 using metallic copper catalysts. In 1910 copper catalysts were replaced by silver catalysts with higher yields [2]. Although the first report of the excellent catalytic behavior of iron molybdates in selective oxidation of methanol to formaldehyde is of 1931, the related industrial process based on them only went into operation in 1940-50 [1]. A recent report [3] shows that iron molybdates and silver catalysts are nowadays equally used as industrial catalysts for formaldehyde production. Many metal molybdates catalyse the reaction under consideration [4] and the active sites are widely believed to be associated with surface Mo atoms in octahedral coordination [5]. Octahedral coordination of Mo is only achieved in Fe-defective iron molybdates what is in accordance with the fact that maximum activity is obtained with catalysts with Mo/Fe atomic ratio greater than 1.5 [6]. The presence of two terminal oxygens double bonded to Mo in such coordination allows the methanol reacting molecules to be bonded simultaneously by two points. The activation of the hydrogen of the hydroxyl group produces methoxy species that are intermediates in the formaldehyde formation. The role of Fe in iron molybdates catalysts would be to act in the transfer of 02 and H20 between surface and gas phase [6] and to hinder the reduction of Mo +6 [7]. Several studies of the catalytic behavior of iron molybdates as a function of their physical and chemical properties show that surface acidity, related with M0 +6 ions, is a necessary condition for the catalyst effectiveness for formaldehyde formation [8,9]. The oxidizing function is also necessary, though the activity is not necessarily determined by this function. In a recently study Sun-Kuo et al [6] demonstrated that catalytic behavior of iron

9To whom correspondence should be addressed

808 molybdates is closely related to atomic ratio Mo/Fe of the catalyst. They found the maximum activity for a catalyst with an atomic ratio Mo/Fe=l.7. This result agrees with the work of Kolovertnov et al 30 years ago [10]. Some researchers claim that the active phase is normal iron molybdate but recognise that an excess of Mo is required to compensate its loss during reaction. Industrial catalysts are always Fe-defective molybdates due to this reason. Up to now several methods have been used to prepare iron molybdates, the most part of them based on coprecipitation techniques. Previous studies [11 ] have evidenced that the catalytic behavior of Mo-Fe oxides depends on many variables of the coprecipitation procedure: starting compounds, concentration of parent solutions, pH and temperature of coprecipitation step, order of addition of parent solutions, ripening etc. In a typical preparation procedure iron molybdate is coprecipitated from solutions of ferric chloride or ferric nitrate and ammonium molybdate [8]. The control of all the above mentioned procedure variables, strongly difficult the preparation of these catalysts and deviations from the preparation recipe can have very adverse effects on the performances of the catalyst from the standpoint of activity, selectivity and stability. Recent sol-gel methods have been recognized as promising procedures to prepare catalysts [12-14]. The sol-gel methods allows a unique way of catalyst design, because they represent an ab initio synthesis of the final solid from well defined molecular compounds [ 13]. By suitable choice of reagents, reaction and drying conditions, such technique allows to predefine pore structure, porosity, composition, surface polarity and crystallinity or amorphicity of metal oxides [12]. In principle, any metal that forms stable oxides can be forced to copolymerise with other metals in sol-gel procedures to provide mixed metal oxides [ 13]. To investigate of influence of preparation methods on catalytic properties of iron molybdates catalysts with controlled excess of MoO3 were prepared by coprecipitation and solgel techniques. Their properties and performances were compared with an industrial catalyst. 2. EXPERIMENTAL

2.1. Catalyst Preparation Iron molybdates with atomic ratio Mo/Fe = 3 were prepared by coprecipitation and solgel techniques. Coprecipitated catalysts were obtained from starting aqueous solution of Fe(NO3)3.9H20 and (NH4)eMo7024.Iron nitrate solution was slowly added to the solution of Mo previously acidified (pH ~2) with HNO3. After the addition of Fe solution the precipitates were ripened in contact with mother liquor at 100~ for 3 hours under vigorous stirring. Afterwards the precipitates were filtrated, dried at 120 ~ overnight and finally calcined. Sol-gel catalysts were prepared in acid medium using appropriate molybdenum precursor and Fe(NO3)3.9H20 or FeCI3. In a typical procedure, first the precursor solutions of Mo and Fe in acid medium were prepared and then ferric solution was added to Mo solution. No precipitate was observed during the addition. The resulting solution was evaporated until dryness. The formed film was removed by adding liquid N2, then was crushed in a glass mortar and dried at 120~ overnight. Calcination, for both techniques, was performed at 648 K during 10 hours. For coprecipitated catalysts calcination was always under air flow, whereas sol-gel catalysts were calcined with or without air flow.

809

2.2. Catalyst Characterisation The BET surface areas were measured using a Perkin-Elmer Shell 212C sorptometer instrument based on the N2 physisorption capacity at 77K. Bulk elemental composition were determined by atomic absorption. X-Ray diffraction patterns were obtained with a D5000 Simens diffi'actometer equipped with a primary beam quartz monochromator (Co Kczl=l.78897A~) at 40kV and 25 mA. The morphology, chemical analysis and homogeneity of the prepared catalysts were examined with a scanning electron microscope (SEM) JOEL JSM840 equipped with a Delta Kevex energy dispersive X-ray analyzer. FT-IR spectra were recorded with a Perkin-Elmer 1600 spectrometer with a range of 4000 - 400 cm-1, and a resolution of 2 cm-1. X-ray photoelectron spectroscopy (XPS) was used for chemical analysis and investigation of reduction state at the catalysts surface. The analyses were performed on a VG ESCA 3 apparatus. The kinetic energy of the emitted photoelectron is given by: E~=1486.6-EBcor, were 1486.6 is the energy of the incident radiation (AI Kcz ray] and Emo~ the corrected binding energy of the electron. The binding energy were calibrated with respect to the signal for adventitious carbon (biding energy: 284.8 eV). Iron phases bulk composition was studied by MOssbauer spectroscopy at room temperature operating in the constant acceleration mode and with a radiation source of 57Co in a Rh matrix. 2.3. Catalytic Tests Methanol oxidation was carried out in a conventional flow apparatus at atmospheric pressure. The feed mixtures were prepared by injecting the liquid methanol into air flow with a Gilson 302 pump. The catalyst was diluted with inert carborundum (1:3 volume ratio) to avoid adverse thermal effects, and placed in a tubular pyrex reactor with a coaxially centred thermowell with thermocouple. The reactor outlet was kept at 403 K, to prevent condensation of liquid products and formaldehyde polymerization, and it was connected with multicolumn Shimadzu GC-SA gas chromatograph with thermal conductivity detector. The column system used (1.5m of Poropak N+l.5m of Poropak T+0.9m of Poropak R) could separate CO2, formaldehyde, dimethylether, water, methylformate, dimethoxymethane and formic acid. The last product was never detected. 3. RESULTS AND DISCUSSION

3.1. Catalyst characterisation BET results in figure 1 show that sol-gel technique yields catalysts with surface areas that are approximately twice of coprecipitated ones. This may be due to formation of a second amorphous phase of MoO3 in sol-gel catalysts [6] instead occupation of lattice interstices, by Mo excess, which occurs in coprecipitated catalysts [ 16]. Industrial catalysts present a lower surface area than coprecipitated catalyst with the same atomic ratio Mo/Fe (=3) which can be attributed to the severe calcination step. The scanning electron micrographs of the catalysts (figure 2) show that catalysts prepared by both methods have the same sponge-like morphology. This result agrees with the fact that Mo excess retards the crystallization of Fe2(MoO4)3 [ 17]. Morphological aspect and surface areas are in good agreement. However our results are in disagreement with recent results of Sun-Kuo et al [6]. These authors have found that highest surfaces areas corresponded to samples formed by associated ordered lamellae. After long activity tests, including catalytic tests with water in reactor feed, catalysts kept their morphological aspect.

810

L

i I

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i I

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: ::- -:..:. ~_:::-::, ~ :.. :::.--~i :~ :- i~ --:-_-:--.~ i i i

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_:- :.~: ~.~.- - :_ .-i-~-._-:-.-.~-:~

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0

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I

2

4

I

6

8

10

12

14

16

18

surface area (m~/g)

Figure 1 - Surface area of catalysts prepared by coprecipitation and sol-gel techniques (from no halogenated precursors), and industrial catalyst. a)

b)

Figure 2 - Scanning electron micrographs of flesh catalysts a)Mo/Fe=3 coprecipitated, b)Mo/Fe=3 sol-gel. Bulk elemental analysis of some catalysts was performed by atomic absorption. The results presented in table 1 show that coprecipitation method yields, within the experimental errors, catalysts with the Mo/Fe atomic ratio of preparation. Catalysts prepared by sol-gel method showed lower Mo/Fe atomic ratios than the expected value. Furthermore calcination of these catalysts under air flow enhances the loss of Mo. Table 1 Bulk elemental composition by atomic absorption Catalyst

M o / F e (atomic ratio)

Coprecipitated

2,9

Sol-Gel (calcined without air flow)

2,4

Sol-Gel (calcinated with air flow)

1,9

Industrial

3,1

811 In figure 3 some representative X-ray patterns of the fresh catalysts are presented. All catalysts look crystalline but coprecipitated catalyst seems to be slightly better crystallized than sol-gel one. They present the diffraction lines corresponding to normal iron molybdate and MoOs. Some diffraction lines of these two phases are superimposed. Furthermore catalysts with a large excess of Mo present x-ray lines that cannot be assigned neither to normal iron molybdate nor to MoOs. According to Abaulina et al [16] those lines (6.9, 3.4, 2.3) correspond to another phase, similar to natural ferrimolybdate, formed by MoOs dissolved in iron molybdate. Catalysts prepared by sol-gel method the using halogenated precursors have the same x-ray patterns. Calcination under air flow has no effect on the phase distribution yielded by this technique. Phase composition and degree of the reduction of catalysts were also examined by MOssbauer spectroscopy at room temperature. In table 2 parameters for flesh catalysts are presented. The results show that Fe-defective coprecipitated catalyst has only one type of iron(III) molybdate which correspond to normal iron molybdate. Fresh sol-gel catalyst with Mo excess have a small amount (3.8%) of reduced phase FeMoO4 which is not detected by XRD.

0

I 15

20

! 45

! 60

Figure 3 - X-ray diffraction patterns of fresh catalysts MOssbauer analysis after long catalytic runs evidences that Mo/Fe=3 catalyst prepared by coprecipitation undergo reduction; after 72h of reaction this catalyst had 2% of FeMoO4. Solgel catalyst lose Mo during long activity runs and show MOssbauer characteristic parameters of" Fe203 phase (5%). The presence of FeMoO4 supports the deactivation mechanism proposed by Pemicone et al [19]: Fe2(MoO4)3+CH3OH ~ 2FeMoO4+MoO3+HCHO+H20 All IR spectra in figure 4 show a strong absorption band centered at 830 cml what is associated with stretching of Mo=O in tetrahedral environment [6,16,19]. The shoulder located at 780 cm-1 can be ascribed to stretching of Mo-O bond of heteropolyanions of Mo with octahedral coordination [ 19]. The weak band at 960cm "l is assigned to Fe-O-Mo vibrations in the ferric molybdate phase [19]. The broad band at 610 cm-1 displayed by Mo/Fe=3 coprecipitated, sol-gel and industrial catalyst is assigned to Mo=O in octahedral environment of Fe-defective iron molybdates [ 16]. The intensities of this band and of the band at 995cm l

812 are proportional to the Mo excess of the catalysts and they tend to disappear with time on stream: the industrial catalyst had almost the same composition than Mo/Fe=3 coprecipitated catalysts, but the intensities of those two bands were smaller. This strange result is possibly due to the fact that the industrial catalyst is 7 years old. Table 2 Mrssbauer parameters of fresh catalysts (~i-isomeric displacement, W-band width, Aquadripolar separation). CATALYST Fe species ~i W A I/Io (%) Phase mln.s -1

mln.s -1

m m . s -~

Industrial

Fe 3+

0,41

0,32

0,21

1O0

Fe2(Mo04)3

Mo/Fe=3

Fe 3+ Fe 2§

0,44 0,89

0,26 0,28

O,19 2,34

99 <1

Fe2(Mo04)3

Fe 3+ Fe 2+ Fe 2§

0,41 1,06 1,19

0,30 0,24 0,26

O,19 2,33 0,93

96,2 2,0 1,8

Fe2(Mo04)3 FeMoO4

Coprecipitated

Mo/Fe=3 Sol-Gel

1

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Figure 4 - IR spectra of fresh catalysts a) sol-gel from no halogenated precursors, b) coprecipitated, c)industrial, d)MoO3. Chemical analysis and reduction state at the surface of fresh and used catalysts were performed by XPS (table 2). Only Mo(VI) was detected, with a binding energy of 232.1-232.4 eV. This value is agrees with the value obtained by Petrini et al [20] for similar catalysts: 232.5 eV. Fe energy band (2P3/2) were complex and can be ascribed to Fe 3+ species in iron molybdate, with a binding energy of 711.9 eV, and to Fe 2+ species of FeMoO4 with a smaller binding energy: 709.9 eV. These results also agree with the results presented by Petrini. Mo/Fe atomic ratios and degree of reduction (Fe:+/Fe3+), at surface, are summarized in table 3. Prepared catalysts present smaller atomic ratio than the expected value:3. During reaction, with or without water in reactor feed, there is migration of Mo from the bulk to the catalyst surface; Mo/Fe ratios for used catalysts were always higher than for fresh catalysts. Water in reactor feed seems to accelerate this process. Surface reduction is stronger for sol-gel than for

813 coprecipitated catalysts. The surface reduction increases during reaction and water also accelerates this process.

3.2. Catalytic Tests For all tested catalysts the products of methanol oxidation were formaldehyde, dimethylether, methylformate, dimetoxymethane and CO. At high temperatures (T>300~ and high level of conversion (>90%) small amounts of CO2 were also observed. Table 3 Atomic ratios and degree of reduction at surface for fresh and used catalysts by XPS (preparation atomic ratio Mo/Fe=3). CATALYST

Coprecipitated

fresh after 72 h of reaction after 24 h of reaction with water

Mo/Fe 2,35 2,60 2,43

Fe2+/Fe3+ 0,19 0,41 0,46

Sol-gel

fresh after 72 h of reaction after 24 h of reaction with water

2,17 2,18 2,43

0,60 0,58 0,72

Catalytic runs for testing the effect of preparation method on catalytic behavior were performed at 573 K with 4% of methanol (in air) in reactor feed. Results (figure 5) show that among tested catalysts, including an industrial catalyst, the more active and selective was the sol-gel catalyst. Comparing the behavior of this catalyst with the industrial catalyst, at several temperatures (figure 6) it is noteworthy that the sol-gel catalyst performs better than industrial catalyst, even at temperatures 50K lower. The calculation of the specific activities, of formation of formaldehyde and CO, from yield versus contact time data provided the results presented in table 4. It is seen that the activity per square meter of surface are for both products, exhibited by sol-gel catalyst is lower than the specific activity of coprecipitated catalyst. Such differences of activity are attributable to the lower Mo contents of sol-gel catalysts (table 1). 100

75

[]Conversion [] Sel CH3OCH3 9 Sel HCOOCH3 [] Sel CO

% 50

25 DII~IUgI~| Igllgll~ll~ll~l 9~ ! ~ !

.

.

.

.

.

.

.

.

i

Mo/Fe=3 coprec.

.

.

.

.

.

.

.

.

.

.

.

.

.

i

Mo/Fe=3 sol-gel no halogenated precursors

i

Industrial

Figure 5- Conversion and selectivities of catalysts prepared by coprecipitation and sol-gel (without halogenated precursors).

814 Results in figure 7 evidence that calcination of sol-gel catalysts under air flow leads to a strong loss of activity, attributable to loss of volatile species of Mo during calcination step. The use of halogenated precursors in preparation of sol-gel catalysts occasions loss of activity, eventually due to the formation of halogenated species on catalysts surface that are inactive or less active. Unfortunately it was not possible to identify such hypothetical species. Table 4 Specific activities (T=598 K) of formation of HCHO and CO reaction mixture

specific activity Otmol mons-tm~ HCHO CO 2,32 0,46 2,48 0,54 1,65 0,36 1,14 0,29 0,74 0,24

0a'a)

Coprecipitated

Sol-Gel

MeOH 0,10 0,18 0,54 0,93 1,30

02 17,74 17,17 18,00 18,25 18,44

H20 3,34 3,73 4,23 4,69 5,02

HCHO 2,33 2,55 2,91 3,27 3,26

0,55 0,60 0,73 0,86

17,70 17,69 17,84 17,94

4,89 4,89 4,57 4,38

2,13 2,14 2,23 2,17

100

0,11 0,49 0,77 0,98

0,25 0,33 0,39 0,43

m

...i -.

_

_

_

~HG'HO

%50

C A - 1 3 O C H 3

~

HO:X)GH3

~CO

m

Industrial T--548K

Industrial T=573K

Imlumial T=623K

Mo/Fc=3

Mo/Fc=3

T---54gK

T=573K

sol-gcl

sol-gel

Figure 6- Conversion and selectivities of industrial and sol-gel catalyst at several temperatures. Runs for testing stability (figure 8) showed that prepared catalysts were stable in used operating conditions. Furthermore, water do not inhibit the reaction and have a benefit effect on the selectivity. Comparing the results of stability and the degree of surface reduction (Fe2+/Fe3+ in table 2) we conclude that there is no correlation between the two set of results and the apparent high degree of surface reduction is probably due to the high vacuum conditions during XPS measurements. 4. CONCLUSIONS Sol-gel techniques provide interesting routes to prepare iron molybdates catalysts for selective oxidation of methanol. Catalysts prepared by this method have higher surface areas than catalysts prepared by coprecipitation techniques what allows to operate at lower temperatures with the advantage of limiting the consecutive oxidation of formaldehyde to CO.

815 In this type of catalysts Mo excess form an amorphous separate phase that provides large contribution to the total surface area, thus increasing the number of methanol adsorption sites. The precursors used in sol-gel methods seem to influence the catalytic properties of the catalyst. In particular halogenated precursors have a beneficial effect on the catalyst selectivity. 1 0 0

_

_

_

_

_

Q

~

9

o/e 50

-

Sel

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no halo~

halog,

precursors

precursors

preempts

preem~ors

halo~

calcination without an"

calcination in air

calcination without air

calcination in air

Figure 7-Effect of precursors on catalytic behavior of catalysts prepared by sol-gel(T=573 K). coprecipitated a) 100

.

.

.

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El -conversion O-sel. HCHO A-sel.CO. Figure 8-Stability tests without (a) and with (b) water in reactor feed.

I

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10 15 Time (It)

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20

_

816 REFERENCES

1. Reuss, G., W. Disteldorf, O. Grundler, A. Hilt, "Formaldehyde", in Ullmann~Encyclopedia of Industrial Chemistry, VoI.A11, p.619, VCH Publishers, 54 Ed. (1992). 2. Satterfield, C. N., "Catalytic Oxidation Methanol to Formaldehyde", in Heterogeneous Catalysis m Pratice, McGraw-Hill, New York (1980). 3. ECN Process Review, p.30, April 1994. 4. Golodets, G. I., J. R. H. Ross, "Heterogenous Catalytic Reactions Involving Molecular Oxygen", m Studies m Surface Science and Catalysis, Vol. 15, Elsevier Science Publishers B. V., Netherlands (1983). 5. Trifir6, F. S. Notarbartolo and I. Pasquon, J. Catal., 22, 324 (1971). 6. Sun-Kou, M. R., S.Mendioroz, J. L. G. Fierro, J. M. Palacios, A. Guerrero-Ruiz, J. Mater. Sci., 30, 496 (1995). 7. Novakova, J. and P. Jiru, J. Catal., 27, 155 (1972). 8. Pernicone, N., J. Less-Comm. Met., 36, 289 (1974). 9. Ai, M., J. Catal., 54, 426 (1978) 10.Kolovertnov, G. D., G. K. Boreskov, V. A. Dzisko, B. I. Popov, D. V. Tarasova and G. C. Belugina, Kinet. Katal., 6, 6, 950 (1965). 11 .Wilson, J. H., Ph.D. Thesis, University of Wisconsin-Madison (1986). 12.Schneider, M and Baiker A., Catal.Rev.-Sci. Eng., 37(4), 515 (1995) 13.Wilhelm F. Maier, F. M. Bohnen, J. Heilm~n, S. Klein, Hee-Chanco, M. F. Mark, S. Thorimbert, I.-C. Tilgner, M. Wiedorn, in Applications of Organometallic Chemistry in the Preparation and Processing of Advanced Materials, Ed. J. F. Harrod and R. M. Laine, Kluwer Academic Publishers, Netherlands (1995). 14.Mizukami, F., S. Niwa, M. Toba, T. Tsuchiya, S. Shimidzu, S. Imai and J. Imamura, "Preparation and Propreties of the Catalysts by a Chemical Mixing Procedure" in Preparation of Catalysts IV, Ed. B. Delmon, P. Grange, P.A.Jacobs and G. Poncelet, Elsevier Science Publishers B.V., Amesterdam (1987). 15.Soares, A.P.V., Ph.D. Thesis, Technical University of Lisbon, Lisbon (1996) 16.Abaulina,L. I., G. N. Kustova, R. F. Klevtsova, B. I. Popov, V. N. Bibin, V. AMelekhina, V. N. Kolomiichuk and G. K. Boreskov, Kinet.Katal., 17, 5, 1126 (1976). 17.Boreskov, G. K., G. D. Kolovertnov, L. M. Kefeli, L. M. Plyasova, L. G. Karakchiev, V. N. Mastikhin, V. I. Popov, V. A. Dzis'Ko and D. V. Tarasova, Kinet. Katal., 17, 1, 125 (1965). 18.Pernicone N., Catal. Today, 11, 85 (1991) 19.Anagha, A. B., S. Ayyappan and A. V. Ramaswamy, J.Chem.Tech.Biotechnol., 59, 395 (1994). 20.Petrini, G., F.Garbassi, M. Petrera and N. Pernicone, "Study of Iron(H) Molybdate as Precursor of Catalysts for Methanol Oxidation to Formaldehyde", in Chemistry and Uses of Molybdenum, Ed. H. F. Barry and P. C. Mitchell, p.437, Climax Molybdenum Company, Ann Arbor, Michigan, USA (1982).

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

817

Oxidation Catalysts Prepared by Mechanically and Thermally Induced Spreading of SbzO3 and V205 on TiO2

U. A. Schubert a, J. Spenglera, R. K. Grasselli a'b, B. Pillep c, P. Behrens c and H. KnOzingera

alnstitut for Phys. Chemie, Ludwig-Maximilians-Universitat, D-80333 Mimchen, Germany bDepartment of Chem. Engineering, University of Delaware, Newark, DE 19716-3116, USA ~ far Anorg. Chemie, Ludwig-Maximilians-Universitat, D-80333 MOnchen, Germany

Vanadium and antimony oxides are essential parts of some industrial catalysts for the selective oxidation of substituted aromatics to the corresponding anhydrides [1] and the selective oxidation of paraffins to the corresponding unsaturated acids and nitriles [2]. These catalysts are generally prepared by impregnation or coprecipitation methods. It was the objective of this study to investigate an alternate method of catalyst preparation, a method based on a solventless ball-milling technique, aimed at adequately dispersing the active catalyst ingredients on a given support material to yield acceptably effective catalysts. Specifically, the spreading and the dispersion of V-oxide, Sb-oxide, and V-Sboxides on Ti02 supports were investigated by means of the ball-milling technique, and the so prepared materials compared to conventionally prepared materials. The tribo-ehemieal process of the former method was followed by spectroscopic techniques including XPS, XANES, and TPR, which revealed that active phase dispersions comparable to those obtained by conventional preparation techniques can readily be achieved by the milling method. It was further observed that the addition of small amounts of water during the milling greatly enhances the rate of the dispersion process. Catalytic tests of a V-oxide-on-TiO2 composition reveal that the selective oxidation of o-xylene to phthalic anhydride (PA) proceeds at comparable levels for compositions prepared by either the milling or the impregnation method. Therefore, it is concluded that the solventless ball-milling technique is also an effective alternative method for the preparation of selective oxidation catalysts.

1. I N T R O D U C T I O N Partial oxidation of hydrocarbons employing mixed metal oxides as catalysts comprises an economically important class of reactions for the upgrading of base feed stocks [3]. An illustrative example of it is the partial oxidation of o-xylene and/or naphthalene to phthalic anhydride (PA) with a world production of 3.2 million metric tons per year, industrially carried out in shell and tube reactors using air as the oxidizing agent [4].

818 Catalysts employed in the oxidative production of PA are V20~-based compositions which are generally of the monolayer type and are supported on titania (anatase). With such catalysts, selectivities in excess of 80 mole % PA are achieved at essentially complete conversion. The utilization of Sb-V-oxide-based catalysts supported on anatase improves the PA selectivity. However, little is known about the intrinsic chemical or electronic effects of Sb203 in such catalytic systems, as well as the chemical and physical characterization of the supported Sb-oxide or supported Sb-V-mixed metal oxide [ 1]. Industrial catalysts are prepared exclusively by impregnation of the support material with aqueous solutions of the active phase materials, followed by subsequent drying and calcination. In the work described here, an alternative mode of catalyst preparation was chosen, based on a solventless, mechano-chemical method of incorporating the active catalytic components on the support carrier via ball-milling. Mechano-chemical treatment of solid materials leads not only to an alteration of its morphology and texture, e.g. formation of surface solid solutions [5], but also in the case of catalytic compositions to an alteration of their catalytic behavior [6]. For example it was demonstrated in the case of a V-oxide/TiO2 - based PA catalyst, that milling of the catalyst components led to an enhancement of PA yields, concomitant with observable spectroscopic and microscopic changes of the catalyst surface [7]. In our study reported here, the tribochemically induced dispersions of Sb-oxide and Voxide on TiO2 were investigated by means of spectroscopic methods for the binary systems Sb-oxide/TiO2 and V-oxide/TiO2, and the ternary system Sb-V-oxide/TiOv One of our research aims was to prepare a catalyst of the ternary system Sb203/V2Os/TiO2 by ball-milling, followed by calcination, having comparable dispersions of the single oxide components as those obtained by conventional methods employing solution impregnation. Another objective in our current study was to determine what influence the addition of water to the materials to be milled might have on the dispersion of the f'mal product, as compared to the product obtained by dry milling only. The catalysts prepared by ball-milling the oxide components with the carrier were evaluated by a test reaction, the oxidation of o-xylene with air, in a specially designed microreactor, determining their activity and selectivity, and by comparing their so determined catalytic properties with those of industrially prepared catalysts.

2. E X P E R I M E N T A L

For catalyst preparation the mechanical treatment was done in a planetary mill with the dry oxidic compounds or with 10 wt. % 1-120. For that purpose the oxides were mixed carefully in a mortar and then milled (145 rpm) in a agate vessel (250 ml) with six agate balls. The time of milling was chosen to be between l h and 20h. In the case of heat treatment of the materials, 400~ for 5h was chosen for calcination. The impregnated samples were prepared from an aqueous Sb-III-acetate suspension followed by drying at 110~ and calcination at 400~ for 5h.

819 The theoretical monolayer loadings were estimated according the formula of Roozeboom [8]. In the case of 1 ML SbxO3/TiO2, 10 wt. % Sb203 was used and for 1 ML V2Os/TiO2, 6 wt. % V20-yrI'iO2 Was used. The XPS measurements were carried out at a modified VSW ESCA 100 with Mg Ka and AI Ka X-ray radiation. The analyzer was operated in the fixed analyzer transmission mode (FAT) at 22 eV pass energy. The resolution of the spectrometer is given by the linewidth of AU4fT/2. For a sputtered and annealed gold sample, 1.65 eV was obtained. The lines corresponding to Cls, Sb3ds/2, V2p3/2, Ti2p and Ols were analyzed. The pressure in the analyzer chamber during spectra collection was always below 5.0 * 10-8 mbar. To determine charging effects, all signals were referred to the C 1s line corresponding to graphitic carbon at 284.4 eV or to the Ti2p3/2 signal from TiO2 at 258.5 eV [9]. The TPR experiments were carried out with the following parameters: reduction in H2 (0.58 ml/min) in N2 as carrier gas ( 11.5 ml/min) at a heating rate of 10~ The catalytic test reactions were carried out in a microreactor under the following conditions: reaction temperature T = 330~ atmospheric pressure with 0.7 % o-xylene in air, and a space time W/F = 5.0 * 10-5 kgeat * 1-1 * h. Under test conditions the catalysts were diluted in quartz (1 : 7 wt. %).

3. RESULTS and DISCUSSION 3.1.1 SlhOa/TiO2

With the help of X-ray-Photoelectron-Spectroscopy (XPS) it is possible, because of the surface sensitivity of the method, to make in addition to the determination of surface oxidation states also some conclusions about the dispersion of surface species. The ratio of the signal intensities of the support material to those of the supported species is a measure of the dispersion [10]. The ratio of the signal intensities of Ti2pv2 and Sb3d3/2 decrease as a function of milling time. This is shown in Figure 1, for the case of a theoretical monolayer (1ML) of Sb2OflTiO2 as a function of the pretreatment, i.e., the duration of the milling. The decrease of the intensity ratio correlates with the increase of the dispersion of Sb-oxide on TiO:. In all instances of milling it is observed that the signal ratios in the beginning, i.e., at short milling times of l h to 5 h, experience a large change while the differences between 10 and 20 hours of milling are relatively small. It follows from this that for the ball mill employed, the milling times of l h to 20 h are sufficient to achieve dispersions which do not increase any further with milling times exceeding 20 h.

820

8t / n

[

--I---

t"

8

ctL,y ~

--V--" with 10 wt. % H ~ .~11_~_

m

--O--" dry ~ ! , ~

+ ..~ 400~ --&--" ~ " ~ ~ / + ~ 400~

-6

"-,,,... i,

C

0

5

9

10

15

20

L-O

duration of .~1]~r~g in h Figure 1:

XP-signal ratios Ti2p/Sb3d of 1 ML Sb203]TiO2 depending on the milling time.

The signal ratios of the dry milled materials are compared in Figure 1 to those of the calcined dry milled materials, as well as to the 10 wt % H20-milled materials, and the conventionally, by means of wet impregnation, prepared material. Where applicable, the calcination temperatures are indicated in the figure. It is apparent that the dispersion of Sboxide on TiO2 of the 20 h dry milled sample is comparable to that of the conventionally prepared sample. Comparable values are also obtained from the dry milled and subsequently calcined samples (5h 400~ With the impregnated sample, complete dispersion is reached already after 1 h of milling. An increase of the Sb-oxide dispersion as a function of milling time cannot be found with the calcined samples. The signal ratios of the milled samples in the presence of 10 wt. % H20 are also illustrated in Figure 1. The influence of the addition of water is clearly noticeable. The signal ratios of the samples milled in the presence of water are clearly different from those milled dry, when the milling times are kept short. Already after 2 h of milling, the materials milled in the presence of water reveal signal ratios comparable to the impregnated samples, the calcined samples or the 20 h dry milled samples. The signal ratios of the samples milled in the presence of 10 wt. % H20 which were

821 subsequently also calcined are not shown in Figure 1. The signal ratios of the latter samples are comparable to those obtained for the dry milled and calcined samples. Similar conclusions about the increase of the dispersion of reducible species such as Sb203 can be obtained by employing temperature programmed reduction spectroscopy (TPR). For a quantitative interpretation of the TPR profiles it is necessary to integrate the recorded signals. In this manner it is possible to assess the oxidation states of the respective samples through the uptake of the reduction equivalent. The observed oxidation changes of the materials investigated are illustrated in Table 1.

Table 1:

From TPR results derived changes in the oxidation state of Sb-oxide in the systeme 1 ML Sb-oxide/TiO2 depending on the pretreatment.

s~steme 1 ML Sb-oxide/'ri02 after 20h dry milling

chan~es of oxidation state

3.1

Sb(III) -~ Sb(O)

1 ML Sb-oxide/TiO2 after 20h dry milling + 20h 450~ in N2

2.9

Sb(III) -~ Sb(O)

1 ML Sb-oxide/TiO2 after 20h dry milling + 20h 450~ in 02

5.0

Sb(V) -~ Sb(0)

For the system 1ML Sb203/TiO2, milled dry for 20 h, it is possible to assess an oxidation change from Sb 3+ to Sb ~ whereas in the 02 calcined samples an oxidation state of Sb5§ was found, as expected. These changes in oxidation state can be also followed by means of XPS. For the dry milled samples there is no significant change to be observed in the binding energy of Sb3d3r2 signal (1 h milled, 539.4 eV; 20 h milled, 539.6 eV). This agrees well with the literature value for Sb203 of 539.4 eV [ 11]. When milling with 10 wt. % HE0, a change in the binding energy is observed between the 1 h and 2h milled samples. After 1 h of milling the Sb is still found to be in the +III oxidation state (Sb3d3/2 being 539.6 eV), while after 2h of milling a value of 541.8 eV is observed. The latter corresponds to the literature value for Sb205 of 541.6 eV for the Sb3d3/2 signal [12]. The same values for the binding energies are obtained for the Sbsignals of the calcined samples. This implies that no tribo-chemically induced oxidation of the Sb3+-oxide could be observed for the dry milled materials, however, oxidation was definitely observed in the calcined samples irrespective of the milling time. In contrast to these findings is the observation that unsupported Sb203 requires a significantly higher temperature, i.e., . 9 600 ~ to oxidize Sb3+-oxide to Sb4+-oxade. Neat SbS+-oxide can be obtained from Sb3+-oxide under dry conditions only when placed under oxygen at high pressures (sealed tubes, autoclaves). However, this oxidation occurs already at room

822 temperature with TiO2 supported samples milled for 1 h in the presence of water. In this case not only the mechanical activation of the oxide, but also the interaction of the Sb-oxide with the support material (TiO~) is of importance, since this tribo-chemical effect does not occur by milling only the neat Sb-oxide. The change in oxidation state of Sb can also be determined for the dry milled and calcined samples by means of X-ray absorption measurements (XANES), using the L~-edge (beamline E4, DORIS III, HASYLAB, Hamburg). Herewith, it is possible to recognize Sb 3+-, Sb4+- and SbS+-compounds through the shitt of the white line (transformation from 2s to 5p) in the reference compounds Sb203, Sb204, Sb205 and Na[Sb(OH)6], shifting from 4.7023 keV for Sb 3+ in Sb203 to 4.7068 keV for Sb 5+ in Na[Sb(OH)6]. The spectra of the 1ML Sb203/TiO2 dry milled samples exhibit the same white line shift as the Sb203-reference, irrespective if the milling was carried out for l h, 5h or 20h. Conversely, the calcined samples exhibit a distinct double peak structure, which suggests the existence of Sb 3+ and Sb5+ species. Through spectral simulation and comparison with reference compounds it is possible to assess the ratio of Sb 3§ to Sb5+. With about a 45:55 ratio, a stoichiometry of Sb2Oa+xis indicated [13], herewith, confirming by the XANES-investigations the XPS-results discussed above. The apparent differences in the oxidation states of the Sb-oxides, as measured by XPS and XANES, can be explained on the basis that the differentiation of Sb4§ and SbS§ in the XP-spectra is most difficult because of the small difference in the binding energies of only 0.3 eV between the two species [ 11]. However, Sb 5§ species in particular which reside on the surface of these systems lead, because of the surface sensitive XPS method, to an overweighting of the Sb 5+ oxide by this method.

3.1.2. VzOs/TiOz For the system V2Os/TiO2it is also possible to show with the help of XPS that milling leads to a distinct increase of dispersion of the V-oxide on the TiO2 carrier. The ratios of the signal intensities of Ti2p~/2 and V2p3/2 are depicted in Figure 2 as a function of the milling time. It is apparent that under dry milling conditions the dispersion increases significantly with time of milling. When l0 wt. % H20 is present during milling, the dispersion achieved after 1 h of milling is about equal to that achieved after 20 h of milling under dry conditions. The signal ratio of the wet milled sample does not change significantly with an additional 20 h of milling from that achieved after only 1 h of milling. In contrast, calcination of the samples enhances significantly the dispersion, which increases still further when mechanical activation preceded calcination. With long milling times the calcination step outweighs dispersion differences, whether the samples were milled dry or wet. In order to determine if a reduction of the vanadium oxidation state takes place with milling, starting with V205 on TiO2, it was necessary to subtract the interfering x-ray satellites in the region of the V2p signals. This is achieved with the aid of spectral fitting programs, which allow the different binding energies of V-oxides to be considered (reference values are: V2p3~: V2Os, 517.2 eV; V204, 516.0 eV; V203, 515.8 eV).

823

9-

8- \

_._:

" \

7_-

/

i ~ ~%/~

I

-A-: -v-:

m

I-9

1-8

~_-,., ,n.th io ,,t. ~, %0 ,,~m,,,a.th~.o,,e. , ~ o + ~ . , 4 . ~ * c

r

9

7 6 5

2 0

2

4

6

8

10

12

14

16

18

20

d u r a t i o n o f rail]~r~g in h XPS signal ratios Ti2p1/E/V2p3/2 of 1 ML V2Os/TiO2depending on the conditions of preparation.

Figure 2:

In the V2pa/2-signals of the V-oxide/TiO2 systems, a distinct increase in the linewidth as compared to the signals of the reference compositions V205, V204 and V203 with 1.6 eV is observed. This can serve as a basis for the differentiation of V 5§ and reduced V-species, since the difference in the binding energy between V 5+- and V4+-oxide is more than 1 eV, and is therefore easily distinguishable in the simulation of the spectra. Conversely, the differentiation between V4 - and V -oxtdes by simulation is, became of the small differences of the binding energies of only 0.2 eV, nvt+possible. Therefore, in the subsequent discussion, distinction will be made only between - and reduced V-species. Significant amounts of reduced V-oxide species are identifiable already at short milling times for the dry milled Voxide/TiO2 material. The signal ratio VS§ r~ remains independent of the milling time at about 30:70. In the 10 wt. % H20 milled samples, the signal ratio shifts to approximately 50:50, whereby the post calcined samples exhibit a reverse ratio of 70:30, irrespective of the presence or absence of water during the milling operation. Calcination at 450 ~ for 5 h causes a portion of the vr~L-Species in the vicinity of the surface to reoxidize to V 5+. Since the mechanically induced reduction of V205 starts at the surface, the ratios of V 5§ to v~d-species -

9

9

+

3+

9

824

is certainly too low in comparison to the total amount of V-species, i.e., V " ~ is with certainty overweighted.

3.1.3. SbzOz/VzOs/TiOz From the signal intensities it is apparent that the milling conditions employed significantly influence the V/Sb surface properties. If the activation is carried out in the dry state, then the Sb-content increases significantly with milling time. With th el0 wt. % H10 containing samples, a constant V/Sb-ratio sets in already atter short milling time, albeit at a lower level than with samples which were ball-milled dry. That implies that a still greater enrichment of Sb occurred on the surface. Comparable results are obtained when the samples are post-treated through calcination after milling, whereby the dry milled and calcined samples exhibit a constant, but smaller V/Sb ratio than the wet milled samples. These results are illustrated in Figure 3.

14 'U m

B

12

z ~

~o..~2%/']:~io 2

-II-

:d~

,,i 11 _,~

12

--V--" dry .~11~,.1 + 5h 400~ --O--: with 10 wt. % ~ O ~ 1 ~ = ~

10"

--A--: with 10 wt. % ~ O .~11=~ + 5h 400~

E 0

.r'l

-10

s~

4~ .

m.

20

Figure 3:

14

V ~ V

6 i :-----_* '

I

5

V

6

-* '

I

10

'

I

15

'

duration of milling in h

I

20

XPS signal ratios V2p3/z/Sb3d3/2of 1 ML Sb203N2Os/TiO2 depending on the conditions of preparation.

In the case of the dry milled and calcined samples, no V/Sb values could be determined for the lh to 3h milled samples, because the signals of V2p and Sb3d are not sufficiently large to be evaluated. Therefore, only the differences of the signal ratios of the individual samples are summarized in Table 2.

825 Ti2pl/z/Sb3d3/2 signal ratios of the Sb203N2Os/TiO2systeme depending on the preparation.

Table 2:

duration of milling

dry milled

dry milled + 5h 400~

milled with 10 wL % HzO

milled with 10wt. % H20 + 5h 400~

lh

9.1

--

1.5

1.7

20h

3.6

7.3

2.0

2.1

From the observed differences of the signal ratios it is recognizable that the treatment of the specific samples has a significant influence on the dispersion of active components on the surface of the carrier material. Particularly for the dry milled and calcined samples, the active components appear after calcination not to be accessible any more for XPS, because of the small signal intensities. This effect can be avoided through the addition of 10 wt. % H20. For the determination of the oxidation states of Sb and V, the spectra were analyzed as discussed above through curve fitting. The binding energies so detemined for the Sb3d3/2 signals are given in Table 3.

XPS Sb3d3/2 binding energy of the Sb203/V2Os/TiO2systeme depending on the preparation.

Table 3:

duration of milling

dry milled

dry milled + 5h 400~

milled with 10 wt. % H20

milled with 10wt. % H20 + 5h 400~

lh

539.6

--

540.0

539.9

20h

539.9

539.8

540.0

539.9

With the exception of the dry milled samples for short periods, the binding enel'gy of 3+ the Sb-signal rises, and values are found which lie between the Sb -oxide (539.6 eV) ~ the 4+ 9 Sb -oxide (540.3 eV) [12], independent of the preconditioning of the samples. For the V2p3a signal it is possible to show through fitting, that in all samples in addition to vS+-species also reduced V-species exist.

826 3.1.4 Catalyst Evaluation

As known from the literature, V2Os/TiO2 is, besides the Sb203/V2Os/TiO2 system, a well established catalyst in the PA synthesis [14, 15].Therefore two supported V-oxide-onTiO2 catalysts, one preparedby the milling method and the other by the conventional impregnation method, were tested for the oxidative conversion of o-xylene. The reaction was carried out in a microreactor under the conditions described above and the results are summarized in Table 4. Table 4:

Sample

results from o-xylene oxidation of a milled V2Os/TiO2- and a conventionally prepared V2Os/TiO2 catalyst.

Conversion in %

Y(o-tol.) in %

Y(PA) in %

Y(phthlide) in %

S(PA) in %

S(total OX.

prod. in % 1ML VEOs/TiO2 lhdry milled 1ML V2Os/TiO2 impregnated

Y(o-tol.): Y(PA): Y(phthalide): S(PA): S(total oxid. prod.):

15

2.8

8.0

1.2

53

80

20

2.9

8.9

1.3

45

65

yield of o-tolualdehyde. yield of phthalic anhydride yield of phthalide selectivity referring to phthalic anhydride selectivity referring to the total amount of oxidation products

It is apparent that the two catalysts give comparable yields of the desired PA product, under comparable reaction conditions. Therefore, it is concluded that the solventless ballmilling method described herein is a viable method of preparing selective oxidation catalysts.

4. S U M M A R Y In this study it was demonstrated that with the aid of solventless ball-milling of catalyst components, dispersions of active components on carrier materials could be achieved, as measured by TPR, XPS and XANES, that are comparable to the dispersions achieved through conventional impregnation techniques of catalyst preparation. Comparable catalytic results are obtained by both preparation methods. Specifically, catalysts supported

827 on TiO2 prepared by ball milling techniques yield phthalic anhydride in the oxidation of oxylene at comparable levels to conventionally prepared catalysts. Therefore, based on catalytic and spectroscopic results presented in this study, it is concluded that selective oxidation catalysts can be prepared by solventless ball-milling techniques which give comparable catalytic properties to those prepared by conventional solution impregnation methods. Addition of small amounts of water during the milling process enhances the rate of active phase dispersion on the catalyst carrier and thereby shortens the milling time required to attain the desired catalytic properties.

ACKNOWLEDGMENTS This work was financially suplx~ed by the Bayerische Forschungsverbund Katalyse FORKAT, by the Deutsche Forsehungsgemeinschaft (SFB 338) and by the Fond der Chemischen Industrie. R.K. Grasselli gratefully acknowledges the Alexander von Humboldt Stiftung for receipt of a Humboldt Research Prize.

REFERENCES

[1] [2]

[3] [4]

[5] [6] [7]

[8] [9] [lO] [11] [12] [13] [14] [15]

S.E. Golunski and D. Jackson, Appl. Cata., 48 (1989) 123. a. A.T. Guttmann, R:K. Grasselli, J.F: Brazdil and D.D Suresh, US Patent No. 4 746 641 (1998). b. A. Andersson, S.L.T. Andersson, G. Centi, R.K. Grasselli, M. Sanati and F. Trifiro, Proc. 10th Int. Congr. Catal., Budapest (Eds. L.Guczi et al.) 1992 A, 691. a. R.K. Grasselli and J.D. Burrington, Adv. Catal., 30 (1981) 133. b. R.K. Grasselli, J. Chem. Ed. 63 (1986) 216. K. Weissermel and H.-J. Arpe, Industrielle Organische Chemie, VCH Verlagsgesellschaft, Weinheim, 1994, 415. P.A. Zielinski, R. Schulz, S. Kaliaguine and A. Van Neste, J. Mater. Res., Vol. 8, No. 11 (1993) 2985. H.S. Horowitz, C.M. Blackstone, A.W. Sleight and G. Teufel, Appl. Catal., 38 (1988), 193. V.A. Zazhiggalov, J. Haber, J. Stoch, L.V. Bogutskaya, I.V. Bacherikova, Proc. 1lth Int. Cong. Catal., Baltimore (Eds. J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell), Elsevier, 1996 B, 1039. F. Roozeboom, T. Fransen, P. Mars and P.J. Gellings, Z. Anorg. Allg. Chem. 449 (1979) 25. D. Briggs and M.P. Seah, Practical Surface Analysis, John Wiley & Sons, Chichester, 1994, Part I: Auger and X-ray Photoelectron Spectroscopy. J.W. Niemantsverdriet, Spectroscopy in Catalysis, VCH Verlagsgesellschaft, Weinheim, 1993, 51. B. Visswanathan, S. Chokkalingam, T.K. Varadarajan and S. Badringarayanan, Surf. Coat. Technol. 28 (1986) 201. R. Izquierdo, E. Sacher and A. Yelon, Appl. Surf. Sci., 40 (1989) 175. U.A. Schubert et al., to be published in J. Phys. Chem.. M.S. Wainwright and N.R. Foster, Catal. Rev. -Sci. Eng., 19 (1979) 211. V. Nikolov, D. Klissurski and A. Anastasov, Catal. Rev.-Sci. Eng. 33 (1991) 319.

This Page Intentionally Left Blank

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

829

The Effect of Preparation Parameters on the BET Surface Area of Z r O 2 Powder YuanYang Wang YanZhen Fan

YuHan Sun*

SongYing Chen

State Key Laboratory of Coal Conversion, Institute of Coal Chemistry, Chinese Academy of Sciences, Taiyuan 030001, P. R. China

1. INTRODUCTION Zirconium oxides are extensively used in many fields (such as ceramic, refractory, sensor and catalysis) due to their propel~es[ 1]. Many methods have been developed for the production of such materials[2]. Among them, Rapid Thermolysis Approach (RTA) is a safe, simple and instantaneous route. Kingsley firstly reported the preparation of alumina and related oxides by RTA[3]. However, the samples derived showed a relative low BET surface area. ZrO 2 soproduced has been shown to be highly active towards CO oxidation and methane combustion. In order to have a full understanding of so-derived zirconia, the preparation of zirconia is in the first instance investigated in the present work. 2. E X P E R I M E N T A L

Zirconium oxychloride (or zirconium oxynitrate) was mixed with urea, nitric acid and water in an evaporating dish, and then thermolyzed in a muffle furnace at a fixed temperature about 10 minutes. This led to samples with a foam structure. The samples were then characterized by physical adsorption of N 2 at 77K (ASAP-2000), TEM (H-600 II) and XRD (D/max-TA). 3. RESULTS AND DISCUSSION In the RTA process, large amount of gases are produced by the decomposition of precursor and fuel (i.e. urea and nitric acid), which leads to a foam structure. Figure 1 illustrates the photograph of so-produced samples. The ZrO 2 foam is observed to brim over the container. 3.1. Effect of Precursor and Fuel Composition on BET Surface Area of ZrO 2 Powder

Table 1 shows that BET surface areas of samples change with the precursor and the fuel composition (i.e. urea, nitric acid or their mixture). Samples from zirconium oxychloride show * To whom all correspondence should be addressed

830 higher surface area than those from zirconium oxynitrate. Furthermore, BET surface area for ZrO(NO3) 2+ urea system are lower than those for ZrOCI~ + HNO 3+ urea system. This is in a good agreement with the reported by kingsley[3], in which BET surface area of pure A1203 sample was 8.30 mVg, and that of 50 wt% ZrO 2 in A1203 sample only 3.12 m2/g when zirconium oxynitrate was employed. These indicate that nitric acid plays important roles in the formation of the foam. Obviously, the foam structure and high BET surface area should be attributed to the synergetic effect of the decomposition of precursor and the combustion of fuel[4]. Figure 1. The photograph of ZrO 2 with foam structure Table 1 The effect of precursor and fuel composition on BET surface area of ZrO 2 powder* Sample No.

Precursor

H20/Zr 4§ HNO3/Z1"4§

CO(NH2)2/Zr 4§

SBzr (m~/g)

P1

ZrO(NO3) 2

10.0

0

4.0

3.82

P2

ZrO(NO3) 2

10.0

0

1.0

6.25

P3

ZI~)C12

27.5

4.0

4.0

8.33

P4

ZrOC12

13.8

4.0

1.0

35.31

* thermolyzed at 773K.

3.2. Effect of Solution Composition on the BET Surface Area of ZrOz Powder The effect of solution composition on the BET surface area of ZrO 2 powder is shown in Table 2. Obviously, with the decrease of the amount of urea or water in the presence of HNO 3, BET surface area ofZrO 2PoWder increases; and the concentration of nitric acid shows a optimum value. This indicates that less urea and water, with moderate nitric acid, are important for higher BET surface area. Figure 2 indicates the difference of two types of ZrO 2 samples in their morphologies. Although the particle size for both are in nanometer-size, their morphologies and BET surface

831 Table 2 The effect of solution composition on the BET surface area of ZrO 2 powder* Sample No.

H20/ZrOC12

HNO3/ZrOC12 CO(NH2)2/ZrOC 89 SBrr (m2/g) i

S1

0

4.0

0

$2

0

4.0

0.5

150.46 51.22

$3

0

4.0

1.0

40.15

$4

0

2.0

1.0

22.20

$5

0

8.0

1.0

24.10

$6

13.8

4.0

1.0

35.31

$7

27.5

4.0

1.0

27.47

* thermolyzed at 773K. area differ from each other seriously. Samples produced in the presence of urea consist of 50 nm of particles with crystal structure whlie that made without urea is highly-dispersed. It is clear that these defferences are closely related to the praparation mechanism. At high temperature, the following reactions take place in the system of precursor and fuel[5]: CO(NH2) 2 + H20 ~ HNO 3

> CO2(g) + 2NH3(g)

> NO2(g) + NO(g) + O2(g) + H20(g)

Figure 2. The TEM graphs of samples S 1 and $7

(1) (2)

832 CO(NH2) 2 + HNO 3

> CO(NH2)2 HNO 3

(3)

CO(NH2) 2 HNO 3 + 3HNO3 > CO2(g) + 4NO2(g) + 2NH3(g) + H20(g) + O2(g) ZrOC12 + 2HNO 3 > ZrO(NO3) 2 + 2HCI(g)

(4)

ZrO(NO3) 2

(6)

> ZrO 2 + 2NO2(g) + 1/2 O2(g)

(5)

In the present of urea, large amount of gases are produced to form a foam structure, and the process is then controlled by combustion mechanism, the reactions (5) and (6) hardly influence the process. Without urea in the sample (see S 1 in Table 2), only reactions (2), (5) and (6) occur, ZrO~ powder is mainly made by the decomposition of ZrO(NO3) 2 and no foam structure appears, the process is controlled by decomposition mechanism. When the samples are prepared through combustion route, the reactions between nitric acid and urea gives rise to a very high temperature which even reaches 1873K or so[3]. Obviously, this high temperature results sintering and aggregation of ZrO2 powder, leading to large particles and then low surface area. With the decomposition route, ZrO2 powder can be produced by the decomposition of zirconium oxynitrate at about 773K. At such a mild temperature, small particles in a highly-dispersed state, of course, show high BET surface area, which is even higher than that of ZrO2 aerogels ( about 100 m2/g) prepared by sol-gel method involving supercritical drying at the same temperature[6]. The amount of water also influences the BET surface area. The BET surface area of ZrO2 is found to decrease with increasing the amount of water (see TaNe 2). This may be due to that water might retard the release of gases and the rapid expansion of the mixture. 3.3. Effect of Thermolysis Temperature on the Texture of ZrO 2 Powder

The thermolysis temperature is found to have a strong influence on the BET surface area ofZrO 2powder. Samples (Tl~and T2) thermolyzcd at 573K and 773K almost show the same BET surface area, and their surface areas are much higher than that of sample T3 produced at 973K (see Table 3). However, it is interesting that sample T2 shows a concentrated pore distribution around 3.5nm while the bimodal pore distribution is observed for samples T1 and T3 at 3.5nm and 6.5nm (see Figure 3). Furthermore, the pore volume of T3 is much lower than those of T1 and T3 (see Table 3). These imply that their BET surface area originates differently from each other. Considering that the preparation of the samples follows the decomposition mechanism, the difference should be related to their dispersion degree and crystal structure. As mentioned above, the sample produced via thermolysis at 773K consists of homogeneous highly-dispersed particles (see Figure 2). Consequently, sample T2 displays a concentrated pore distribution around 3.5nm, which produces high surface area. With the sample thermolyzed at high temperature(i.e. 973K), however, the sintering and aggregation of particles

833 Table 3 The effect of thermolysis temperature on the texture of ZrO 2 powder* Therm.

SB~

Pore

Pore vol.

Crystal size**

t/m**

temp.(K)

(m2/g)

size(nm)

(cm3/g)

(nm)

ratio

T1

573

146.16

7.31

0.27

. . . . . . . .

T2

773

150.46

5.19

0.20

7.60

3.13

T3

973

44.31

9.71

0.11

12.40

2.52

Sample no.

* H=O:HNO3:CO(NH2)2:ZrOC12= 0:4.0:0:1.0 ** calculated based on the XRD results. takes place. This leads to inhomogeneous growth of highly-dispersed particles and then a bimodal pore distl'ibution with most of pores around 6.5nm, which causes the decrease of the BET surface area of the sample. In the meantime, the phase transformation of tetragonal to monoclinic produces more monoclinic phase in samples T3 (see figure 4 and Table 3), and leads to the formation of large partiicles due to volumetric expansion[7]. Both result in a sharp decrease of BET surface area.

.....

O -- T e t r a g o n a l X -- M o n o c l i n i c

TI

. . . . . . . . . T2 ~

0.6

E 0,4

."" .,

.,'\

-!

(}.2

/"~

~

:-,'; : i,' ', : ," '~

" t

;; I'

'. '.

T3

o

,.

~. I

o

A

o

[

I

I T3

|~

1o

20

30

40

Pore diameter(nm)

Figure 3. The pore distribution of samples with different thermolysis temperature

50

6O

70

20

Figure 4. The XRD patterns for samples with different thermolysis temperature

In the case of sample T1, no such an aggregation occurs because of its lower thermolysis temperature of 573K even if it shows a similar pore distribution to sample T3. This may be caused by the inhomogeneity of amorphous ZrO2 particles due to the thermolysis temperature less than the crystallization temperature of 743K (see Figure 4). Therefore, the high BET surface area should be attributed to the high dispersion of amorphous ZrO2 particles.

834 4. CONCLUSION Rapid Thermolysis Approach is one of best routes to produce ultrafine ZrOr The samples so-derived are of nano-sized particles with high BET surface area. BET surface area of powder changes with the system of precursor and fuel. Synergetic effect of the decomposition of precursor and the combustion of fuel leads to foam-structured powder with high BET surface area. The presence of urea and nitric acid is very important. Thermolysis temperature has a strong influence on the texture of ZrO~ powder. 773K is the proper temperature for the preparation of ZrO~ powder with high BET surface area and concentrated pore distribution. ACKNOWLEDGEMENT

The authors acknowledge the National Natural Science Foundation of China for its financial supports. REFERENCES

1. S. S. Prakashi, C. J. Brinker, A. J. Hurd, Nature, 374(1995)439. 2. D. A. Ward, E. I. Ko, Chem. Mater., 5(1993)956. 3. J. J. Kingsley, K. C. Patil, Mater. Lett., 6(1988)427. 4. J. J. Kingsley, K. Suresh, K. C. Patil, J. Mater. Sci., 25(1990)1305. 5. A. M. Wynne, J. Chem. Educ., 64(1987) 180. 6. H. W. Xiang, B. Zhong, S. Y. Peng, Mole. Catal.(China), 8(1994)263. 7. P. D. L. Mercera, J. G. V. Ommen, E. B. M. Doesburg, Appl. Catal., 78(1991)79.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

Preparation of VOHPO4.0.5H20 and (VO)2P207 performance for maleic anhydride synthesis

835

and

their

catalytic

T. Miyake and T. Doi Tosoh Corporation, Yokkaichi Research Laboratory, 1- 8 Kasumi, Yokkaichi- shi, Mie 510, JAPAN

Firstly, influence of alcohol solvents on characters of VOHPO4.0.5H20 prepared from V409 and ortho-HaPO4 was studied by valence of vanadium, XRD, SEM, TG/DTA, ~ P-NMR and Raman spectra. It was revealed that morphology and bulk characters of VOHPO4.0.5H20 differed among alcohols used. Then, oxidation of n-butane to maleic anhydride on (VO)2P207 obtained by transformation of VOHPO4-0.5H20 was investigated. Selectivity of maleic anhydride was independent of the ratio of surface areas; namely, so-called selective to non-sective surface of (VO)2P207.

1.INTRODUCTION Vanadium-phosphorus oxides are known as the effective catalysts for the production of maleic anhydride by oxidation of n-butane [1-3]. Especially a crystalline (VO)2P207 is claimed to be the active and selective phase [4,5], though some research groups insist that VOPO4 plays an important role [6]. Generally three methods are known to synthesize (VO)2P207; aqueous solvent method [4], organic solvent method [5] and VOPO4.2H20 reduction method [7]. It is reported that VOHPO4.0.5H20, precursor of (VO)2P207, and (VO)2P207 prepared by these methods have different surface and bulk characters. For example, Horowitz et al. [7] reported that VOHPO4-0.5H20 of different morphologies were obtained by varying P/V ratio and organic solvents in the VOPO4-2H20 reduction method. In the organic solvent method, Kesteman et al. [8] concluded that the primary factor which influenced crystallinity and morphology of precursor was the choice of alcohol used during dehydration-condensation of vanadium compound and ortho-H3P04. From an industrial viewpoint, as the selectivity of maleic anhydride is around 60%, the improvement is strongly desired. Recently, Igarashi et al. [4] and Bordes [9] proposed that the selectivity depended on the ratio of selective crystalline face to non-selective one of (VO)2P207. More precisely, on (200) mainly maleic anhydride was produced and on (042) thus-produced maleic anhydride was over-oxidized to carbon oxides. When this hypothesis is accepted, the morphology of (VO)2P207, and therefore that of the precursor VOHPO4"0.5H20, might play a very important role for the catalytic performance. In this study, at first in order to obtain a pure and crystalline precursor having various (001)/(220) ratio, precursors were synthesized using various alcohol solvents

836 during dehydration-condensation of V409 and ortho-H3P04. Then, the relation between the (200)/(042) ratio of (VO)2P207 and the selectivity of maleic anhydride was investigated with the pure (VO)2P207 transformed from the precursors.

2. EXPERIMENTAL 2.1. Preparation of VOHPO4.0.5I-hO Typical preparation procedure was as follows [10]: Firstly, V205 (91 g) was reduced to V409 in isobutyl alcohol (280 g) under reflux for 48 hours in N2 stream. After filtration and washing with acetone, the black solid V409 was dried at 313 K for 15 hours. Then the solid V409 (20 g) was suspended in 100 g of various flesh alcohols (isopropyl alcohol, isobutyl alcohol, 2-butyl alcohol, 2 - m e t h y l - l - b u t y l alcohol, cyclohexyl alcohol or 2 - e t h y l - 1 - h e x y l alcohol). The suspension was heated until 383K. (In the cases of isopropyl alcohol and 2-butyl alcohol, the suspension was heated until reflux.) Then, here added an ortho-H3P04 (18.2 g) in the same alcohol (7.2 g) dropwisely for 30 minutes. The ratio of P to V was 1.13 for preparation of the precursor. After H3P04 addition, this condition was kept for 5 hours (in the case of isopropyl alcohol, for 24 hours). Then the obtained slurl~ was filtered. Thus obtained solid was washed with acetone and was dried at 313 K for 15 hours. Precursors were also prepared by the aqueous solvent method and VOPO4.2H20 reduction method according to the literatures [4, 7]. To obtain pure precursors, special care was given to mixing, washing and so on. 2.2. Characterization Elemental analysis for V and P of the precursor and (VO)2P207 was carried out using inductively coupled plasma spectroscopy (ICP) on Kyotokouken UOP-2. According to the method of Hodnett [11], the average valence of vanadium was determined by the double titration method using KMn04. Powder X - r a y diffraction patterns were recorded with Mac 'Science M18XHF diffractometer using C u - K c~ radiation (40 kV, 100mA). Crystallite size of VOHPO40.5H20 and (VO)2P207 was calculated by Scherrer equation. Scanning electron microscopy (SEM) analysis was carried out with Shimadzu TPM 810. Thermogravimetfic (TG) analysis was carried out using SEIKO TG/DTA-220. 31 p_ NMR measurements were carried out with Nippondenshi JNM-GSX-270WB. Raman spectrum was recorded on I. S. A. Jobin Yvon Ramnor-U1000 spectrophotometer. The emission line at 514.35 nm from Ar laser was used for excitation. 2.3. Catalytic performance Precursors were transformed to (VO)2P207 by heat-treatment in N2 stream or n - b u t a n e - a i r stream. Into the tubular reactor, a 10 g portion of (VO)2P207 was placed. When gas component was arranged to 1.5% n-butane at space-velocity of 1500 h - 1 , the temperature was raised to the desired one. Products were analyzed by GC and LC.

837 3. RESULTS AND DISCUSSION 3.1. Influence of alcohol solvents on characters of VOHPO4-0.5H20 The P/V ratio and the average valence of vanadium of precursors prepared by the organic solvent method in various alcohols are shown in Table 1. In every case, the P/V ratio was 1.09 + 0.08. The average valence of vanadium of precursors were substantially 4+ irrespective of the alcohol used. Although there remained a small portion of alcohol between the layers (vide infra), these alcohols had substantially no influence on the measurement of valence (<0.05).

Table 1 Properties of precursors prepared in various alcohols by the organic sol vent method Solvent P/V atomic Average valence Weight decrease* ratio / of V / wt% 3.96 10.3 Or-1 isobutyl alcohol 1 01 4. O0 10.4 Or-2 isopropyl alcohol 1 10 3.99 10. 6 Or-3 2-methyl-l-butyl alcohol 1 08 4. O0 12.0 Or-4 cyclohexyl alcohol 1 16 3.97 11.8 Or-5 2-ethyl-l-hexyl alcohol 1 08 3.98 12.0 Or-6 2-butyl alcohol 1 15 ;9 Calibrated TG weight decrease between 473 and 803 K (cf. Figure 2)

Figure 1 shows the XRD patterns of precursors. When isobutyl alcohol, isopropyl alcohol and 2 - m e t h y l - 1 - b u t y l alcohol were used, only (220) peak was sharp and all the other peaks were broad. This suggested that the precursors prepared in these alcohols were relatively thin perpendicular to (001) face. On the other hand, when cyclohexyl alcohol, 2 - e t h y l - 1 - h e x y l alcohol and 2-butyl alcohol were used, all the XRD peaks of the precursors were very sharp and crystallinity seems to be improved. In addition, peak intensity of (001) became relatively high to that of (220) and this suggested that the precursors prepared in these alcohols were thick compared to those obtained in the former alcohols.

Or-6 *~ [>" Or-5 "~ Or-4

L.

o

-I

_ ~,t ~

10

_ 20

~.~ ..........

~

30

40

50

2 0/degree

Figure 1. XRD patterns of precursors prepared by the organic solvent method. (Notation; see Table 1)

838

Figure 2 shows the result of TG/J3TA measured in N2 stream. It is well known that transformation of VOHPO4-0.5H20 to (VO)2P207 occurs between 520 to 770 K [12]. Although this weight decrease is clearly seen in Figure 2, the rate differed among precursors. In the cases of precursors prepared in isobutyl alcohol, isopropyl alcohol and 2 - m e t h y l - 1 - b u t y l alcohol, the weight decrease curve was not sharp and this suggested that crystallinity of the precursor was not so good. To the contrary, in the cases of precursors prepared in cyclohexyl alcohol, 2 - e t h y l - 1 - h e x y l alcohol and 2-butyl alcohol, the weight decrease curve was very sharp and this suggested that the crystallinity was rather high. Differences in sharpness of the TG curves were well related to those of XRD peak intensities; the stronger the XRD peak intensity was, the sharper the weight decrease curve was. In Table 1, the TG result is summarized. Calibration was carried out by neglecting the weight decrease below 520 K and the theoretical weight decrease which corresponds to the transfomation is 10.5 wt%. T h e calibrated weight decrease for the latter three alcohols (namely, cyclohexyl alcohol, 2 - e t h y l - 1 - h e x y l alcohol and 2-butyl alcohol) was slightly higher than the theoretical value and this suggests that a small portion of the alcohol solvent remained betWeen the layers.

r~ r162

95

-o 90

Or-2

",

Or-3

~ 85 [-.,

Or-4

80

273

473

673

873

273

473

673

873

1073

Temperature / K

Figure 2. TG/DTA analysis of precursors prepared in various alcohols by the organic solvent method. (Notation; see Table 1)

Figure 3 shows the 31 p_ NMR spectra of precursors prepared in various alcohols. The precursors prepared in alcohols other than isobutyl alcohol showed the characteristic peak at - 1 3 8 ppm which could be assigned to VOHPO4-0.5H20 and no peak assigned to VO(H2P04)2 was observed. Therefore, we could say that purity of these precursors was high and short range order of VOHPO4-0.5H20 crystal was good, especially for 2-butyl alcohol and 2 - e t h y l - 1 - h e x y l alcohol. When intensities of the peak are compared, they are proportional to the crystallinity estimated from XRD. Here we would like to consider 3, P - M R result together with those from XRD and TG/DTA. Although we ignored the TG weight decrease below 520 K, especially in cases of isobutyl alcohol, isopropyl alcohol and 2 - m e t h y l - 1 - b u t y l alcohol, substantial weight decrease seems to exist. If we estimate that these weight decreases are from the alcohols remained between layers, precursors prepared in these alcohols could have

839

2 - butyl alcohol

LI~[~

2 - m e t h y l - 1 - butyl alcohol

.

j

isopropyl alcohol -. J . ~.L ,i , k...ua..,,.lt,dll.lk.lu~ld.klhahlh,di~k~ljlhk~l~iAl~lJd~htt.,.,I,,jLti,,,

~,~w'T"""*q"'"'"'~'"r ......r r l ~ , , W g ~ ~

,

isobutyl alcohol

300

200

100

0 -100 Chemical shift

-200

-300

Ppm

Figure 3. a, p_ NMR spectra of precursors prepared by the organic solvent method.

isobutyl alcohol

]

cyclohexyl alcohol

2 - e t h y l - I - hexyl

rather poor XRD peak intensities and ~1 P - N M R peak intensities. Figure 4 shows the Raman spectra of precursors prepared in various alcohols. In the case of 2 - e t h y l - 1-hexyl alcohol, it was difficult to measure Raman spectra of high resolution because of the fluorescence from the precursor. From the spectra observed, it should be emphasized that all the peaks ascribed to the precursor [13] are observed and no other peaks stem from impurities such as VOP04 phases were seen. Generally speaking, the intensity of the main peak assigned to V=O was stronger for alcohols which gave poor XRD peak intensities and ~ P - N M R resolution and vice versa. For example, Raman peak intensity of precursors prepared in isobutyl alcohol or cyclohexyl alcohol was strong while 31 P - N M R peak intensity of the precursor prepared in either of these alcohols was weak. In case of 2-butyl alcohol or 2 - e t h y l - 1 - hexyl alcohol, the opposite relation was observed. Considering that both XRD and ~1 p _ NMR are characterizing the bulk and that Laser Raman is sensitive to the surface, this phenomenon might be reasonably explained; the precursors having intensive XRD peak are thick and gives weak Raman intensity because of relatively low abundance of V=O alcohol on the surface.

q,)

2- methyl- I-

[

. . . . 2 - butyl alcohol

T

~a~k_-~,, ,-~,.,-,~ L ~ , . ,I ~',,~-~,wr~,~'~r't IW~"q~il

r,.,~r r ~

Figure 4. Raman spectra of precursors prepared by the organic solvent method. 100

700

1300 100

Raman shift / cm -'

700

1300

840

.............

....

;:~::~::~;~:;%~::::::~ii!:~i'.:,ii i ;:.4ii!~!ii i:~:~ii i;~i~F~i~.::::Niiiii? ili~i ::~U i~ ~!i~i ' ~iiii!iiiii',~.... iii~ii ...... ~,~

i-~-

i~iiiiN,

"

.......

isobutyl alcohol

cyclohexyl alcohol

isopropyl alcohol

2- ethyl- 1- hexyl alcohol

,.

.

.

.

m

..........~..-.

.

...... i:ii~!i ::i:~i:}~i:!i!:;ii~.~{::!~iii '?ii:i!i~:.:i~:~ ........

~:~:~,. :

~:::

2 - methyl- 1- butyl alcohol

i

m

2- Butyl alcohol

Figure 5. SEM photographs of precursurs prepared by the organic solvent method.

841 Figure 5 shows SEM photographs of precursors prepared in various alcohols. It is apparent that in every case the particles are of the same form and no other forms from impurities were observed. In the cases of precursors prepared in isobutyl alcohol, isopropyl alcohol and 2 - m e t h y l - 1 - b u t y l alcohol, the morphology of the precursor was petal- or rose-like. On the other hand, the morphology of the precursors prepared in cyclohexyl alcohol, 2 - e t h y l - 1 - h e x y l alcohol and 2-butyl alcohol was plate-like and rather flat. This was in accordance with the XRD results. In addition, it should be noted that the morphology of precursors were very similar to that of precursors prepared in the corresponding alcohol by the VOPO4-2I-hO reduction method [7]. From these results, it can be said that precursors are pure and uniform.

3.2. Relation between the ratio of surface area of selective face to non-selective one and maleic anhydride selectivity It is sometimes reported that to obtain steady catalytic performance hundreds of hours are necessary for activation of a catalyst under n-butane oxidation reaction [14]. Therefore, firstly the selectivity of the catalyst whose activity became stable at high n-butane conversion (>90 %, ca.100 hours) was compared with that of the catalyst activated under conventional procedure for about 500 hours (Figure 6). As the same maleic anhydride selectivity was obtained between these two catalysts at the same 60% n-butane conversion, we adopted this short-time activation procedure.

80

Sel.

6o

.s

--0--,-0-

;>

40

0

~0~.0---0

2o

m

--0--'0-

Conv.

m,,

320

q)

0 0

340 ~

~

360 ~

I

I

I

I

100

200

300

400

Reaction time / hr Figure 6. Activity and selectivity change during reaction Pressure; 101.3 kPa, n-Butane; 1.5%, GHSV; 1500 hr 1

--~

500

842

In Table 2, the catalytic performance at 60% n-butane conversion and the properties of catalysts after the reaction are summarized.

Table 2 Properties of catalysts and their catalytic performance Preparation Solvent TransV a l e n c e Surface Surface-area Temp. Conv. Sel formation of V Area(m2/g) ratio(-) ~ % % Aqueous H20 N2, 650~ 4.0 15.2 1.08 410 65.1 59.9 Organic isobutyl alcohol Butane-air 4.3 29.1 0.64 400 64.6 63.5 solvent 2-Butyl alcohol Butane-air 4.5 17.3 0.52 410 62.3 65.7 c-Hexyl alcohol Butane-air 4.1 20.2 0.83 400 59.5 63.0 MOP04 Benzyl alcohol N2, 800~ 4.0 7.2 0.82 440 61.0 58.0 Reduction Benzyl alcohol N2, 650~ 4.0 29.7 1.45 360 60.3 59.3 Benzyl alcohol N2, 650~ 4.0 28.2 1.44 380 60.3 61.4 Benzyl alcohol N2, 650~ 4.0 24.6 1.35 375 58.1 60.1 2-Butyl alcohol N2, 650~ 4.0 14.8 1.20 430 58.8 53.0 Surface-area r a t i o Crystal was tentatively considered to be a disc whose diameter and thickness were equal to the size of (042) and (200), respectively. Then, surface-area ratio is calculated as follows. Surface-area ratio = 2 * n ((042)/2)- / n 9(042) * (200) = (042) / 2 9(200)

As is shown in Figure 7, conversion was dependent on preparation procedure; more (Figure 8). This indicates that

it was indicated that the temperature for 60% n-butane the surface area of (VO)2P207 irrespective of the catalyst precisely, dependent on the surface area of (200) face (200) face is active and other faces are rather inactive.

440 420 ~

400

~ 380 => 360

~ 8:340

A O; aqueous solvent method r-l; organic solvent method

320 - A; VOPO4"2H20 reduction method 300

I

5

I

I

I

10 15 20 Surface area / m2g1

I

25

Figure 7. Relation between the temperature for 60 % n-butane conversion and surface area of the catalyst. Pressure; 101.3 kPa, n-Butane; 1.5%, GHSV; 1500 hr -1

30

843

440 420 ,.Q

a~

~176 400

a

0

0 ~ -~ 380

r

~>360 g

~

340

~

320 30O

D

O; aqueous solvent method !-1; organic solvent method A. VOPO4. 2H20 reduction method I

!

~"~'-

!

5 10 15 Surface Area of (200) face / m2g-1

20

Figure 8. Relation between the temperature for 60 % n-butane conversion and surface area of (200) face of the catalyst. Pressure; 101.3 kPa, n-Butane; 1.5%, GHSV" 1500 hr -1

No correlation between the maleic anhydride selectivity and the ratio of surface areas of selective to non-selective faces was observed (Figure 9). This suggests again that maleic anhydride is formed only on the so-called selective (200) face and it is not possible to improve the selectivity simply by increasing the surface area of (200) face.

844

4. CONCLUSIONS Depending on the alcohol solvent used for preparation, VOHPO4.0.5H20 of different characters were obtained from V409 and ortho-H3P04. Especially the precursor having stronger XRD peak intensities gave sharper TG weight decrease curve and higher 3~ P - N M R peak intensity, which meant that the precursor of this kind was pure, uniform and highly crystalline. Raman spectra indicated purity of precursors was high. Maleic anhydride selectivity did not depend on the ratio of the surface area of selective to non-selective faces and it was suggested that the maleic anhydride selectivity is not improved simply by increasing this ratio.

9 70

>

65

-~ 60

O

~.

ZX

55 O; aqueous solvent method .~ 50 -[2]; organic solvent method ZX; VOP04" 2H20 reduction method I 45 0.5

a

1.0 Ratio of selective surface area to non-selective one / -

Figure 9. Relation between maleic anhydrie selectivity and the ratio of selective surface area to non-selective one.

1.5

845 REFERENCES

1. 2. 3. 4. 5. 6. 7.

E. Bordes and P. Courtine, J. Catal., 57(1979)236. G. Centi and F. Trifiro, Chem. Rev., 88(1988)55. G. J. Hutchings, Appl. Catal.,72(1991)1. H. Igarashi, K. Tsuji, T. Okuhara and M. Misono, J. Phys. Chem., 97(1993)7065. G. Busca, F. Cavani, G. Centi and F. Trifiro, J. Catal., 99(1986)400. G. Centi, Catal. Today, 16(1993)5 and references therein. H. S. Horowitz, C. M. Blackstone, A. W. Sleight and G. Teufer, Appl. Catal., 38(1988)193. 8. E. Kesteman, M. Merzouki, B. Taouk, E. Bordes and R. Contractor, Scientific Bases for the Preparation of Heterogeneous Catalysts 6th. Intl. Symp., Poster Session No. 1, 301 9. E. Bordes, Catal. Today, 16(1993)27. 10. T.Miyake and T.Doi, Appl. Catal., 131(1995)43. 11. B. K. Hodnett, Ph. Permanne and B. Delmon, Appl. Catal., 6(1983)231. 12. F. Cavani, G. Centi, F. Trifiro and G. Poll, J. Them. Anal., 30(1985)1241. 13. F. B. Abdelouahab, R. Olier, N. Guilhaume, F. Lefebvre and J. C. Volta, J. Catal., 134(1992)151. 14. M. Abon, K. E. Bere, A. Tuel and P. Delichere, J. Catal., 156(1995)28.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

847

H y d r o x y l a t i o n o f B e n z e n e on Z S M 5 T y p e Catalysts M.H~ifele, A.Reitzmann, E.Klemm and G.Emig Lehrstuhl for Technische Chemie I, Universit~it Erlangen-Ntimberg, Egerlandstr. 3, 91058 Erlangen, Germany

ZSM5 type zeolites were used as catalysts for the one-step synthesis of phenol by benzene partial oxidation with nitrous oxide. Isomorphous substitution of A13+ ions by other trivalent metal ions revealed a high catalytic performance of the H-Ga-ZSM5 in a wide temperature range. Systematic variation of the partial pressures of the reactants led to satisfactory preliminary kinetic models. Deactivation could be reduced by postsynthetic catalyst silylation which is believed to block the strongest acid sites responsible for coke formation. 1. INTRODUCTION With worldwide phenol consumption exceeding 5 million tons in 1995, optimizing production routes of this essential chemical becomes very important. As an alternative to the traditional cumene process, a one-step-synthesis of phenol from benzene is highly desirable. With a ZSM5 type zeolite in its acid form as catalyst and nitrous oxide as oxidant, benzene may be directly oxidized to phenol [ 1-4]: C6H6 + N20

ZSM5)C6HsOH + N2

(Eq. 1)

The vast economical potential has led Monsanto to consider commercializing the process [ 17]. Discovering the high potential of a one step phenol synthesis (Eq. 1), Ono and coworkers [2] improved the original process that was based on vanadium pentoxide catalysts (comparing nitrous oxide and oxygen as oxidants [8]). It is generally agreed upon that ZSM5 type catalysts are outstanding among a wide range of metal oxide and zeolite catalysts in this gasphase-reaction [9-11 ]. Besides H-A1-ZSM5 [ 1] especially H-Fe-ZSM5 and H-A1/Fe-ZSM5 [4] zeolites were used as catalysts. Only little attention has been paid to H-Ga-ZSM5 [5,12] even though element electronegativity and ion size of Ga 3+ suggest a similar behaviour of H-GaZSM5 zeolites compared to H-A1-ZSM5. Depending on the catalysts used, several quite different mechanisms of the hydroxylation reaction have been suggested: The group of Ono [2], Burch and Howitt [1] and Tirel et al. [3] using H-A1-ZSM5 proposed an acid catalyzed mechanism. Panov et al. [4] using iron containing ZSM5 zeolites advocate a redox mechanism mainly catalyzed by iron species. While the suggested mechanisms are quite different, similar experimental results were obtained concerning the dependence on temperature and partial pressures of the reactants [ 1,2,13]. Furthermore, appreciable coke formation led to a loss in catalyst activity in all cases [2,14,15].

848 In addition to the effects of different catalysts on the hydroxylation reaction, which are still discussed controversially in the literature [ 16], the knowledge of the influence of the reaction conditions, especially temperature and reactant partial pressure, is also important. In this paper detailed reaction engineering investigations with a H-Ga-ZSM5 are hence presented [5-7]. 2. EXPERIMENTAL All zeolite samples were synthesized by the hydrothermal method described in detail in [6]. The experiments were performed in a completely automated laboratory setup including an integrally operated plug flow tubular reactor. Reaction components were analyzed by on-line gas chromatography with FID and TCD [5-7]. Table 1 summarizes the reaction conditions for the benzene hydroxylation on the H-Ga-ZSM5 catalyst. Nitrogen was used as balance.

Table 1." Reaction conditions (W=weight o vst, F=total molar flow) catalyst modified residence time ~ total reactor benzene nitrous oxide composition W/F pressure temperature concentration concentration [g'min/mol] [Pa] [~ [%] [%] SiO2/Ga203=80

0-126

103

350--450

2-12.5

4-26

3. RESULTS AND DISCUSSION Similar observations were made with all catalysts used. Hydroxylation occurred with high selectivities and yields to phenol. Except phenol only benzoquinone was produced in higher yields. Formation of CO2 was not observed below 400~ Initial phenol yields were about 25% but deactivation by coke formation quickly led to a decreasing activity. 3.1 Comparing different ZSM5 zeolites Hydroxylation of benzene with nitrous oxide on ZSM5 type zeolites was strongly influenced by catalyst composition and modification, e.g. type of framework metal ion, type and strength of acid sites, molar ratio of Si/A1 and pretreatment conditions [1-5]. Substitution of the framework aluminum by other trivalent metal ions (especially gallium and iron) had a significant influence on reaction performance. Activity increased in the sequence H-AIZSMS
849

--

. - - x - - . H-AI-ZSM5

H-Ga-ZSM5 - ~

- H-Fe-ZSM5

25 o

,x

W/F:

2O

92 g min/mol X

o (" (D 15 rQ.

06H6:42 %

"'"X" "

%

N20:26

t i m e on stream:

40 min

X"

O 10 "13 (D >5

0 325

q

~

i

I

t

350

375

400

425

450

Temperature [~

475

Figure 2 Comparison of phenol yield on H-Fe-ZSM5, H-A1-ZSM5 and H-Ga-ZSM5 as function of temperature The effects of a variation of the main reaction parameters temperature, nitrous oxide partial pressure and benzene partial pressure were studied as a first step towards a detailed kinetic model, as this is required for the design of an industrial hydroxylation reactor. The choice of the proper reaction conditions can significantly increase phenol production [6]. 3.2. Effect of reaction temperature Increasing the reaction temperature from 300 to 450~ led to an increase in reaction rate. Benzene conversion reached values up to 50% at 450~ with the fresh catalyst (Figure 3). x 350 ~

9375 ~

a 400 ~

9425 ~

o 450 ~

5O

|

H-Ga-ZSM5 W/F: 92 g min/mol

40

N

N20:26

~ -Q

3o

t~

20

o

10

"

C6H6:4.2 %

O g

%

t

0 0

20

40

60

80

'

,,,, t 100

Time-on-stream [min]

IX 120

I 140

/ 160

Figure 3 Benzene conversion as function of catalyst time on stream at different temperatures

850 Like the hydroxylation rate, the deactivation rate, too, increased with temperature. After 150 minutes time-on-stream of the catalyst, benzene conversion at 4500 C was very close to the conversion at 400~ Complete regeneration was achieved by heating the catalyst to 500~ under a flowing oxygen/nitrogen mixture (ratio 1:5). A comparison of benzene conversion on nondeactivated catalysts (W/F= 92 g min/mol, N20=26%, C6H6=4%) shows that conversion increased from about 15% to near 50% by increasing the temperature from 350 to 450~ Figure 4 shows yields and behavior of the product distribution at the beginning of the reaction (time-on-stream = 40 min) when varying reaction temperature. --o-- phenol

~

benzoquinone - x -

catechol ~

resorcinol ~

carbon dioxide

30

5

25

4o"e 00 to

20

3 -o

,_.__.

-o 15 >..

o t~ >, t__

m

2

. = =

10

"

H-Ga-ZSM5 W/F: 92 g min/mol N20:26 %

C6H6:4.2 % time on stream: 40 min

o -o

=.=.,

1 .-~ >.. 0

0 325

350

375

400

425

450

475

T e m p e r a t u r e [~

Figure 4: Influence of reactor temperature on product distribution Phenol was the main product over the whole temperature range, butthe formation of byproducts became important above 400~ (Fig. 4). At this temperature the expected exponential increase of phenol yield versus temperature did not occur because consecutive reactions of phenol lead to the formation of di-hydroxy-benzenes (catechol, resorcinol) and benzoquinone as well as by total oxidation. Surprisingly, no hydroquinone was found, suggesting a fast consecutive oxidation of the para-isomer to benzoquinone [5-7]. Additionally, above 400~ the conversion of nitrous oxide increased because secondary products generated by these further reactions need higher stoichiometric amounts of nitrous oxide (e.g. total oxidation requires 15 molecules of nitrous oxide) [6]. The increase in undesired products, especially the steep rise of total oxidation, confined the maximum operating temperature to 400~ 3.3 Effect of feed concentrations An increase of nitrous oxide feed concentration led to an increase inreaction rate and benzene conversion. On the other hand benzene conversion decreased with increasing benzene feed concentration (Table 2). Phenol yield basically changed in the same way as benzene conversion in all cases. Upon variation of the feed concentrations, the highest phenol production was obtained at 26% nitrous oxide and 12.5% benzene at T=400~ W/F=92 gmin/mol and a catalyst time on stream of 40 minutes [5-7].

851

Table2." Benzene conversion as a function of nitrous oxide and benzene partial pressures (temperature=400 ~ W/F = 92 ~.min/mol, catalyst time-on-stream=40 minutes) feed concentration of nitrous oxide [%] 8.3 16.7 26.0

feed concentration of benzene [%] 4.2 36.1 % 19.3 % 6.6 %

2.1 4.2 12.5

39.8 % 20.7 % 8.4 %

42.6 % 23.8 % 10.7 %

46.8 % 29.6 % 11.7 %

Rather than the relative magnitudes of conversion and yield, the absolute values prove more helpful in understanding the experimental observations. Both the absolute amount of converted benzene and formed phenol increased with increasing nitrous oxide feed. In contrast to decreasing relative benzene conversion in Table 2 the amount of absolute converted benzene increased with increasing benzene feed concentration. Phenol formation displayed an even stronger dependence on the benzene feed concentration [6]. [] 4.2% benzene

912.5% benzene

.2.1% benzene

100 90 ,,...-

H-Ga-ZSM5 W/F: 92 g min/mol T= 400~

80

r o

70

..r

60

o

50

._~ >

4o

"6 "~

3o 20 lO I

I

I

I

I

10

15

20

25

30

Feedconcentration of nitrous oxide [%]

Figure 5 Selectivity to phenol as a function of feed concentration of nitrous oxide at different benzene feed concentration These effects were reflected by the change of selectivity to phenol upon a variation of the feed concentrations in Figure 5: Benzene selectivity to phenol shows the strong influence of benzene feed concentration on the product distribution. Upon increasing the benzene feed concentration from 2.1% to 12.5%, the selectivity to phenol increased from about 45% to nearly 95%. The influence of nitrous oxide feed concentration on selectivity to phenol is less pronounced. An increase of the nitrous oxide partial pressure led to a decreased selectivity to phenol. In the same way selectivity to benzoquinone increased. But due to a stoichiometric consumption of three molecules nitrous oxide per molecule of benzoquinone it is obvious that its selectivity is even more strongly dependent on the nitrous oxide partial pressure [6]. In summary, increasing feed concentration of nitrous oxide led to an increase of reaction rate with only little loss of selectivity to the desired product phenol. Absolute benzene conversion increased with increasing benzene feed concentration especially at high nitrous oxide

852 concentrations. Selectivity to phenol increased with increasing benzene feed concentrations. In the same way this led to a rise in phenol production exceeding the increase of converted benzene. Upon proper variation of these parameters space-time-yields of more than 1 kg phenol per kg catalyst and per hour may be reached. These observations can be explained by considering adsorption effects. Sorption simulation calculations (Cerius 2) resulted in higher benzene adsorption amounts compared to nitrous oxide. Excess benzene concentration on the catalyst surface especially at low nitrous oxide partial pressure leads to a lower influence of benzene partial pressure variation on the reaction rate compared to nitrous oxide. The increase of phenol selectivity with rising benzene feed concentration on the other hand can be explained by competitive adsorption between benzene and phenol. Increasing benzene partial pressure changes the sorption equilibrium, and forces desorption of phenol to play an increasing role. Further hydroxylation and polymerisation reactions are suppressed and selectivity increases. This explanation is supported by adsorption simulation calculations. The effect of benzene concentration on selectivity to benzoquinone is controlled by two competing effects. High benzene feed concentration suppresses benzoquinone formation by forced desorption of phenol, while low benzene feed concentration leads to consecutive reactions of benzoquinone. Both have a comparable influence on phenol selectivity. 4. KINETIC MODELING Based on these reaction engineering results several of kinetic models were developed. The first step was to develop a power law model describing the experimental data. Especially, the activation energy of benzene hydroxylation to phenol was of particular interest. The following reaction scheme was suggested C6H6 + N20 --> C6HsOH + N2

(Eq. 2)

C6HsOH + VN20 N20 --> consecutive products

(Eq. 3)

resulting in a set of differential equations (Eq. 4-6). Because of the measured higher consumption of nitrous oxide for consecutive reactions of phenol, the parameter VN20 was introduced in the modell (Eq. 2, Eq. 5). It is a global factor for NEO-Consumption for all consecutive products and not correlated with the order q refering to xmoin Eq 5. Table 3 summarizes the results of the parameter estimation. All values of the kinetic parameters - reaction orders, rate constants and activation energies - were estimated by nonlinear regression analysis based on numerous experiments. -EA,I

O~C6H 6

&

(~2 O

-

--

O~C6H5OH

&

-

kol

kol

--

-

--

kol

9e

RT

. X m C6H6

(Eq. 4)

n

"

XN20

-EA,1 e RT 9

-EA,2 rn XC6H6 9

n XN20 9 -- VN20

ko2 9

-EA,1 9e

RT

e9

RT

. Xp C6H5OH

Xl~12 9 O

(Eq. 5)

-EA,2 . Xm

C6H6

N20 -- ko2

. X n

e9

RT

. X p

C6HsOH

Xl~12 9 O

(Eq. 6)

Using experimental data at the beginning of the reaction (time-on-stream = 15 min) it can be assumed that deactivation is negligible.

853

This power law model is able to describe the experimental results with good significance (standard deviation in Table 3).

Table 3

Valuesof parameters for the power law model

parameter

value

standard deviation

n

0.26

0.03

m

0.3

0.02

kol

0.955 mol/(g min)

0.136 mol/(g min)

EA,1

40.6 kJ/mol

3.2 kJ/mol

P

0.18

0.08

q

0.43

0.07

k02

55 mol/(g min)

32 mol/(g min)

EA,2

65.4 kJ/mol

7.3 kJ/mol

VN20

3.03

0.36

9benzene

+10%

m "0

o

9

0.25

o

E

0.20

v

t-

s9 ~6 o 0

E

9phenol

a nitrous oxide

0.03

0.30

0.02

t-

0.15

O ,m

~6

0.10

0.01

o

o.o5

E

0.00 0.00

0.10

0.20

mole fraction (experiment)

0.30

0.00 0.00

0.01

0.02

mole fraction (experiment)

0.03

Figure 6 Power law modell" Parity plots of calculated and observed mole fractions Except for a few points, the parity plot of benzene and nitrous oxide mole fractions displays a satisfactory agreement with a maximum deviation of + 10% (Fig. 6 left). The higher deviation between calculated and measured phenol values (Fig. 6 right) stems from the simplicity of the model for the description of more complex consecutive reactions of phenol. Values of reaction orders (Table 3) and competitive sorption effects mentioned in Section 3.3 could be quantified more satisfactorily using a adsorption kinetic model. The model was derived under the assumptions of constant volume reaction, nitrous oxide reacting from the gas phase (because of the small influence of nitrous oxide feed concentration on phenol selectivity), three times higher phenol sorption constant than benzene sorption constant (Kc6HSOH~3KC6H6, as derived from sorption simulation calculations), but without considering the dependency on temperature (reaction temperature 400~ [5,7]. In addition, the consecutive oxidation of phenol to benzoquinone was regarded as a two step reaction with hydroquinone as intermediate, thus avoiding the the need to postulate formulation of a trimolecular collision. It was further assumed that hydroquinone was

854

immediately oxidized to benzoquinone because no hydroquinone was found in the product stream. When modeling the reaction rate at 400~ total oxidation can be neglected. Similar to the above mentioned power law rate model, only the first data points were used for parameter estimation, assuming no deactivation at this time-on-stream. ~PCsH6

&

= -k 1

aPN20 o~

= -kl

9

K C6H6 Pc6H6 9 1 + K C6H6 p C6H6 9

"1-

PN20 9

(Eq. 7)

3 K C6H6 p C6H5OH 9

KC.H. PC6H. PN20 "

1 + Kc.H.

PC.H6 + 3 KC6He PC6H5OH

(Eq. 8)

3. K C6H6 9PC6H5OH-PN20 - 2k2 aPc6H5OH c~

Pc6H6 -I" 3

K C6H6 " P C 6 H 5 O H

9

KC6H6 PC6H6 PN20 9 1 + K 06H6 9Pc6H6 q- 3. K C6H6 9PCsHsOH

= +k 1 -k

1 + K C6H6

2

3. K C6H6 .PC6HsOH

9

9

(Eq. 9)

PN20

1 + K C6H6 "Pc6H6 -'l- 3. K C6H6 "PC6H5OH

With the values from parameter estimation given in Table 4 a satisfactory description of the experimental data with the kinetic model at a temperature of 400~ can be achieved (Fig.7), though the number of parameters is lower compared to the power law model. This demonstrates the particular effect of the competetive adsorption of the aromatic components on the reaction resulting in increasing selectivity at high benzene feed concentrations.

Table 4: Results of parameter estimation at 400~ for the adsorption model parameter

value

standard deviation

kl [mol/(g min)]

1.7E-03

9.1E'05

k2 [mol/(g min)]

1.3E-03

1.5E-04

KC6H6 [1/bar]

176

66

KC6H5OH [1/bar]

528 9phenol

9 benzene & nitrous oxide 30 1 m ll) "O O

+

"o o vE

25

E

3.0 2.5

20

v

~

15

(/) ,.~ ~.__,

I~.

10

a. (~

1.0

"t::

5

1:::

0.5

.~_

,0 ,._,

0 0

10

20

30

partial pressure (e~enment) [kPa]

2.0

0.0 0.0

1.0

2.0

3.0

partial pressure (expedment) [kPa] Figure 7 Adsorption model: Parity plots of calculated and observed mole fractions

855 Especially, the satisfactory prediction of phenol values confirms that the adsorption model provides a better description of the kinetics than the power law model. (Fig. 7, right). 5. REDUCING DEACTIVATION For an industrial application it is necessary to achieve stable space time yields of about 1 kg phenol per hour per kg of active catalyst material [6]. Suppressing the strong deactivation, by adding of oxygen to the feed is not successful because the phenol yield decreases with increasing oxygen concentration [5]. To reduce coke formation a modification of the catalyst properties is necessary. It is assumed that stronger acid sites are responsible for deactivation [18]. A selective poisoning of the strongest acid sites by silylation should hence improve long term stability with little loss of activity towards benzene hydroxylation [6]. Figure 8 shows the improvement in long term stability for H-A1-ZSM5 zeolite with a SiO2 to A1203 ratio of 100 at 425~ As expected the initial activity of the silylated zeolite was reduced compared with the untreated catalyst but phenol yield was higher after 50 minutes time-on-stream. Optimizing this pre-treatment procedure could probably lead to a catalyst with strongly reduced deactivation. [] non-silylated

9silylated

20 18~

14

0 t'-

12

"o9 ._~ >..

8 6

W/F= 92 gmin/mol

06H6; 4.2 % N20:26 %

T= 425 ~

(i.) 10 t"

4

2+ ot 0

t

t

t

I

I

t

I

50

1O0

150

200

250

300

350

400

time-on-stream [rain] Figure 8 Effect of silylation on long term stability of H-A1-ZSM5 (SIO2/A1203 ratio=100) 7. CONCLUSIONS Our investigations showed that H-Ga-ZSM5 is an active catalyst for benzene hydroxylation with nitrous oxide. The high activity observed does not support the hypothesis that the presence of iron, possibly included as impurities during synthesis of zeolite, is necessary for benzene hydroxylation. On the other hand the acid strength of Br~nsted acid sites decreases in the sequence H-A1-ZSM5>H-Ga-ZSM5>H-Fe-ZSM5 and this is in contrast to the activity gain by exchanging aluminum for gallium or iron [5,6]. So in mechanistic questions no conclusive answer is possible. Further investigations to explain the nature of the active sites in the complex zeolitic system are necessary.

856 In contrast to the indistinct nature of active sites, our investigations gave detailed information about the reaction pattern. The influence of reaction temperature on product distributions and phenol yield could be quantified. The appearance of total oxidation products confined the maximum operating temperature to 400~ in the case of H-Ga-ZSM5 catalysts. It seems to be advantageous to produce phenol working at high feed concentrations of nitrous oxide and benzene. High nitrous oxide partial pressure leads to high reaction rate with only little loss in selectivity towards the main product phenol. An increase in benzene partial pressure leads to an increase in both reaction rate and selectivity to phenol. High selectivity at high reaction rate improves the space time yield important for industrial application. In [6] promising space time yields of 1 kg phenol per hour per kg of active catalyst material are reported. Kinetic modeling resulted an initial satisfactory approache to describe the experimental data. While with a power law model, activation energy could be estimated, with the adsorption model the competitive sorption effects at various feed concentrations could be described. The main problems of this new route to phenol are the high costs of nitrous oxide and the strong deactivation of zeolite catalysts due to coke formation. Nevertheless less expensive nitrous oxide may be recovered from the waste gas stream of adipic acid plants, attempts to reduce deactivation by a modification of the zeolite catalysts appear promising. REFERENCES

[1] R.Burch, C.Howitt, Appl.Catal.A, 103 (1993) 135 [2] E.Suzuki, K.Nakashiro, Y.Ono, Chem.Lett.,6(1988)953 [3] M.Gubelmann, P.Tirel, J.Popa, 9th International Zeolite Conference, Montreal, July 1992. [4] G.Panov, A.Kharitonov, V.Sobolev, Appl.Catal.A, 98 (1993) 1 [5] M.H/ifele, A.Reitzmann, D.Roppelt and G.Emig, Erdrl Erdgas Kohle, 12 (1996) 512. [6] M.H/ffele, A.Reitzmann, D.Roppelt and G.Emig, Appl.Catal.A, accepted for publication (1996). [7] A.Reitzmann, M.H~ifele and G.Emig, Trends in Chemical Engineering, Research Trends, Council of Sci. Res. Trivandrum, Vol.3 (1996) 63. [8] M.Iwamoto, J.Hirata, K.Matzukami and S.Kagawa, J.Phys.Chem., 87 (1983) 903. [9] M.Gubelmann and P.-J.Tirel, EP 341165 (1989). [10]R.Burch and C.Howitt, Appl.Catal.A, 86 (1992) 139. [ 11] G.I.Panov, G.A. Sheveleva, A.S.Kharitonov, V.N.Romannikov and L.A.Vostrikova, Appl.Catal. A,82(1992)31. [ 12] M.Gubelmann, J.-M.Popa and P.-J.Tirel, EP 406050 (1990). [ 13] G.I.Panov, A.S.Kharitonov and G.A.Sheveleva, WO 95/27691 (1994). [14]R.Burch, and C.Howitt, Appl.Catal.A, 106 (1993) 167. [15] A.S.Kharitonov, G.A.Sheveleva, G.I.Panov, V.I.Sobolev, Y.A.Paukshtis and V.N.Romannikov, Appl.Catal.A, 98 (1993) 33. [16] V.I.Sobolev, K.A.Dubkov, E.A.Paukshtis, L.V.Pirutko, M.A.Rodkin, A.S.Kharitonov and G.I.Panov, Appl.Catal.A, 141 (1996) 185. [17] Petrochemical News, 1996, 35 (53), 2. [ 18] R.Barrer, R.Jenkins, G.Peeters, Amer.Chem.Soc., Molecular Sieves II, 258-270,1977

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

857

Direct Hydroxylation of Benzene to Phenol by Nitrous Oxide Anthony K. Uriarte a, Mikhail A. Rodkin a. Michael J. Gross a, Aleksandr S. Kharitonov b and Gennady I. Panov b a Fibers Strategic Business Unit, Monsanto, P.O. Box 97, Gonzalez, FI 32560, USA b Boreskov Institute of Catalysis, Novosibirsk, 630090, Russia

1. ADIPIC ACID MANUFACTURE AND NITROUS OXIDE BYPRODUCT About 5 billion p o u n d s per year of adipic acid are m a n u f a c t u r e d worldwide by the nitric acid oxidation of cyclohexanone a n d / o r cyclohexanol (KA). The KA to adipic acid yields are near 94% of theory. Glutaric and succinic acids are the major byproducts and account for most of the yield loss. Monsanto and some other adipic acid producers recover or upgrade these to salable byproducts resulting in an overall KA utilization efficiency that approaches 99%. However, the nitric acid efficiency is lower because approximately 1 mole N20 is produced per mole of adipic acid, in addition to the easily recyclable NOx that is generated as a result of nitric acid reduction. In the early 1990's it was reported that N 2 0 emissions from adipic acid producers could contribute to atmospheric ozone depletion and global warming [1]. It was estimated that adipic acid production may account for up to 10% of the a n n u a l increase in the atmospheric N20. This report sparked an a b a t e m e n t initiative among the major adipic acid producers. Successes have been a n n o u n c e d and implementations are scheduled for 1996-98 [2]. However, Monsanto was already practicing complete a b a t e m e n t by a thermal reduction process and elected to p u r s u e a more desirable path of value-added utilization. Two general areas of utilization were considered: I) oxidation of N20 to NO and s u b s e q u e n t conversion to nitric acid; and 2) use of N20 as a selective oxidant. The latter had the potential of satisfying the criterion of value addition. There were several reports on the selective hydroxylation of aromatics. Because of the economy of scale, the full use of the available N20

858 would be an important factor in the consideration of the potential options. Monsanto's adipic acid production at Pensacola, Florida is over 600 million p o u n d s per year which equates to almost 200 million p o u n d s of N20. For the hydroxylation of aromatics at stoichiometry, this a m o u n t is equivalent to 400 million p o u n d s of phenol. Based upon a projected a n n u a l growth in the phenol m a r k e t of 3-5%, paper studies were initiated to evaluate the use of N20 to produce phenol. 2. HYDROXYLATION OF BENZENE AND NEW ADIPIC ACID P R O C E S S CONCEPT Reports on use of N20 for hydroxylation of benzene to phenol appeared as early as 1983 [3] - M. Iwamoto used V2Os/SiO2 as a catalyst and at 550oC achieved 10% conversion and 70% selectivity of benzene hydroxylation. The other reports that followed concentrated on usage of ZSM-5 type catalyst for this transformation [4-7]; see also a comprehensive review on the subject [8]. A brief s u m m a r y of catalysts used and process p a r a m e t e r s is given in Table 1. Fast catalyst deactivation was reported for most cases and the values of productivity usually refer to initial stages of the reaction. Table 1 Catalyst

V205/SIO2 H-ZSM-5 H-ZSM-5 H-ZSM,5 Fe-ZSM-5

SiO2/A1203 Selectivityof benzene conversion to phenol, % 70 85 NA >90 95 33 98 100 99

Productivity,mmole phenol per gram of catalyst per hour 1.0 0.3 3.2 1.8 3.0

Ref.

[3] [4] [5] [6] [7]

In summary, benzene can be reacted with nitrous oxide in the vapor phase at elevated temperatures over ZSM-5 or similar catalysts to give phenol and nitrogen (eq. (I)).

300-500oc + N20

7_,SM-5

~~,/OH + N2

(1)

859 The reaction has very high selectivity of benzene conversion to phenol (>99%). A further step was taken to incorporate the phenol scheme into an overall adipic acid process. Eq. (2) summarizes one such possibility.

%,.iC00Ii + N20

A--

2}Iz

i (2}

The process would use N20 to hydroxylate benzene to phenol. The phenol would be hydrogenated to cyclohexanone using available technology. The final step is the currently practiced nitric acid oxidation of cyclohexanol and cyclohexanone, which r e t u m s N20 for use in the front end of the process. The stoichiometric balance is close; however, either some additional on-purpose N20 or KA would likely be required for a stand alone plant. The successful commercialization of the overall process concept depended on the viability of the first step which is a breakthrough technology. The data reported in the literature showed high selectivity of benzene conversion to phenol and good productivity. However, the catalyst coked quickly - in most reported cases the catalyst lost its activity in a matter of a few hours. Another problem of the reported chemistry is the low N20-to-phenol selectivity. In fact, the stoichiometry of benzene oxidation to CO2 by N20 implies that 1% of benzene selectivity loss to deep oxidation is accompanied by 15% selectivity loss in N20 conversion. Considering that the supply of nitrous oxide is limited, the efficiency of its utilization is very important for the commercial operation. To develop a viable commercial process, Monsanto and The Boreskov Institute of Catalysis (BIC) formed a joint R&D team in 1994. Understanding of the fundamental aspects of the reaction played an important role in bringing this concept from a lab curiosity to what, we hope, will become an industry standard in short time.

860 3. REACTION MECHANISM The heart of the benzene hydroxylation is the catalyst active site known as alpha-site (a term coined by BIC). The proposed mechanism is outlined is equations (3-5) below:

N20 + ( )a

~

~~,,~OH +

~ O H

(3)

N2 + (O)t:t

a

(o)a

a

._

~~,~O]H

+ ( )a

(4)

{5}

At the initial stage nitrous oxide is believed to decompose on an alpha-site, loading it with a unique form of oxygen, called alpha-oxygen, and releasing dinitrogen. This active alpha-oxygen then reacts with a benzene molecule inserting oxygen into the C-H bond and yielding adsorbed phenol. Desorption of the product is the final stage, which frees the active site for further reaction. How real is the proposed mechanism? It turns out that its stages can be modeled separately. In the absence of benzene, decomposition of nitrous oxide in a closed system at temperatures below 300oC leads only to evolution of nitrogen - all of the released oxygen under these conditions is left on the catalyst in the form of alpha-oxygen. This process was used as a basis for one of the procedures developed for measurement of alpha-site concentration [9]. When a catalyst loaded with alpha-oxygen is isolated from nitrous oxide, cooled and reacted with benzene vapors at or below room temperature, phenol is extracted from the catalyst as the only product, thus supporting a model for the second and third steps of the proposed mechanism [10].

861 4. P R O C E S S CHEMISTRY AND E N G I N E E R I N G

4.1 Catalyst Deactivation Deactivation of the catalyst and benzene selectivity loss to coking were major concerns at the initial stages of the program. Dramatic improvements have been achieved in this a r e a - see Fig. 1

t

Productivity, mmoles of phenol

Time

Fig. 1. Change in catalyst deactivation profile The activity of the first generation catalysts dropped to one-half of their initial value within 3-5 hours. With catalyst and process improvements the catalyst half life has been increased to 3-4 days. Fig. 1 also shows that we were able to change the original exponential decay into linear deactivation with the new catalyst system and process design.

4.2 Catalyst Productivity In the early experiments, the system showed a good productivity of ca. 1 mmole of phenol per gram of catalyst per hour. Further process development led to more t h a n 10-fold increase in productivity. In the current design base case we a s s u m e the average productivity of the catalyst over 48 hours to be above 4 mmole of phenol per gram of catalyst per hour, which is at the top end of the best industrial catalysts.

862 4.3 C a t a l y s t R e g e n e r a t i o n

Once the catalyst activity goes below the acceptable level, the catalyst can be regenerated by passing oxygen-containing gas through the catalyst bed at elevated temperatures, which completely restores the catalyst activity. No performance deterioration has been observed after multiple reactionregeneration cycles (over 100) and longer term effects are under investigation. 4.4 R e a c t o r D e s i g n

Based on the early performance data (rapid catalyst coking and high reaction exothermicity), the initial choice for the process design was a fluidized bed reactor. However, further studies revealed that the reaction selectivity is remarkably insensitive to the temperature rise present in an adiabatic reactor. This observation, along with a substantial improvement in catalyst stability and the need for quick scale-up, led to the selection of a simple adiabatic plug flow reactor design. The pilot plant, Fig 2, has been in operation since May 1996. The unit design includes continuous recycle of the vapor and liquid streams with a full complement of on-line analyzers. (column bypass), _ .......

-~--~I!

!

Air

9 ~.-

Reactor [

Vent

I

:~..~ ....

Nitrogen'-" Feed Phenol Storage

Benzene Storage

Compressor , (Vent Gas Recycle! |

IIR~alyz~l :, |

,

N20 Storage Preheate

____~4 ..........

|

!

-~ ...... ~ ...........

[02 Analyzer I : .,~ . . . . . . . . . . .~

Flg. 2. Simplified scheme of the benzene-to-phenol Pilot Plant in Pensacola, Florida

863

4.5 Overall P r o c e s s P e r f o r m a n c e Table 2 shows typical performance p a r a m e t e r s achieved in the pilot plant d e m o n s t r a t i o n runs. Table 2 Typical performance d a t a Performance p a r a m e t e r s Reaction Temperature, ~ C

400-450

Contact Time, seconds

1-2

Benzene to Phenol, mol %

97-98

Benzene to COx, mol %

0.2-0.3

Benzene to Diols, mol %

1

Nitrous oxide to Phenol, mol %

85

Productivity, mmole P h e n o l / g c a t a l y s t / h r

4

4.6 P r o c e s s Safety The described reaction Offers o u t s t a n d i n g safety features for process design. It is a g a s - p h a s e reaction with very short residence time, therefore there is a minimal inventory of flammable material. Operating conditions have been defined t h a t e n s u r e the whole process is non-flammable throughout: the reaction step, the separation systems a n d the recycling loops. Another safety feature of the proposed technology is the absence of any highly-reactive intermediates. All active catalyst surface species are immediately c o n s u m e d a n d their concentration is minuscule.

4.7 E n v i r o n m e n t a l F e a t u r e s The process h a s a very high selectivity toward the target phenol. Unreacted benzene is completely recycled. The separation is a simple distillation - there is minimal a q u e o u s waste, no inorganic salts a n d some of the by-products can be isolated a n d sold. The total organic waste is expected to be less t h a n 2% of the phenol m a n u f a c t u r e . And last, b u t not least - the process u s e s waste nitrous oxide, a b a t e m e n t of which currently c o n s u m e s n a t u r a l gas a n d emits m u c h more CO2 t h a n is expected to be emitted by the benzene to phenol process.

864

REFERENCES 1. M.H.Thiemens, W.C. Trogler, Science, 251 (1991) 932. 2. R.A. Reimer, C.S. Slaten, M.Seapan, M.W.Lower, P.E.Tomlinson, EnvironmentaI Progress, 13(2)(1994) 134. 3. M. Iwamoto, J.I. Hirata, K. Matsukami, S. Kagawa, J. Phys. Chem., 87 (I 983) 903. E. Suzuki, K. Nakashiro, Y. Ono, Chem. Lett., (1988) 953. 5. G u b e l m a n n , P.J. Tlrel, Eur. Pat. Appl. EP 341165 A1 8 Nov 1989, 4 pp (US Patent No. 5,001,280); M. Gubelmann, J.M. Popa, P.J. Tirel, Eur. Pat. Appl. EP 4 0 6 0 5 0 A2 2 J a n 1991, 9 pp (US Patent No. 5,055,623). 6. R. Burch, C. Howitt, Appl. Catal., A, 86 (1992) 139. 7. A. S. Kharitonov, T.N. Aleksandrova, L.A. Vostrikova, K.G. lone, G.I. Panov, USSR Pat. No. 4 4 4 5 6 4 6 (1989); A. S. Kharitonov, G. I. Panov, K. G. Ione, V. N. Romannikov, G. A. Sheveleva, L. A. Vostrikova, V. I. Sobolev, U.S. US Patent No. 5110995 A 5 May 1992, 8 pp. 8. G. I. Panov, A. S. Kharitonov, V. I. Sobolev, Appl. Catal., A, 98 (1993) 1. 9. G. I. Panov, V. I. Sobolev, A. S. Kharitonov, J. Mol. Cata/., 61 (1990) 85. 10. V.I. Sobolev, A.S. Kharitonov, Ye. A. Paukshtis, G.I. Panov, J. Mo/. Cata/., 84 (I 993) 117. .

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

865

R a p i d c a t a l y t i c o x y g e n a t i o n of h y d r o c a r b o n s w i t h perhalogenated ruthenium porphyrin complexes

John T. Groves*, Kirill V. Shalyaev, Marcella Bonchio and Tommaso Carofiglio Department of Chemistry, Princeton University, Princeton, NJ 08544, USA

Abstract. Perhalogenated ruthenium porphyrins were found to be efficient catalysts for the oxygenation of hydrocarbons including secondary alkanes and benzene in the presence of 2,6-dichloropyridine N-oxide under mild conditions in aprotic media. Up to 15,000 turnovers and rates of 800 TO/min were obtained. A mechanism where Ru(III) - Ru(V) intermediates play an important role is proposed and discussed. The search for new methods for the catalytic oxygenation of hydrocarbons is one of the most important directions of modem chemistry [1]. Among the metalloporphyrinmediated oxidations, ruthenium catalysts display remarkable activity for aerobic oxidations [2] (Fig. 1) and promising reactivity with N20 [3]. O

O O

O

o

o

[ ~ )

Figure 1. Catalytic cycle for olefin epoxidation with dioxygen using ruthenium porphyrins.

~-~r 2

866 Highly efficient oxygenation reactions with ruthenium porphyrin complexes and aromatic N-oxides in the presence of strong mineral acids have been described by Hirobe et al. [4]. We have recently reported that electron deficient perhalogenated ruthenium porphyrins catalyze the oxygenation of a variety of even unreactive substrates under mild conditions (40- 65~ in the presence of 2,6-dichloropyridine N-oxide in aprotic media [5]. Unusually high rates and turnover numbers (TO) were obtained. Carbonyl (5,10,15,20-tetrapentafluorophenylporphyrinato) ruthenium(H) Run(TPFPP)(CO) has shown unusually high activity with 2,6-dichloropyridine N-oxide as the oxygen donor (Table 1). Here we describe further studies of the mechanism of this remarkable process. CO

~

oo

PFPP

,,

~

F Adamantane and cis-decalin were hydroxylated with high selectivity, complete stereoretention, extraordinarily high rates (up to 800 turnovers/min), and high efficiency (up to 15,000 turnovers, Table 1, entries 1-3). Similar conversions were obtained when RuVI(TPFPP)(O)2 and RuVI(TPFPBrsP)(O)2 were used as catalysts. Oxygenation of less reactive substrates such as benzene and cyclohexane proceeded with lower but still significant turnover numbers (100-3,000, Table 1, entries 4-6). Tertiary versus secondary selectivity in adamantane oxidation adjusted for the number of carbons-hydrogen bonds of each kind was above 210. No rearrangement products were detected in cis-decalin hydroxylation. The kinetics of product evolution in a typical reaction of adamantane hydroxylation showed an initial induction period followed by a fast, apparently zero-order phase with the maximum rate and highest efficiencies (Fig. 2). Deviation from linear behavior took place only after 90% oxygen donor and 80% of the substrate had been consumed. When II RuVI(TPFPP)(O)2, prepared by reaction of Ru (TPFPP)(CO) with 3-chloroperbenzoic acid was used as the catalyst, no induction time was detected and zero-order kinetics were observed as well. The well defined and characteristic UV-vis spectra of metalloporphyrins provide an invaluable tool for the mechanistic studies. Thus, monitoring the state of the metalloporphyrin catalysts during the course of both model reactions by UV-vis spectroscopy revealed that the initial form of the catalyst remained the predominant one throughout the oxidation, i.e. in the Run(TPFPP)(CO) catalyzed reaction c.a. 80% of the porphyrin catalyst existed as Run(TPFPP)(CO) and in RVI(TPFPP)(O)2 catalyzed reaction more than 90% of

867 the catalyst was still in the form of RuVI(TPFPP)(O)2 despite the high turnover numbers reached (-- 400 TO). The fact that Run(TPFPP)(CO) demonstrates similar and even higher maximum turnover rate of 4.9 TO/min in adamantane hydroxylation versus 4.0 TO/min for RuVI(TPFPP)(O)2 under the same conditions indicates that an active catalyst species other than RuVI(TPFPP)(O)2 is involved in the fast catalytic hydroxylation. Table 1. Hydrocarbon oxidations a catalyzed by [Run(TPFPP)(CO)]. Time (min.)

Product (% conv.) b

adamantane

20

1-adamantanol (76.2) adamantane- 1,3-diol (7.3)

91

72

2e

adamantane

120

1-adamantanol (61.0) adamantane-l,3-diol (6.0)

97

800

3

cis-decalin

25

(Z)-9-decalol (79.6) (Z)-decal-9,10-diol (4.2)

90

64

4

trans-decalin

60

(E)-9-decalol (25.8) secondary alcohols (4.3) ketones (13.9)

70

4.4

5e

cyclohexane

180

cyclohexanol (1.6) cyclohexanone (6.7)

95

22

6f

benzene

12h

1,4-benzoquinone (13.3)

40

7

1-octene

60

1,2-epoxyoctane (96)

96

11 (36)g

8

1-octene/adamantane

60

1,2-epoxyoctane (54) 1-adamantanol (28)

90

9.5 4.8

9h

cyclohexene

320

cyclohexene oxide (18.2)

-

0.38

#

Substrate

1

Yield c Max. rate d (%) (TO/min)

ii a [substrate] - [pyC12NO] - 0.02 M, [Ru (TPFPP)(CO)] = 50 I~Jl. All reactions in CH2C12 at 65 ~ in sealed containers, b% conversion based on substrate consumed. Products were identified by GC-MS and compared to authentic samples, c % yield based on pyC12NO consumed. d maximum oxidation rate measured as the slope of the zero order phase of the kinetic plot. e [substrate] = [pyC1ENO] = 0.2 M, [Run(TPFPP)(CO)] = 10 ~ M . f [benzene] = 2 M, [pyC1ENO] = O.i02 M, [RulI(TPFPP)(CO)] = 50 ~/I. g [1-octene] = 0.O1 M, [pyC1ENO] = = 0.04 M, [Ru (TPFPP)(CO)] = 50 ~lM, maximum rate = 36 TO/min. h [cyclohexene] = 0.04 M, [pyCl2NO] - 0.02 M, [Run(TPFPP)(CO)] = 50 ~M, reaction was not complete at 320 min.

868

Figure 2. Adamantane Hydroxylation Catalyzed by RuI~(TPFPP)(CO) and RuVI(TPFPP)(O)2, [adamantane] = [pyC1ENO] = 0.02 M, [catalyst] = 50 ~tM, CHEC12, 40~

TO = moles of product/moles of catalyst

400

~

u

v

~

0

300TO CO

200-

100-

0

50

100 Time, min

150

200

Therefore, the classical t r a n s - d i o x o R u ( V I ) - oxoRu(IV) catalytic cycle [2] (Fig. 1) can be ruled out as the primary reaction pathway in case of rapid catalytic oxygenation. The apparent zero-order kinetics observed are consistent with a steady-state catalytic regime accessible from different initial states of ruthenium metalloporphyrin. Indeed, common oxidants, other than aromatic N-oxides, such as iodosylbenzene, magnesium monoperoxyphthalate, Oxone | and tetrabutylammonium periodate produced the t r a n s dioxoRu(VI) species from RuII(TPFPP)(CO) under reaction conditions but were ineffective for the rapid catalysis. A two-electron oxidation of Run(TPFPP)(CO) would produce oxoRu(IV) porphyrin and eventually dioxoRu(VI). What is the alternative pathway for Run(TPFPP)(CO) activation? It is known that ruthenium(II) ~-cation radicals are formed from the corresponding carbonyl compounds by chemical or electrochemical one-electron oxidation [6]. Such species have been shown to undergo intramolecular electron transfer upon axial ligation and removal of CO to give ruthenium(III) porphyrins [6c, 7]. An emerald green solution of Run(TPFPP)(CO) § radical cation with a strong EPR signal (g = 2.00) was quantitatively obtained when Run(TPFPP)(CO) was oxidized with ferric perchlorate in methylene chloride. An EPR signal typical of a ruthenium(HI) species (gll = 2.55, g_l_= 2.05)

869 [5, 6d,e] was detected after the addition of 2,6-1utidine N-oxide to the solution of the radical cation. Photo-stimulation of RuII(TPFPP)(CO) catalyzed reactions with red-orange light (> 560 nm) which shortened the induction period was observed. Interestingly, RuII(TPFPP)(CO) has only residual absorbance in this region of the spectrum but RuH(TPFPP)(CO) +" radical cation shows a strong band at 635 nm. We conclude that the effect of the red-orange light is consistent with photoejection of the carbonyl ligand from the radical-cation to produce a Ru(III) species. We propose that Ru(III) and oxoRu(V) species are the key intermediates in the catalytic cycle of "rapid oxygenation" which can be viewed as a part of the general scheme of diverse oxidative chemistry of ruthenium porphyrins (Fig. 3). The aerobic oxygenation pathway (involving the even oxidation states Ru(II), Ru(IV) and Ru(VI)) we have previously described [2] is shown in the left half of Figure 3. The new fast catalytic process on the right half of the figure reveals chemical interconnectivity between the fast and slow catalytic regimes.

Figure 3. Mechanisms of Ruthenium Porphyrin Oxidation Catalysis

-e-

~

7+"

~

CO

CO

OD 02 D X

SO

OD

so

OD

slow

fast

0

s

~ 1

RDS 0

X

SO O

7+"

S - substrate OD - oxygen donor O Thus, one-electron oxidation of RuVI(TPFPP)(O)2 would give the known dioxoRu(VI) cation radical. Oxygen transfer from this species could enter the fast cycle by producing the proposed transient oxoRu(V) complex [8]. Likewise, a one-electron reduction of dioxoRu(VI) porphyrin would also result in formation of the same reactive species.

870 The competitive oxidation of a 1:1 mixture of adamantane and adamantane-d16 catalyzed by RuVI(TPFPP)(O)2 showed a kinetic isotope effect, kn/kD = 4.8 at 40~ (Fig. 4A). Significantly, the deuterated and undeuterated substrates displayed similar turnover rates (kH/kD= 1.2) for hydroxylation in separate reactions (Fig. 4B). 120

120

A

B

100

100-

80

8060 -

60 TO 40

TO 40-

20

~

d16

II

20-i ~"

I

0

i

10

i

i

I

0

i

i

i

i

i

20 30 40 50 0 10 20 30 40 50 Time, min Time, min Figure 4. (A) Competitive hydroxylation of adamantane/adamantane-d16. [ad.] = [ad.-d16] = 0.01 M, [PyC12NO] = 0.02 M, [RuVI(TPFPP)(O)2] = 50 IxM. (B) Separate hydroxylation [substrate] = 0.02 M, [PyC12NO] = 0.02 M, [RuVX(TPFPP)(O)2] = 50 lxM, CH2C12,40~

m

AH~ = 19 kcal/mol

~

R = 0.9967

9 ~

f

4-

I

n

2

I

2.9

3.0

I

I

3.1 3.2 1/T x 1000, 1/K

I

3.3

3.4

Figure 5. Eyring plot of adamantane hydroxylation catalyzed by RuVI(TPFPP)(O)2.

871 Similar results were obtained with Run(TPFPP)(CO) as the catalyst [5]. A kinetic scheme consistent with this observation is as follows: as long as the oxidant OD is present in excess, the active Ru(III) catalyst will exist as the adduct Ru(III)-OD. The reactive oxoRu(V) would then be formed in the rate determining step, which is independent of the concentration of the oxidant (OD) (Fig. 3). The temperature dependence of adamantane hydroxylation yielded a linear Eyring plot for the observed rate constants determined over the range of 25 - 65~ with an apparent AH-~ = 19 kcal/mol (Fig. 5). Olefins such as 1-octene and cyclohexene demonstrated unusually low turnover rates compared to adamantane and cis-decalin (Table 1, entry 7), although in competitive oxidations 1-octene was twice as reactive as adamantane (Table 1, entry 8). Apparently, the catalyst is inhibited in the presence of olefins. Interestingly, upon a stoichiometric reaction of RuVI(TPFPP)(O)2 with cyclohexene in methylene chloride, rapid transformation of the ruthenium porphyrin into Run(TPFPP)(CO) was observed (Fig. 6). 1.0

o

0.9

co

I< p,pp

0.8 0.7

o

0.6 O

0.5

< 0.4 0.3 0.2 0.1 0.0 450

470

490

510

530

550

570

590

610

630

650

Wavelength, nm

Figure 6. Transformation of RuVI(TPFPP)(O)2 into Run(TPFPP)(CO) in the presence of cyclohexene, [cyclohexene] - 0.04 M, [RuVI(TPFPP)(O)2 ] - 50 I.tM, CH2C12, 40~ 30 min. Trans-dioxoRu(VI) complexes are known to react with olefins according to the classical oxo-transfer mechanism [2] (Fig. 1). The oxoRu(IV) intermediate produced in this process disproportionates readily to give dioxoRu(VI) complex and Ru(II) porphyrin which has strong affinity even towards trace amounts of carbon monoxide. A similar process realized as a side reaction in the "rapid oxygenation" system would constantly and effectively tie up the catalyst in the catalytically inactive form of Run(TPFPP)(CO). Indeed, no noticeable changes had been detected in the UV-vis spectrum of the ruthenium porphyrin during the course of Run(TPFPP)(CO) catalyzed oxidation of cyclohexene.

872

Conclusions High selectivities, rates and yields demonstrated in hydrocarbon oxygenation make the catalytic system of perhalogenated ruthenium porphyrins and 2,6-dichloropyridine N-oxide promising for a number of practical applications. Unusually high potency of the active oxidant observed in the substrate studies and results of the kinetic experiments clearly indicate that an active species other than RuVI(TPFPP)(O)2 is implicated in catalysis. We propose that Ru(III) and oxoRu(V) species are the key intermediates in the catalytic cycle. Relatively low rates of olefin epoxidations are explained by the constant buildup of inactive RuII(TPFPP)(CO) during the reaction. Further studies on the mechanism of this remarkable oxygenation reaction and a search for access to the rapid catalytic cycle from other oxidants are in progress.

Acknowledgements Partial support of this research by the Monsanto corporation and the National Science Foundation for the purchase of an NMR spectrometer are gratefully acknowledged.

References 1. (a) Selective Hydrocarbon Activation: Principle and Progress; Davies, J. A., et al., Eds; VHC: New York, 1994. (b) Activation andfunctionalization of alkanes; Hill, C. L., Ed.; John Wiley & Sons: New York, 1989. (c) Metalloporphyrins in Catalytic Oxidations; Sheldon, R. A., Ed.; Marcell Dekker, Inc.: New York, 1994. 2. (a) Groves, J. T.; Quinn, R. J. Am. Chem. Soc. 1985, 107, 5790-5792. (b) Groves, J. T. and Han, Y. Z. in Cytochrome P450: Structure, Mechanism, and Biochemistry; Ortiz de Montellano, P. R., Ed.; Plenum Press: New York, 1995. 3. Groves, J. T.; Roman, J. S. J. Am. Chem. Soc. 1995, 117, 5594-5595. 4. (a) Ohtake, H.; Higuchi, T.; Hirobe, M. Heterocycles 1995, 40, 867-903. (b) Ohtake, H.; Higuchi, T.; Hirobe, M.J. Am. Chem. Soc. 1992, 114, 10660-10662. 5. Groves, J. T.; Bonchio, M.; Carofiglio, T.; Shalyaev, K. V. J. Am. Chem. Soc., 1996, 118, 8961-8962. 6. (a) Dolphin, D.; James, B. R.; Leung, T. Inorg. Chim. Acta 1983, 79, 25-27. (b) Leung, T.; James, B. R.; Dolphin, D. Inorg. Chim. Acta 1983, 79, 180-181. (c) Barley, M.; Becker, J. Y.; Domazetis, G.; Dolphin, D.; James, B. R. Can. J. Chem. 1983, 61, 2389-2396. (d) James, B. R.; Dolphin, D.; Leung, T. W.; Einstein, F. W.; Willis, A. C. Can. J. Chem. 1984, 62, 1238-1245. (e) James, B. R.; Mikkelsen, S. R.; Leung, T. W.; Williams, G. M., Wong, R. lnorg. Chim. Acta (B) 1984, 85, 209-213. (f) For a review see James, B. R. in Fundamentals of Research in Homogeneous Catalysis, Shilov, A. E. Ed. Gordon Breach, NY, 1986, 309-324. 7. Barley, M. H.; Dolphin, D.; James, B. R. J. Chem. Soc., Chem. Commun. 1984, 1499-1500. 8. (a) Che, C. M.; Ho, C.; Lau, T. C. J. Chem. Soc. Dalton. Trans. 1991, 1259-1263 and references therein. For leading references on non-porphyrin ruthenium oxidation catalysts see also: (b) Dobson, J. C.; Helms, J. H.; Doppelt, P.; Sullivan, P.; Hatfield, W.; Meyer, T. J. Inorg. Chem. 1989, 28, 2200-2204.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

873

Ethanol Oxidation Using Ozone over Supported Manganese Oxide Catalysts: A n In Situ Laser R a m a n Study Wei Li*, and S. Ted Oyamat, Laboratory for Environmental Catalysis and Materials, Departments of Chemical Engineering and Chemistry, Vir~nia Polytechnic Institute and State University, Blacksburg, VA 24061-0211 Ethanol oxidation using ozone was investigated over ahmina and silica supported manganese oxide catalysts at temperatures t~om 300 to 550 K by in situ laser Raman spectroscopy coupled with reactivity tests. Ethanol was found to react with ozone at lower temperatures than with oxygen, and also with a lower activation energy. This is in agreement with the stronger oxidizing ability of ozone compared to oxygen. The oxidation reactivity was found to be closely related to that of ozone decomposition, suggesting an important role of ozone decomposition in the reaction mechanism In situ laser Raman spectroscopic studies showed the existence of adsorbed ethoxide species on the catalyst surface under reaction conditions, however, at a much lower concentration than when oxygen alone was used as the oxidant. Transient experiments provided direct evidence that surface peroxide (an adsorbed species due to ozone) and surface ethoxide (an adsorbed species due to ethanol) reacted with each other on the catalyst surface.

1. INTRODUCTION Ozone has been attracting increasing attention recently as an alternative oxidant in the oxidation of volatile organic compounds (VOCs) due to its strong oxidizing ability and hence lower reaction temperatures [1-7]. Ozone was generally found to be effective at enhancing the conversion of VOCs, especially at low temperatures. The kinetics of complete oxidation of benzene by ozone on MnO2 was investigated by Naydenov and Mehandjiev [7]. They found that the activation energy for benzene oxidation with ozone (30 kJ mol"1) was much lower than that with oxygen (88 kJ moll), but was similar to that of ozone decomposition (32 kJ moll). It was concluded that the rate determining step for benzene oxidation by ozone was ozone decomposition. Klimova et al. [4] studied the oxidation of lower aliphatic alcohols by ozone over silica and ahnnina. It was found that the main products of the oxidation reactions were: acetaldehyde and carbon dioxide from ethanol, propionaldehyde and carbon dioxide from n-propanol, and acetone l~om isopropanol. It was determined that for the oxidation of 1 mole of the alcohols about 1 mole of ozone or less was consumed. By varying the reactor vohnne while keeping a constant surface area of the catalyst it was demonstrated that the reaction proceeded primarily on the catalyst surface. Spectroscopic

Department of Chemical Engineering $Departments of Chemical Engineeringand Chemistry

874 study of ozone oxidation reactions has been rare (8), and no work has been done under reaction conditions. This paper reports a comparative study of ethanol oxidation reaction with an ozone/oxygen mixture or oxygen alone over supported manganese oxide catalysts using/n situ laser Raman spectroscopy coupled with reactivity measurements. Alumina and silica supported manganese oxide catalysts were chosen because manganese oxide is an excellent catalyst for complete oxidation, while the alumina supported catalyst has a si~ificantly higher activity for ozone decomposition than the silica supported sample. 2. EXPERIMENTAL 2.1. Catalyst Preparation and Characterization

Ahtmina and silica supported MnO2 catalysts were prepared by incipient wetness impregnation of supports (Degussa, Aluminoxide C and Cabosil L-90) using manganese acetate (Aldrich) as the precursor and were calcined in air at 773 K for 3 h prior to use The crystal phases of MnO2 in the catalysts were identified using X-ray diffraction (XRD) and their surface areas were determined by N2 physisorption using the BET method. Ozone Generator 02~

~ Purifier ~

~

Check Valve

Vent MFC

He ~ Purifier MFC

ValveCheck

Carrier GaSGas Chromatography Sample ]Check Valve

r'-~-~l]o'ns]~ [aere .......... ' , ! Porap.ak .Q.S E Vent~9. . . . . . . . . . . . . . . . . . . . . . . . "

Ozone Analyzer Motor Ethanol (Syringe Pump) Heating Wire

Collecting Notch Focusing Lens Filter Lens

Temperature Ar§ Ion Laser Controller (514.5 nm)

Figure 1. Schematic of the in situ laser Raman system

CCD Detector

875

2.2. In Situ Laser Raman Spectroscopy The in situ laser Raman spectroscopic studies were carded out with a high throughput spectrometer using a Raman sample cell (Fig. 1), which allowed the spectrum to be acquired under reaction conditions. The catalysts (200 mg) were pressed into thin wafers 15 mm in diameter and about 1 mm in thickness, and were held on a ceramic rod by a stainless steel cap The rod was spun at 1800 rpm to avoid local overheating by the laser. The temperature was controlled by a programmable temperature controller (Omega, CN2010), and was measured by a thermocouple placed 3 mm away from the sample The laser Raman spectrometer was equipped with an Ar ion laser (Lexel, Model 95, wavelength = 514.5 nm) as the exciting source, a single-stage monochromator (Spex, 500 M) and a charge-coupled device (CCD) detector (Spex, Spectrum One). A key feature of the Raman system was a holographic notch filter (Kaiser, Super-Notch Plus), which effectively rejected Rayleigh scattering, while allowing > 80% of the Raman signal to pass through. Three sets of reactivity tests were performed on the catalysts: ethanol oxidation using an ozone/oxygen mixture, using oxygen alone, and ozone decomposition. The samples were pretreated in situ at 773 K in oxygen for 2 h before each set of measurements. The reaction feed for oxidation using an ozone/oxygen mixture contained 7.8 mol% oxygen, 0.16 mol% ozone, 0.8 tool% ethanol with helium as the balance gas, and the total flow rate was 110 cm3 rain"~ (82 ~ o l s'~). For the oxidation using oxygen alone, the feed composition was essentially the same except that no ozone was used and the total flow rate was kept at 110 cm3 rain"~ by increasing the flow rate of oxygen. For ozone decomposition, no ethanol was injected into the stream and the total flow rate was kept at 110 cm3 rain ~ by increasing the flow rate of hefiun~ Some bypassing of the gas around the sample likely occurred, but the rate data should be accurate for conversions of 10% or less (differential conditions). Ozone was produced by passing oxygen (Air Products, Extra Dry, > 99.6%) through a corona-discharge ozone generator (OREC, Model V5-0), and the inlet, and outlet ozone concentrations were measured using a UV absorption ozone monitor (Safety Caution" Ozone is highly toxic, hence leak checking and purification of the exhaust stream with an ozone decomposition filter should be carried out.) The reaction products were analyzed by an on-line gas chromatograph equipped with thermal conductivity and flame ionization detectors. Separation of 02, CO, and CO2 was achieved by a Carbosphere column (Alltech) while separation of organic compounds was carried out with a Porapak QS column (Alltech).

3. RESULTS

3.1. Catalyst Characterization Although both alumina and silica supported catalysts had 10 wt% loading of manganese oxide, they exhibited considerably different MnO2 crystallinity (Fig. 2). The alumina supported catalyst essentially only had a well-dispersed manganese oxide phase, the silica supported sample showed strong di~action peaks due to crystalline manganese oxide (13MnO2, JCPDS 24-735). The surface area of the ahxmina and silica supported catalysts was found to be 88 and 75 m2 g-1 respectively.

876

3.2. Ethanol Oxidation Reactivity As expected, ethanol was found to be more reactive with ozone than with oxygen (Fig. 3), especially at low temperatures. The reactivity difference became less si~ificant at higher temperatures, and eventually disappeared above 500 K. MnO2/SiO2 showed fimilar behavior (Fig. 4), however, the reactivity difference between ozone and oxygen was more pronounced in this case, and only disappeared around 530 K. For both catalysts CO2 and H20 were the main products, with a small amount of CO produced at higher temperatures (> 450 K), and no organic products were detected.

02 9

I

,

I

,

I

9

'

,

I

SiO2 I

I

I

I

I

10% MnO2/AI203

o

I

I

I

I

I

20

40

60

80

100

20/Degree Figure 2. XRD patterns of the supported MnO2 catalysts and the supports. The measurements of inlet and outlet ozone concentrations allowed the calculation of the ratio of converted ozone to converted ethanol (Fig. 5). Converted Ozone

%Ozone Conversion x Ozone Feed Pate

Converted Ethanol

%Ethanol Conversion • Ethanol Feed Rate

(1)

This ratio were found to decrease from 14 - 10 at 300 K to over unity at > 500 K (not shown). The ratio on the silica-supported sample was slightly higher than that on the alumina-supported catalyst. Because of the absence of crystalline manganese oxide, the reaction turnover rates on MnO2/Al2Eh were calculated as~ming 100% dispersion of manganese oxide. Activation energies of ethanol oxidation using oxygen and an ozone/oxygen mixture were calculated from the Arrhenius plots of the turnover rates (Fig. 6). The activation energy was much higher (89 kJ molq) when oxygen alone was used as the oxidant than when an ozone/oxygen

877 mixture was used. In the latter case the activation energy was only 3.7 kJ mol"1 at lower temperatures (< 400 K), and 48 kJ tool~ at higher temperatures (> 400 K). 40 40 o

~o 30

"~ 30 O

20

-~ 20

Ozone/O~ge~....~

,

gen

n_m__n_____._._i~ ~m~~'O/

300

350

400 450 500 Temperature / K

550

300

Figure 3. Ethanol oxidation reactivity using 03 or 02 on 10% MnO2/AI203.

O

gen

350

400 450 56o Temperature / K

Figure 4. Ethanol oxidation reactivity using 03 or 02 on 10% ]VlllO2/SiO2.

20

n/U~n,.,

~

MnO2/AI203

MnO2/SiO 2

O

~e

O

rj

0

,, , , ,

360

3;0 460 Temperature / K

4~o

Figure 5. Converted ozone to ethanol ratio on the supported MnO2 catalysts

-7a = 48 kJ

-S

b~ ~-9

~~ta

mol-1

= 89 kJ mo1-1

---10

-11 -12 1.8

.... 2.0 2.2

;~,.,. 2.4

EaT3:7.kJ,m~

2.6 2.8 3.0 103 T-1 / K-1

3.2

' 3.4

5;0

3.6

Figure 6. Arrhenius plots of the turnover rate of ethanol oxidation on MnO2/A1203

878 3.3. In Situ Laser Raman Spectroscopic Studies The Raman spectra of the catalyst samples were compared under various conditions (Fig. 7). The spectrum under oxygen exhibited a broad signal peaking at 658 cm-1 (Fig. 7a). With the introduction of ozone the 658 cm"l peak ~hiRed to 634 cmq and a new sharp signal at 878 cmq appeared (Fig. 7b). When ethanol was introduced in the absence of ozone, a new species with Raman signals at 884, 2878, 2930, 2970 cmq was observed which was assigned [9] to an adsorbed ethoxide species (Fig. 7c). Under the coexistence of ethanol and ozone (Fig. 7d), the intensity of both the new signal at 878 cmq and the ethoxide species decreased dramatically, and the signal at 658 cmq shifted to 640 cmq.

2936

640

2930

d

2970 A

c ~

b b a I

I

I

,

I

3000 2900 2800 R a m a n S h i f t / c m -1

I

,

I

~

I

~

I

,

I

1000 800 600 400 200 Raman S h i f t / c m -1

Figure 7. Raman spectra of 10% MnO2/AI203 sample under various conditions (a. oxygen; b. ozone; c. ethanol; d. ethanol and ozone). In addition to steady state experiments, transient measurements were also performed on the ahtmina-supported catalyst to investigate the interaction between the adsorbed species. The transient experiments were carded out by starting with a surface with preadsorbed ethoxide species. Ozone was then suddenly introduced and a set of Raman spectra was acquired at regular time intervals. The intensity of the 878 cm"l peak increased with time, while the adsorbed ethoxide peaks in the higher wavenumber region (2800 - 3200 cm"l) decreased with time (Fig. 8).

879

150 rain 120 rain

150rain ~

90rain

120 rain

r~

60 rain

90 rain

30rain

60rain 30rain

0min

0min

31'00 ' 3 0 ~ " 29~ ' 2s'o0

'12100'1000'

Raman Shift/G~I1-1

800 ' 600 ' 400

Raman Shift/cm-1

Figure 8. Transient experiment Raman spectra on 10% MnO2/A1203. 3.4. Ozone Decomposition Ozone decomposition reactivity was measured to investigate its role in the ethanol oxidation reaction (Fig. 9). On MnO2/AI203, with increasing temperatures ozone conversion increased slowly at temperatures < 400 K, then increased sharply above 400 K, and finally approached 100% around 500 K. While on MnO2/SiO2, ozone conversion was lower at low temperatures (< 320 K), but steadily increased with increasing temperatures, and reached 100% at a similar temperature (500 K) as on MnO2/A1203. The remover rates on MnO2/AI203 were also calculated assuming 100% dispersion of manganese oxide, and the Arrhenius plot showed two kinetic regions (Fig. 10). The lower temperature region which was dominated by the catalytic decomposition on MnO2 gave an activation energy of 3.2 kJ tool1, while the higher temperature region which was dominated by the gas phase decomposition gave an activation energy of 41 kJ m o r I. 100

-~r

o 80

-7.6

/;;:F

-7.8 a

41 kJmolq

b~ -8.0 ~

_~ -s.2

- . -

20

0

o~

- - - ~o~/~a~o~

./ 3;0

' 3;0

4;0

' 4;0 ' 560 Temperature / K

' 5;0

Figure 9. Ozone decomposition reactMty on supported Mn02 catalysts.

-8.4

E a = 3.2 kJmo1-1

-8.6 212'214'2:6'218'310'312'314' 103 T -1 / K d

Figure 10. Arrh~ttius plots of turnover rate of ozone decomposition o n MnOjA1203.

3.6

880 4. DISCUSSION

4.1. Reaction Stoichiometry For most reactions, reaction stoichiometry is easy to determine, however, this is not the case for oxidation reactions using ozone. One reason is that it is not clear whether oxygen is also involved in the reaction as ozone is usually used as an ozone/oxygen mixture. Another reason is that it is uncertain how many oxygen atoms of each ozone molecule contribute to the oxidation reaction. If only ozone is involved in the reaction, the reaction equation can be written as (assuming only CO2 and H20 are produced):

GH, OH + 203 = 2CQ + 3H20

(2)

However, when each ozone molecule only contn'butes one oxygen atom while producing an oxygen molecule, the reaction equation becomes: q

r,o/7 + 603 = 2 c Q + 6Q +

(3)

In addition to these reactions the ozone decomposition reaction may occur independently. 03 = 3/202

(4)

By measuring the inlet and outlet ozone concentrations, the converted ozone to ethanol ratio was determined, which may give insights on the reaction stoichiometry. The low ratio value at high temperatures (> 400 K) was consistent with a si~ifieant involvement of molecular oxygen in the reaction. In other words, ethanol could initially be activated by ozone to form intermediate oxidation products (aldehydes, earboxylic acids, etc.), and these could subsequently be oxidized by oxygen. At even higher temperatures (> 500 K), the oxidation was likely dominated by oxygen because of the rapid decomposition of ozone and because of the higher activation energy of ethanol oxidation by oxygen compared to that by ozone. At lower temperatures when the oxygen involvement was probably low, the measured ratios ranged fIom 10 to 20, which were closer to the stoichiometry of equation 3. The measured ratio on the silica supported catalyst was found to be slightly higher than that on the alu~mina catalyst, which is consistent with the observation that the silica catalyst was not as active as the a~mina one for ozone decomposition. In a study on the oxidation of lower aliphatic alcohols by ozone on alumina and silica from 293 and 363 K [4], it was found that the converted ethanol to ozone ratio was about unity. However, in that study the main oxidation products were acetaldehyde and ketones, while the only products observed on the supported manganese catalysts of this study were carbon dioxide and water.

4.2. In Situ Laser Raman Spectroscopic Studies Spectroscopic studies of the interaction between ozone and organic molecules on catalyst surfaces have been very rare. Mariey et al. [8] reported a Fourier transform in~ared (FFIR) study of ozone interaction with phenol adsorbed on silica and celia. Ozone was found to be reactive toward phenol and carboxylic acids and aldehydes were detected as possible intermediates. However, the study was carded out from 77 to 220 K, which is far from

881 reaction conditions (usually > 300 K) for those surfaces. Recent work from this laboratory has demonstrated that in situ laser Raman spectroscopy is an excellent tool to study surface intermediates at reaction conditions [10,11]. The signal at 650 cm-1 under oxygen can be ascribed to Mn304, however, this observation does not mean it is the only phase present. Other phases like, and Mn20 3, MnO show only very weak Raman signals while MnO 2 is completely Raman inactive. The introduction of ozone formed an adsorbed peroxide species with a Raman signal at 878 cm"l, while the introduction of ethanol generated an adsorbed ethoxide species with Raman signals at 884, 2878, 2930, 2970 cml. Under reaction conditions with the coexistence of ozone and ethanol, the intensities of both these adsorbed species dramatically decreased, indicating that these two species reacted with each other on the catalyst surface. This was also supported by the transient experiment results. When ozone was introduced on a surface preadsorbed with ethoxide species, the intensity of the ethoxide species decreased gradually due to the reaction with ozone (gas phase or adsorbed), and that of the peroxide species increased with time due to the removal of ethoxide species from the surface. However, if the reaction of ethoxide species was mainly due to gas phase ozone, under steady state conditions, the surface should be covered by adsorbed peroxide species. The in situ Raman spectra indicated that the reaction of ethoxide species was primarily due to reaction with adsorbed peroxide species because the concentrations of both adsorbed species decreased dramatically in the presence of both ethanol and ozone. Thus, a Langmuir-Hinshelwood type mechanism appears to be operating:

o3 + 2 " EtOH + 2 " E t o * + o~ *

2 0 * (or 0 2 . ) O*(orQ*)

+ 2H*

) o~*+o*

(5)

> EtO * + H *

(6)

:~" > c 0 2

) 02 + * (or Z*)

:,,n ." 1_120

(7) (8) (9)

Overall the results are consistent with both ethanol and ozone adsorption being the slow steps in the overall reaction. The surface reaction and desorption steps are fast, and result in a surface that is close to bare.

4.3. Role of Ozone Decomposition As indicated in equation 3, ozone decomposition can be closely associated with the oxidation reactions. It is well accepted that ozone decomposes to produce active oxygen species, which can activate the organic molecules at lower temperatures. It has also been reported by several groups [7, 12, 13] that at low temperatures the activation energy for oxidation reactions by ozone are similar to that of ozone decomposition, which suggests that the rate determining step for the oxidation reaction is probably ozone decomposition. The results from this study are consistent with that conclusion. However, ozone decomposition likely involves steps (5) and (8). Since at low coverage step (5) will be the rate-determining process, the measured activation energy probably corresponds to that step. In addition, the in situ laser Raman spectroscopic study provides direct evidence that surface ethoxide reacted with peroxide species formed by ozone decomposition

882 5. CONCLUSIONS Ethanol oxidation using ozone was investigated over supported manganese oxide catalysts at temperatures from 300 to 550 K by in sire laser Raman spectroscopy coupled with reactivity measurements. Ethanol was found to react with ozone at lower temperatures than with oxygen, and also with a lower activation energy. This is in agreement with the stronger oxidizing ability of ozone compared to oxygen. The oxidation reactivity was found to be closely related to that of ozone decomposition, suggesting an important role of ozone decomposition in the reaction mechanism In sire laser Raman spectroscopic studies provided direct evidence for the reaction between the peroxide (due to ozone) and the ethoxide species (due to ethanol) on the catalyst surfaces. ACKNOWLEDGMENT We gratefully acknowledge the financial support for this work by the Director, Division of Chemical and Thermal System of the National Science Foundation, under Grant CTS9311876. REFERENCES 1. A. Gervasini, G.C. Vezzoli, and V. gagaini, Catal. Today, 29 (1996) 449. 2. A. Gervasini, C.L. Bianchi, and V. gagaini, in Environmental Catalysis, J.N. Armor (ed.), ACS Symp. Set. 552; ACS: Washington, DC, 1994, 352. 3. W. Li and S.T. Oyama, in Heterogeneous Hydrocarbon Oxidation, B.I~ Warren and S.T. Oyama (eds.), ACS S3anp. Set. 638; ACS: Washington, DC, 1996, 364. 4. M. N. Klimova, B.I. Tarunin, and Yu.A. Aleksandrov, Kinet. Katal., 26 (1988) 1143. 5. K, Hauffe and Y. Ishikawa, Chem. Ing. Techn. 5 (1974) 1035. 6. V. Ragaini, C.L. Bianchi, G. Forcella, and A. Gervasini, in Trends in Ecological Physical Chemistry, Bonati, L. (eds.) Elsevier: Amesterdam, 1993, 275. 7. A. Naydenov and D. Mehandjiev, Appl. Catal., A 97 (1993) 17. 8. L. Mariey, J. Lamotte, J.C. Lavalley, N.M. Tsyganenko, and A.A. Tsyganenko, Catal. Lett., 41 (1996) 209. 9. W. Zhang and S.T. Oyama, J. Phys. Chem., 99 (1995) 19468. 10. W. Zhang and S.T. Oyama, J. Phys. Chem, 100 (1996) 10759. 11. W. Zhang and S.T. Oyama, J. Am. Chem See., 118 (1996) 7173. 12. N.A. Kleimenov and A.B. Nalbandian, Proceedings of Academy of Science of USSR, 122 (1958) 635. 13. 1L Del gosse, C. Mazzocchia, and P. Centoh, React. Kinet. Catal. Lett., 5 (1976) 245.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

883

Generation o f singlet oxygen from the catalytic system H202/Ca(On)2 and applications to the selective oxidation o f tmsaturated compounds J. M. Aubry* and V. Nardello Equipe de Recherches sur les Radicaux Libres et l'Oxyg~ne Singulet, URA CNRS 351 Facult~ de Pharmacie de Lille, 3 rue du Professeur Laguesse, BP 83 F-59006 Lille Cedex, France

A new chemical source of singlet molecular oxygen (Io2, lAg), H202/Ca(OH)2, has been investigated in detail. First, the formation of ~O2 has been proved unambiguously by resorting both to the specific detection of the IR luminescence at 1270 nm of Ioz and to the chemical trapping of the excited species with a new cationic water-soluble trap. The process has been shown to be catalytic and the influence of several parameters (pH, concentrations and purities of reactants) on the initial rate of ~O2 formation has been examined. Finally, the ability of the system H20~/Ca(OH)2 to oxidize various water-soluble electron-rich substrates has been assessed.

INTRODUCTION

Excited molecular singlet oxygen, (IO2, ~Ag) is a powerful and highly selective oxidant very useful in organic synthesis. ~'2 The usual way to generate IO2 is photochemical but many chemical reactions are also able to produce this excited species by decomposition of a peroxo compound. 3 In 1985, we showed that about thirty inorganic oxides, hydroxides or oxo-anions induce the decomposition of hydrogen peroxide into ~O2.4 Among all these new chemical sources of ~O2, the homogeneous system H202/MoO42 appeared particularly efficient since it generates quantitatively a high flux of 102 at room temperature (Eq. 1) :

2H2Oz

MoO~" water, pH 10

>

2H=O + 'O2

(1)

Since then, much work has been devoted to this system and both the mechanism of the reactions and the ab'~ty to oxidize various organic substrates, either in water6 or in microemulsions,7 have been investigated in detail.

884 The screening experiments of the periodic classification revealed another particularly interesting chemical source of IO2 9the system H202/Ca(OH)2 (Eq. 2) which appears both attractive for its environmental friendly feature and amazing with regard to the simplicity of the catalystinvolved.

2 H20~,

Ca(OH)= water >

2 H20 + ~ ~0 2 +(1 - (x) z02

(2)

For these reasom, we have undertaken a detailed study of this system by confirming, in a first step, the generation of 102 by chemical trapping and by detection of its specific IRluminescence. Then, the kinetics of the reaction were examined by studying the influence of several parameters such as pH and concentrations and purities of reactants on the initial rate of IO2 formation. Finally, the ability of this system to oxidize various water-soluble electron-rich substrates such as polycyclic aromatic, cyclohexadienic and acrylic derivatives was assessed.

1. EXPERIMENTAL PART 1.1. Reagents Calcium hydroxide (98 %), calcium oxide (99.9 and 99.995 %), and tiglic acid 7 were purchased from Aldrich Chemie. Calcium peroxide was l~om Air Liquide. Hydrogen peroxide (50 % Rectapur) was from Prolabo, Paris. Deuterium oxide (98 % D) was l~om CEA (Commissariat ~ l'Energie Atomique, Saclay). bis-(4'-trimethylphenylammorfium)-9,10anthracene dichloride 1 (BPAA), s potassium 9,10-anthracene dipropionate 3 (ADP) 9 and sodium 1,3-cyclohexadiene-l,4-diethanoate (CHDDE) ~~5 were prepared according to known procedures. 1.2. Instrumentation Disappearance of BPAA was monitored by UV/visl'ble spectroscopy at 373 nm with a Milton Roy Spectronic 3000 spectrophotometer. High performance liquid chromatography (HPLC) analyses were carried out with a CAIson pmnp model 303 by using a 25-cm cohmm packed with Spherisorb RP18-50DS. A mixture of H20 and CH3OH was used as eluent and UV detection was performed with a variable-wavelength monitor (Gilson Holochrom H/MD). The nuclear magnetic resonance spectra (~H and 13C NMR) were obtained on a Bruker AC 300P FT-spectrometer. 1.3. Singlet Molecular Oxygen Monomol Emission (1270 nm) Infrared emission of 102 was measured with a liquid nitrogen cooled germanium photodiode detector (Model EO-817 L. North Coast Scientific co., Santa Rosa. CA) sensitive in the spectral region from 800 nm to 1800 nm with a detector of 0.25 cm2 and a saphire window. Measurements were carried out in a cuvette with mirrored walls (35 m m x 6 m m x 55 nun).

885

1.4. Chemical Detection of Singlet Molecular Oxygen 90 ~tl H202 30 % (0.2 M) were added, under stirring, to an aqueous solution (I)20) of BPAA 1 (10.2 M) and CaO 99.995 % (0.1 M). The disappearance of BPAA was monitored by UV spectroscopy at 373 nm from acidified (H3PO4) samples and the final reaction medium was analyzed by ~H and ~3C NMR spectroscopy alter centrifugation. 1.5. Deuterium Solvent Effect 25 ~tl H202 50 % (0.1 M) were added under stirring to an aqueous solution (I)20 or H20) of BPAA (10"3 M) and CaO 99.995 % (0.05 M). The disappearance of BPAA was monitored by UV spectroscopy at 373 nm from acidified (H3PO4) samples. A similar procedure was used for the study of the influence of pH, concentrations and purities of reactants. 1.6. Peroxidation of Tigfic Acid 7 5.75 ml H202 50 % (10.0 M) were added under stirring to an aqueous solution (H20) containing 1 g tiglic acid 7 (1 M) and 1.136 g CaO 99.99 % (2 M) at 35 ~ The pH of the heterogenous reaction medium was equal to 7.75. The course of the reaction was followed by HPLC at 210 nm from filtered aliquots (1 rrd) and the final reaction medium was analyzed after 4 h by 1HNMR spectroscopy. Comparisonwith a genuine sample obtained photochemicaUy showed that 50 % of the corresponding hydroperoxide 8 was formed. A similar procedure was used for the peroxidation of ADP 3 and CHDDE 5. The experimental conditions are given in Table 1.

2. RESULTS AND DISCUSSION. 2.1. Evidences for the generation of singlet oxygen (~Oz, tag) from the system HzO2/Ca(OH)2 The first aim of our study was to prove unambiguously the generation of 102 from the system H202/Ca(OH)2. One method used was the chemical trapping of this excited species. Since the current water-soluble chemical traps of IO2 bear carboxylate or sulfonate functions, II which are likely to interact with the calcium ions, a new cationic water-soluble trap was designed:the bis-(4'-trimethylphenylammonium)-9,10-anthracene dichloride 1 (BPAA). This trap reacts efficiently with IO2 giving the endoperoxide BPAAO2, 2 (Eq. 3). This trapping process com~tes with the solvent-induced quenching of 102 (Eq. 4).

886

N(CH3) cr

N(CH3)"cr

(3)

,~ (so oc)

! N(CH3) cr

(CH3) cr

BPAA, 1

BPAAO2, 2

102

water

>

302

(4)

The formation of BPAAO2 was confirmed both by ~H and ~3C NMR spectroscopy by comparison with a genuine sample photochemically obtained and by thennolysis at 80 ~ of the final reaction mixture which led to the regeneration of the initial trap according to a specific property of some polycyclic aromatic endoperoxides. 12However, although the formation of the endoperoxide BPPAO2 constitutes a convincing proof of 102 involvement, it is useful to strengthen this result by another independent test based on the longer lifetime of the excited species in deuterated water (67 I~s in D20 instead of only 4.4 ~ in H20). ~3 In order to understand an increase in rate in D~O, it is a s s u n ~ that the main process is the deactivation of ~O2 by solvent i.e. BPAA must be introduced at a concentration well below [3 = l~/k, where Iq is the pseudo-first order rate constant for the deactivation of IO2 by water (Eq. 4) and 1~ is the second-order rate constant for the chemical reaction between '02 and BPAA (Eq. 3). Therefore, two experiments were carded out by using 10.3 M BPAA to the system H202/Ca(OH)~ in D20 or in H20. The disappearance of BPAA, monitored by UV spectroscopy at 373 nm, is reported in Figure 1. Comparison of the initial rates of the disappearance of BPAA and of the final amounts of BPAAO2 leads to an increase in rate by a factor of 13 when the reaction is conducted in deuterated water, giving support for a process involving 102.

887

1 0,8 rb

----"--'-*

A

E e.

,.'~ 0,6 t~ v 0

<

<

0,4

0,2

0

10

20

30

40

50

60

70

Time (h) Figure 1. Comparison of the disappearance of BPAA 10"3 M in ordinary (e) and in deuterated (o) water (room temperature, natural pH, [CaO]o = 0.05 M, [H202]o = 0.1 M). The involvement of ~O2 during the process under study has also been confirmed by resorting to the detection of the IR hnninescence of ~O2 at 1270 rim. This method is specific, direct and rapid but requires work in deuterated water and at a higher ten~rature than room temperature in order to increase the stationary concentration of ~O2. When hydrogen peroxide (1 M) was added under stirring to an aqueous solution (D20) containing CaO (0.2 M) at 50 ~ a significant signal of IO2 luminescence (4.3 mV) was detected. The loss of the signal could be observed by addition of sodium azide, an efficient quencher of IO2, and when the reaction was conducted in ordinary water, the intensity of the signal was about ten-fold lower. All these results, obtained by two complementary methods, confirm unequivocally that singlet oxygen is actually generated during the decomposition of hydrogen peroxide induced by calcium hydroxide. 2.2. Stoiehiometry of the reaction In order to rule out the assumption that the generation of ~02 might come from the disproportionation of H202 induced by a mineral impurity contaitm~ in one of the reactants, several control experiments were carried out. Neither the nature of water (distilled, boiled or degased with argon) nor the origin of hydrogen peroxide (50 % Normapur prolabo or 30 % perhydrol Merck) exhibit any influence on the reaction. On the contrary, comparison of two samples of calcium hydroxide of different purity (Ca(OH)2 98 % and CaO 99.995 %) showed s'nnilar initial rates of 102 generation but a noteworthy increase in the cumulative amount of IO2 calculated from Eq. (5) 14when the purest reactant was used (Fig.2).

888

[102]t =

[BPAA]o [BPAA]o-[BPAA]t+ ~ In [BPAA]t

(5)

This difference is readily explained by the occurrence of side reactions involving the disproportionation of H202 into ground state oxygen induced by impurities contained in Ca(OH)2 98 %. This observation confirms that the generation of ~O2 acttmlly results from an interaction between H202 and Ca(OH)2. 50 411

:Z A

0 x

30 20

O

10

0

5

10

15 Time (h)

20

25

30

Figure 2. Comparison of cumulative amounts of ~O2 generated during the disproportionation of H202 (0.4 M) in the presence of Ca(OI-I)2 98 % (0.05 M) (o) or CaO 99.995 % (0.05 M) (e) ([BPAA] = 104 M, T = 25 ~ natural pH). The decomposition of variable hydrogen peroxide concentrations (0.1 to 0.5 M) by calcium oxide (0.05 M) was also studied. Titration of H202 at the end of the reaction showed that in all cases, it was almost completely consumed by an amount of CaO about ten times lower. Therefore, the process under study is catalytic. Moreover, the final molar amount of H202 titrated in the solid phase was found to be equal to the initial molar amount of CaO whereas no more 1-1202was present in solution. This other finding suggests that, at the end of the reaction, a stable calcium monoperoxide is formed, releasing one equivalent of H202 during the titration in acidic medium. This compound is very likely to be the well-documented CaO~. 8H20, known to be stable under our experimental conditions. ~s 2.3. Kinetics of the reaction In order to study the influence of pH, concentrations of reactants, temperature and nature of the catalyst, we preferred to resort to the chemical trapping method which allows work in ordinary water at room temperature and which leads to a reliable quantification of ~O2.

889 The initial rates of ~O2 formation were then inferred from the amount of ~O2 determined from equation (5). As for the system H202/MoO42", 16 the reaction exhibits a pH dependence with a bellshaped curve and a maximum around pH 10.5. In very alkaline media, H202 dissociates and the rate of ~O2 formation decreases (Fig. 3).

0,006 0,005 ',- 0,004

0,003 0,002 0,001 0'.

9

9,5

10

10,5

11

11,5

12

pH

Figure 3. Influence of pH on the initial rate of IO2 formation by the system H202/Ca(OH)2 ([BPAA] = 10"4 M, [CaO] = 0.05 M, [H,O2] = 0.4 M, T = 25 ~

buffer).

The reaction was found to be first order with respect to calcium hydroxide. On the other hand, the rate law appears more intricate in the case of H~O2 (Fig. 4) and cannot be interpreted in terms of order of reaction. We can just assmne that an increase in HaOa concentration induces an increase in the rate of los formation suggesting that the intermediate respons~le for 102 generation is the most peroxidized species, in contrast to the system H202/MoO42".5

890

A

"7 J=

A

N

O > X t~

O

0

1

2

3

[HaOj (M)

4

5

6

Figure 4. Influence ofH202 concentration on the initial rate of 102 formation in the presence of CaO 0.05 M ([BPAA] -- 104 M, T - 25 ~ pH - 10.0 + 0.1). 2.4. Oxidation of water-soluble substrates by the system H202/Ca(OH)2 Three typical water-soluble substrates including the polycyclic aromatic hydrocarbon 3, the cyclohexadienic derivative 5 and the olefin 7 were oxidized in order to assess the efficiency of the oxidizing system H202/Ca(OH~. These subsWates represent standard types of singlet oxygen reactions, namely [4 + 2] cycloaddition and the ene reaction. The disappearance of the substrates was monitored by HPLC and the products were identified by comparison with genuine smnples, photochemically obtained. Experimental conditions and results are summarized in Table 1. The assays carded out with potassium 9,10-anthracene dipropionate 3 and sodium 1,3cyclohexadiene-l,4-diethanoate 5 confirm once again the involvement of IO2 in the process under study since the corresponding endoperoxides, respectively ADPO2 4 and CHDDEO2 6, can exclusively be obtained by a process involving '02. On the other hand, they show the efficiency of the catalytic system H202/Ca(OH~ towards oxidation since, in a relatively short time, in both cases, all the starting material is consumed giving the expected oxidation products in quantitative yields. Nevertheless, it has recently been shown that the reaction of CHDDE 5 with '02 photochemically generated leads not only to the endoperox~e 6 (88 %) but also to the corresponding hydroperoxide (12 %).~0 Here, this latter product was not detected, probably on account of an adsorption on the calcium peroxide which was filtered before HPLC and NMR analyses.

891

Table 1 Oxidation of various water-soluble derivatives by the system H202/Ca(OH)2 ([substrate] = 10 "2 M, [1"I202] 1 M, [CaO] = 0.2 M, D20, T = 25 ~ (a) [substrate] = 1 M, [CaO] = 2 M, [8202] 10 M~ H20, T = 35 ~ =

" - -

Substrates

Oxidation products

CH~,CH2COOK

At

Yields (%)

10.4

25 min

100

9.8

1.5 h

100

7.8

4.2 h

50(-)

CH2CH2COOK

CH2CH2COOK

CH2CH2COOK ADP, 3

pH

ADPO2, 4

--COONa

OONa

COONa

L--COONa

CHDDE, 5

CHDDEOz, 6

I"13C~.

H3C~~1-13 --H COOH

HO0 / H

ffCH2\

Tiglic acid, 7

Hydroperoxide, 8

COOH

The most interesting result concerns the oxidation of tiglic acid 7. This substrate is actually an electron-deficient olefin and its reactivity towards Io: is relatively low. 6 Nevertheless, the oxidation of a high concentration of tiglic acid (1 M) in H~O provided the corresponding hydroperoxide 8 in fair yield (50 %).

CONCLUSION. The generation of singlet molecular oxygen from the catalytic system H202/Ca(OH~ is now straightforward. This new chemical source of ~O2 is able to oxidize various water-soluble substrates in goods yields and could be an attractive alternative to the photochemical generation of ~O2. The application of this oxidizing system to hydrophobic organic substrates is

under study.

892 REFERENCES.

1. M. Matsumoto, Singlet Oxygen, Frirner, A. A. Ed. CRC Press, Boca Raton, Fl. 2 (1985) 205. 2. H. H. Wasserman and IL W. Murray, Singlet Oxygen, Academic Press, New York, 1979. 3. J. M. Aubry, New chemical sources of singlet oxygen, in <<Membrane lipid oxidation ~, Vol. II, Ed. Vigo-Pelfrey, CRC Boca Raton, (1990) 65-101. 4. J.M. Aubry, J. Am. Chem. Soc., 107 (1985) 5844. 5. V. NardeUo, J. Marko, G. Venneersch, J. M. Aubry, Inorg. Chem., 34 (1995) 4950. 6. V. Nardello, S. Bouttemy, J. M. Aubry, J. Mol. Catal., 117 (1997) 439. 7. J. M. Aubry and S. Bouttemy, J. Am. Chem. Soc. accepted for publication. 8. V. Nardello, J. M. Aubry, unpublished.. 9. B. A. Linding, M. A. J. Rodgers, A. P. Schaap, J. Am. Chem. Sot., 102 (1980) 5590. 10. V. Nardello, N. Azaroual, I. Cervoise, G. Vermeersch, J. M. Aubry, Tetrahedron, 52 (1996) 2031. 11. V. Nardello, D. Brault, P. Chavalle, J. M. Aubry, J. Photocher~ Photobiol. B., in press. 12. N. J. Turro, M. F. Chow, J. Rigaudy, J. Am. Che~ Soc. 103 (1981) 7218. 13. A.A. Gormann and M. A. J. Rodgers, Singlet Oxygen in Handbook of Organic Photochemistry, Scaiano, J. C. Ed. CRC Press, Boca Raton, Vol. II, (1989) 229-247. 14. F. Wilkinson, W. P. Helman, A. B. Ross, J. Phys. Chem. Ref. Data, 24 (1195) 663. Reprint n~ 489. 15. S. Z. Makarov and N. K. Grigor'eva, ZE Pdkl. Khim. 32 (1959) 2184. 16. J. M. Aubry and B. Cazin, Inorg. Chem 27 (1988) 2013.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

893

T o l u e n e gas p h a s e o x i d a t i o n to b e n z a l d e h y d e and p h e n o l over V - c o n t a i n i n g micro- and m e s o p o r o u s materials G. Centi, a F. Fazzini a and L. Canesson, b A. Tuel b a Dip. Chimica Ind. e Materiali, V.le Risorgimento 4, 40136, Bologna, Italy b Inst. de Rech. sur la Catalyse, CNRS, av. A.Einstein 2, 69 626 Villeurbanne, France

Gas phase heterogeneous oxidation of toluene using gaseous 02 over V-containing microand meso-porous materials (MFI, 13 and MCM-41 structure) was studied. Benzaldehyde is the main product, with selectivities up to about 60%, but only at low conversion. Both the nature of the zeolite and its acidic characteristics considerably influence the selectivity to benzaldehyde. Isolated vanadium species in distorted octahedral coordination and the absence of strong acid sites are preferable for selectivity. At high temperature, phenol also forms (selectivities up to about 6%), especially on mesoporous materials with the higher vanadium loading. The formation of phenol by gas phase oxidation of toluene with gaseous 02 is a unique characteristic of the reactivity of these catalysts as well as an interesting example of a new kind of selective oxidation reaction. 1. INTRODUCTION The selective partial oxidation of alkylaromatics is commercially done using complex multistep processes with the formation of considerable amounts of wastes as well as problems and costs of separation and in some cases also of corrosion. For these reasons alternative and cleaner processes are being sought. Heterogeneous selective oxidation using 02 is a convenient solution in terms of both ease separation and reduced waste production, but this method suffers from low selectivity due to the considerable difference between the enthalpies of the weakest bond dissociation in reactant and products giving rise to a fast consecutive oxidation to COx [ 1]. The rate of the consecutive oxidation can be controlled by site isolation, which, however, requires the possibility of obtaining a high dispersion of the active component in a matrix with uniform pore size distribution. Transition metal containing microporous materials are thus an interesting class of catalysts for this investigation which offer the further advantage of possible shape selective effects to better tune the catalytic properties. The preferable choice for the transition metal is vanadium, since most of the catalysts reported in the literature for selective side chain oxidation of alkylaromatics (toluene especially) are based on vanadium as the key component [2-6]. Although V-containing zeolites have already been studied in gas-phase selective oxidations (propane oxidative dehydrogenation, for example [7]) data on the reactivity of these catalysts on the heterogeneous gas-phase selective oxidation of alkylaromatics are not available. Alkylaromatic oxidation has also been done in the liquid phase using H202 as oxidant [8]. The

894 scope of this work was to carry out a preliminary study of the reactivity of these catalysts using toluene to benzaldehyde as the test reaction, in order to determine the outlook for and drawbacks of the use of this class of catalysts in alkylaromatic selective oxidation and identify the relationship between nature of vanadium species and catalytic behavior. In particular, the study was centered on the analysis of the behavior in toluene selective oxidation of a series of V-containing micro- and mesoporous materials [mesoporous (MCM-41), MFI (ZSM-5) and 13type]. The formation of phenol as a by-product has also been observed in some cases. This product does not form in toluene oxidation over vanadium oxide supported on A1203 or TiO2 [2-6] and is an interesting first example of the possibility of direct synthesis of phenol from toluene using gaseous 02. A further objective of this work therefore was to identify the key aspects in this reaction as well as the possible reaction mechanism. 2. EXPERIMENTAL

2.1 Synthesis of the Catalysts. Vx-HMS: Vx-HMS samples (MCM-41 type, mesoporous structure; the symbol x indicates the amount of vanadium in wt % in the samples) were synthesized using tetraethyl orthosilicate (TEOS), vanadyl acetylacetonate, hexadecylamine and alcohols according to a procedure reported in the literaure [9]. Two solutions are prepared : (A) solution A is prepared by mixing TEOS (0.1 mol) with 0.65 mol of ethanol and 0.1 mol of isopropyl alcohol. Then vanadyl acetylacetonate is added and the resulting solution stirred for about 30 minutes. (B) solution B contains hexadecylamine (0.027 mol), water (3.6 mol) and hydrochloric acid (0.002 mol). Solution A is added to solution B under vigorous stirring. Stirring is maintained for about 1 hour and the mixture is kept at room temperature for 12 hours under static conditions. The resulting light green solid is recovered by filtration, washed several times with distilled water and airdried at room temperature. The yield in V-HMS is always very high (typically > 95 %). To remove occluded organics, samples are calcined in air at 600 ~ for 12 hours. The chemical composition and textural characteristics of the samples are summarized in Table 1. Table 1. Chemical composition and textural characteristics of Vx-HMS samples. Sample Si/V S(m2/g) ~p(/~) Vx-HMS Gel Product 1 100 109 1124 34 2 50 57 1015 34.5 3 33 30 953 35 5 20 19 781 34 Op(A) is the mean pore diameter corresponding to the maximum in the pore size distribution curve.

Vx-MFI and Vx-fl: Differently from the above catalysts, in V-MFI (ZSM-5 structure) and V-~ (Beta structure) the zeolite was prepared first and then the vanadium component was added by incipient wet impregnation at r.t. using a peroxide complex of vanadium obtained by reaction at r.t. of V205 with H202. The use of the V-peroxide complex increases the solubility of the vanadium complex, avoiding the possible precipitation at contact with the zeolite crystals and thus avoiding deposition on the outer surface.

895 ZSM-5 (SIO2/A1203 = 90) was synthesized with the following method. To 43.44 g of Na2SiO3 solution are added 1.46 g A12(SO4)3 in 52 cc H20 under vigorous stirring ; then 4.82 g NaOH are added. Successively TPABr (5.48 g) is added to the clear solution and then the solution is acidified with 6 ml H2SO4 in 20 ml H20 up to complete gelification. The gel is transferred to an autoclave where it is maintained under stirring at 120~ for 5 days. The synthesis leads to microcrystalline samples with a Si/A1 ratio close to that of the starting compounds. Beta zeolite (SIO2/A1203= 30) was synthesized using the following procedure. To 5.8 g of NaAIO2 solution are added under stirring 58 ml TEAOH and then 145.35 g of Silica Sol (Ludox). The gel is put in an autoclave and maintained under stirring for 6 days at 150~ After extraction from the autoclave and washing, both zeolites are calcined at 600~ for 3h, using a slow rate of heating (about 13 h from r.t. to 600~ to reach the final temperature, in order to avoid loss of crystallinity. The structure and crystallinity of the samples were checked by XRD and IR analyses. 2.2 Characterization of the Catalysts. The catalysts were characterized by X-ray diffraction (Philips PW 1710 diffractometer, (CuKa radiation) and FT-IR (Perkin Elmer 1750 instrument) analyses. N2 adsorption/desorption isotherms were recorded at 77~ using a Catasorb apparatus. NHa-TPD (thermodesorption) tests were made in a flow system (0.3 g sample) after adsorption of ammonia (5% NH3 in He) at 100~ up to complete saturation, flushing with He for 10 min and then heating up to 550~ at a rate of 10~ The catalysts were pretreated in a pure He flow at 400~ up to complete removal of adsorbed water. The NH3 concentration was followed with an on-line mass quadrupole detector. 5~V wide-line NMR spectra were collected on a Bruker MSL 300 FT NMR spectrometer operating at 78.9 MHz and equipped with a special probe head for measurements in the absence of air. UV-Visible-NIR diffuse reflectance spectra (DRS) were measured in air using a Perkin Elmer Lamda 19 instrument after transformation of reflectivity (R) according to the Kubelka-Munk function: F(R) = (1-R)2/2R**.BaSO4 was used as a reference. 2.3 Catalytic Tests Catalytic tests were made using a quartz continuous flow reactor (5 mm inner diameter) loaded with 0.5 g of catalyst in the form of small pellets (0.1-0.3 mm range). The feed composition was 3% toluene with a stoichiometric O2/toluene ratio (1.5). The space-velocity was set to 14500 h -~ to avoid diffusional limitations. Tests were made in the 250-550~ range, but above 500~ severe deactivation of the catalysts occurs. Before the tests the catalysts were conditioned at 450~ for 6 h in the presence of standard feed. The axial temperature profile was determined by a thermocouple inserted in the catalytic bed. Preliminary experimental tests were made to ensure the absence of mass and heat diffusional limitations on the reaction rates. The analysis of the feed composition and of the reaction products was made using two gas chromatographs (GC), the first on-line for the analysis of incondensable gases (N2, 02, CO, CO2). All organic products and water were adsorbed in two traps containing acetone as the solvent and maintained at about -5~ The products in the absorption traps were analyzed using a GC equipped with a mass quadrupole detector after adding tridecane as an internal standard. The first GC uses a TCD detector and a 3m long Carbosphere packed column, and the second a HP-5 5% PhMe Silicone 30m long capillary column.

896 75

o....

~

Conv. TOL

o.. ...... o...

.>

-o-9 SeI-BA ~ SeI-CO2

.....

........

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---~-- Sol. CO --49 Sel.-PhOH "*'o. "-.... -..o.. ""'o

._g ..~..

~

-,..-~ - - ~

350

400

~-r

450

~*'"~

" T - .--'~ 500

Reaction temperature, *C

Fig. 1 Effect of reaction temperature on the catalytic behavior of V2-HMS in toluene oxidation.

0I

,.

/

~

conv. TOL Sel. BA Sel. PhOH

\

'~ 40

c 0

~ --o~-

20

O

1

2

3

4

5

Vanadium loading, % wt.

Fig. 2 Effect of vanadium loading in Vx-HMS catalysts on the selectivity to benzaldehyde at 400~ and to phenol at 500~ and on the conversion of toluene at 450~

3. RESULTS AND DISCUSSION

3.1 Catalytic Behavior The effect of the reaction temperature on catalytic behavior in toluene oxidation of V2-HMS is shown in Figure 1. Due to the high space-velocity used in these tests to ensure absence of diffusional limitations, the conversion of toluene (TOL) is relatively low (1-14%) in the 350-500~ temperature range. Higher reaction temperatures lead to a significant rate of catalyst deactivation due to formation of heavier, condensed products as shown by the fact that the zeolite becomes darkgrey. Benzaldehyde (BA) is the main reaction product with selectivities in the 41-72% range, but the selectivity decreases with increasing temperature of reaction and conversion of toluene. The formation of benzoic acid was not observed, differently from vanadium-oxide supported over TiO2 or A1203 [2-6], but the formation of phenol (PhOH) in the 0-3% range was detected as the main byproduct. The selectivity of phenol increases as temperature increases with a consequent more marked increase in its yield. A series of minor by-products with selectivities below 1% were also observed that will be discussed later.

The effect of the loading with vanadium in the 1-5% wt. range (as V2Os) is shown in Figure 2 for the Vx-HMS series of catalysts using as indexes of the catalytic behavior the selectivity to benzaldehyde at 400~ and to phenol at 500~ and the conversion of toluene at 450~ Increasing the loading with vanadium nearly linearly increases the activity of the catalyst, apart from loading above about 3% where the slope decreases. This indicates that the rate of toluene depletion is proportional to the vanadium content when good dispersion is achieved, but above about 3% vanadium loading, nanoparticles of V-oxide probably start to form. The selectivity to benzaldehyde passes through a maximum for a loading of vanadium of 2%,above which further increases in the content of vanadium in the zeolite result in a drastic decrease in selectivity. The selectivity to phenol instead increases as the amount of vanadium

897 60

(5 50

--IF-- 1% V-oxide --W- 3% V-oxide

5

.~

in the catalyst increases, reaching selectivities of about 6% for V5-HMS.

06

The effect of the nature of the zeolite for vanadium O loadings of 1% and 3% is t3 |e ao / summarized in Figure 3 reIX. porting, as before, the se/ en ~ 20 lectivity to benzaldehyde at ._~ /! .~_ 400~ and to phenol at _.~ 10 .__ .___ ~_.~ ......~_..... ~o 500~ as indexes of cataffl lytic behavior. The data 0 0 show the considerable efBeta MFI HMS fect of the zeolite on the Type microporous material catalytic behavior regarding the selectivity to both benzaldehyde and phenol. The Fig. 3 Effect of the nature of the zeolite on the selectivity to benzalcatalysts based on Beta dehyde at 400~ and to phenol at 500~ zeolite show poorer catalytic behavior, whereas those based on MFI have intermediate selectivity. Although in going from HMS to Beta samples the formation of by-products increases, the amount of these remain below 5% (as global selectivity) and the worsening in the catalytic performance is mainly due to increases in the formation of carbon oxides. It should be noted that the order of selectivity (Fig. 3) does not follow the order of the pore dimensions: MFI < Beta < Mesoporous. ~ID

4 8t O

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.t..,

~

~-

v

|

i

3.2 Reaction network

The data in Fig. 1 show a drastic loss of selectivity with increasing reaction temperature and conversion, indicating a great sensitivity of benzaldehyde towards consecutive transformations, but do not allow discrimination between the effect of temperature and conversion of toluene and oxygen. Therefore, experiments were carried out at isotemperature and isoflow rate with different amounts of the catalyst or the O2/toluene ratio in the feed. The results are summarized in Figures 4A and 4B, respectively. As the amount of the catalyst increases (Fig. 4A) the conversion of toluene does not follow a nearly linearly increase as expected for a first order rate equation of hydrocarbon depletion (the low hydrocarbon conversion indicates pseudo-differential conditions with respect to hydrocarbon axial profile), because the concentration of oxygen also affects the reaction rate. This is confirmed in Fig. 4B showing the dependence of the conversion of toluene on the O2/toluene ratio in the feed. The rate of toluene depletion thus is not limited by the rate of adsorption of the hydrocarbon, but probably by the consecutive steps of oxidation. The selectivity to benzaldehyde shows a maximum with respect to the increase in contact time. The initial slight increase derives from the fact that the decrease in the O2/toluene ratio along the axial direction of the catalytic bed (caused by the catalytic reaction) causes an increase in the selectivity to benzaldehyde (Fig. 4B), but the further increase in contact time leads to further consecutive oxidation of benzaldehyde with a consequent lowering of the se-

898

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c o n y . TOL Sel. BA Sol. PhOH

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- - e - - c o n v . TOt. Sel. BA - - - 4 - Sol. PhOH -O 9 - SeL CO 9. ~ , . . S e l . C O , .& . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ...-"

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.

.

.

.

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.

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. 3

O:~/Toluene Ratio in F e e d

Fig. 4 Effect of amount of catalyst (A, left) and O2/toluene ratio in the feed (B, fight) on the catalytic behavior of V3-HMS at 400~ and 500~ respectively. lectivity. The sensitivity of benzaldehyde towards consecutive oxidation is also shown in Fig. 4B reporting the effect of the OJtoluene ratio in the feed at isotoluene concentration. The increase in 02 concentration leads to a linear increase in the conversion and corresponding decrease in the selectivity to benzaldehyde, whereas the selectivity to phenol increases. It is also noticeable that the selectivity to CO is nearly constant, whereas that to CO2 increases corresponding to the decrease in benzaldehyde selectivity. This suggests that CO probably derives from a side reaction of toluene conversion, whereas CO2, or at least part of it, derives from the consecutive conversion of benzaldehyde. Analysis of the nature of the by-products gives further interesting indications on the reaction network. The by-products formed can be classified into two groups: (i) products of alkylation/condensation reasonably connected to acid catalyzed reactions, and (ii) products of oxidation or their consecutive transformation/decarboxylation. The relative amount of the products of the first type increases when the zeolite support has acid properties and thus decreases in the order Beta > MFI > HMS, and increases with increasing vanadium content due to the introduction of medium-strong acid sites associated with this element. The order of acidity was determined by NHa-TPD experiments (not reported for sake of brevity). Both the number and acid strength of the acid sites decrease in going from Beta to MFI and HMS zeolites, but the addition of vanadium creates a medium-strong acidity proportional to the vanadium content. An increase in reaction temperature leads to an increase in the formation of these by-products, although very low amounts are usually detected even for reaction temperatures of about 350~ for the more acid samples. At temperatures above about 500~ the zeolites also become darkgrey indicating the formation of carbonaceous deposits. At lower temperatures (about 450~ howwever, deposits of polyenic or polyaromatic products were detected by UV-VIS-NIR/DRS analysis. The following products of alkylation/condensation were detected: diphenyl-methane, methyl-diphenylmethane and xylenes plus minor amounts of styrene, benzofuran, indene, naphthalene, stilbene, anthracene and fluorene. Climent et al. [ 10] observed the formation of diarylmethanes as one of the primary products of acid-catalyzed condensation of benzaldehyde

899 with benzene over zeolites. As discussed below benzene is another significant by-product of reaction which forms either by transalkylation (with consequent formation of xylenes) or by decarboxylation of the oxidation products of toluene. The formation of diphenyl methane and derivates thus indicates that an additional cause of loss of selectivity to benzaldehyde derives from its reactivity towards condensation reactions with other aromatics. Another interesting by-product, although in very low amounts, is stilbene which suggests the possibility of oxidative coupling of two toluene molecules via probable formation of the corresponding radical by abstraction of one H atom from the methyl group. The attack of the tolyl radical on the aromatic ring would produce instead methyl-diphenylmethane, another of the by-products of reaction detected. This confirms the existance of radical pathways of reaction inside the Vcontaining zeolite. The products of selective oxidation detected were phenol, benzophenone, phenylbenzoate, benzene, and benzoquinone, besides benzaldehyde and carbon oxides. Benzophenone derives probably from the further oxidation of diphenylmethane [2] and benzene from the decarboxylation of benzoic acid, although the latter was not detected. Benzene also forms by transalkylation as discussed above. The further oxidation of benzene gives rise to benzoquinone. It is more difficult to explain the formation of phenylbenzoate and phenol, the latter furthermore being the main by-product of the reaction. Miki et al. [ 11 ] studying the vapor phase oxidation of benzoic acid on a Ni-Fe-oxide catalyst, observed a significant formation of phenol and suggested a mechanism in which phenyl benzoate is a main intermediate to phenol. The phenyl benzoate forms from an intermediate complex involving two benzoic acid molecules coordinated at a Ni ion. In our case, benzoic acid formation was not observed, but possibly due to its fast consecutive transformation, even if supported vanadium oxide selectively oxidizes toluene to benzoic acid [ 12]. In order to confirm the above mechanism, tests were made feeding benzaldehyde instead of toluene, but maintaining constant all other reaction conditions. The resuits for V3-HMS are shown in Table 2 as a function of the reaction temperature. Table 2 Catalytic behavior of V3-HMS in benzaldehyde oxidation.

Reaction Temp., ~ Conv. BA, % Select. Benzoic acid, % 350 4 65 400 25 51 450 43 29 500 30 3 Experimental conditions as for Fig. 1, but using benzaldehyde (BA) instead of toluene in the feed. Above a temperature of about 450~ the catalyst deactivates due to the formation of heavier products which remain on the zeolite, explaining the lowering of the selectivity above this temperature, but at lower temperatures benzaldehyde selectively transforms to benzoic acid. Phenol and benzene are the main by-products of benzaldehyde oxidation with selectivities up to about 10% each. The other by-products observed in toluene oxidation, including phenylbenzoate, were also detected in benzaldehyde oxidation. The overall reaction network suggested by these experiments is shown in Scheme 1. The higher formation of phenol from toluene on Vx-HMS catalysts, especially those with the higher

900 vanadium content, probably derives from a combination products of alkylation ~ ~ benzophenone of various factors: 0 phenylbenzoate (i) a higher selectivity in benzaldehyde formation, and faster consecutive ,9 / + I-I20 toluene "~COx / benzaldehyde / benzoicacid -, oxidation to benzoic acid, (ii) lower ~".,,, ~ / ',., acidity which cata+COz / + +CO,z lyzes side reactions benzene / and (iii) larger di"'", / phenol benzoquinone "- .................. J mensions of pore Scheme 1. Reaction network in toluene oxidation cavities which allow formation of the complex intermediate responsible for phenylbenzoate formation.

,

' 6e

3.3 Relationship between nature of vanadium species and catalytic behavior The nature of the vanadium species as a function of the nature of the micro- or mesoporous support and the loading of vanadium was studied by diffuse reflectance spectra in the UV5.1, Visible-NIR region 4.:

(UVN-DRS). The resuits obtained are summarized in Fig. 5.

s

4.0 3.5

All the samples are characterized by a FR2.5 LCT (lowest energy 2.(~ charge-transfer) band at about 380 nm 1.5which indicates the 1.(~ presence of nearly o.~ V3-MH _ isolated V s+ ions in a 7'00 distorted octahedral '50 300 200 5'00 ;50 & environment [ 13]. NM Fig. 5 UV-Visible Diffuse reflectance spectra in air of Vx-HMS (1, 3 and However, while the 5% vanadium) and of V3-MFI and V3-1]. intensity of the band is nearly proportional to the content of vanadium in the 0-3% range for the Vx-HCM series, above this value the intensity of the band [F(R) function is the equivalent of absorbance in the transmission mode] is no longer proportional. It may be noted, furthermore, that at equivalent vanadium content, the intensity of the 380 nm band decreases in the order V3-HCM > V3-13 > V3-MFI. In the latter sample a band centered at 480 nm becomes evident. This band may be assigned to an outershell (delocalized) charge transfer and is thus typical of systems in which the hopping of electrons is possible such as in polynuclear V-oxide species [13]. The E of this band is much lower 3.0

.

.....

901 than that of LCT V5+oct.explaing why a lowering of the intensity is noted in V3-13 and V3-MFI with respect to V3-HCM. The presence of V-oxide nanoparticles increases in the order V3MFI > V3-13 > V3-HCM and probably also becomes significant in the Vx-HCM series for loadings of vanadium above about 3%. A second band centred at 260 nm is observed in the UVN-DRS spectra. An LCT band near this value is expected for tetrahedral V 5§ species [13], but higher energy CT bands of VS+oct.species overlap making a clear analysis difficult. However, it should be noted that the ratio of intensities of the bands at about 260 and 380 nm increases in the order V3-MFI > V3-~ > V3-HCM. This suggests that especially in the former sample some tetrahedrally coordinated V 5§ species is present. No evidence was instead found for the presence of V 4§ species (broad bands in the 1000-800 nm region due to d-d transitions). Further indications on the coordination of vanadium in the Vx-HMS series were obtained by 51V wide line NMR spectroscopy. The results are summarized in Figure 6. In the dehydrated state, all samples show a broad signal at about -550 ppm characteristic of tetrahedrally coordinated V(V) species. An additional band is observed at c.a. -290 ppm for V3-HMS, which shows that this sample also contains octahedrally coordinated species, even in the dehydrated state. All samples very quickly rehydrate upon contact with air, which results in the appearance of a broad NMR signal at about -290 ppm.The process is completely reversible and the signal at 290 ppm disappears when samples are outgassed in vacuum at 250~

Hydration

-

/

/

~

"'y 0

9

9 . -200

9 . . . . -400 --600

! -800

0

-

.

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,

.

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,

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6

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rr

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51V wide-line NMR spectra of Vx-HMS samples in the dehydrated (left) and hydrated states (fight).

Based on these data and on the characterization of the acidity of zeolites by NH3-TPD cited previously, but not reported for brevity, it is possible to draw the following conclusions on the relationships between nature of vanadium species, zeolite characteristics and catalytic behavior: (1) When the formation of V-oxide nanoparticles starts, the selectivity to benzaldehyde decreases considerably due to faster consecutive oxidation. Both the loading of vanadium and the nature of the zeolite influence the formation of these V-oxide nanoparticles. The inner surface area of zeolites is very high (up to 1000 m2/g as in the case of Vx-HCM samples; see Table 1), and thus the formation of V-oxide nanoparticles starts much below the vanadium loading necessary for the formation of a monolayer (around 10% wt. for a surface area of 100 m2/g), indicating that a very low dispersion of vanadium was achieved in

902 the zeolites. (2) The presence of acid sites in the zeolite is unfavourable for the selectivity which decreases in the same order, due to three main effects: (i) acid sites catalyze consecutive transformations of benzaldehyde and (ii) toluene gives rise to a variety of by-products, although in low amounts, and (iii) the formation of carbon oxides is enhanced probably favouring the chemisorption of products/intermediates. Medium-strong acid sites able to catalyze these side reactions are created by the addition of vanadium itself, even on the non acidic mesoporous materials. (3) Isolated V5+oct. sites are probably responsible for the selective behavior in toluene oxidation to benzaldehyde as indicated by the characterization of Vx-HMS catalysts. It should be remarked that unsupported or supported vanadium oxides oxidize toluene to benzoic acid [ 12] or a mixture of benzaldehyde and benzoic acid [2], whereas all V-containing zeolites tested form only benzaldehyde. The nature of the zeolite influences the nature of isolated species, as well as the ratio between isolated to polynuclear vanadium species. (4) Phenol forms by consecutive transformation of benzaldehyde probably via benzoic acid which, however, quickly transforms inside the zeolite along different possible routes (Scheme 1). The formation of phenol requires the formation of large intermediates and thus is favoured in mesoporous zeolites. The lower formation of phenol in ~ zeolite than in MFI, however, defives from the faster side reactions in ~ zeolite due to its higher acidity. ACKNOWLEDGEMENTS This work was realized within the framework of an EuRam-Brite Ill project (BRPR-CT950062) financially supported by the E.C., support which is gratefully acknowledged.

REFERENCES 1. C. Batiot and B.K. Hodnett, Appl. Catal. A: General, 137 (1996) 179. 2. J. Zhu and L.T. Andersson, J. Catal., 126 (1990) 92. 3. B. Jonson, B. Rebenstorf, R. Larsson, S.L.T. Andersson and S.T. Lundin, J. Chem. Soc. Faraday Trans. 1, 82 (1986) 767. 4. K. Mori, A. Miyamoto and Y. Murakami, J. Chem. Soc. Faraday Trans. 1, 83 (1987) 3303. 5. Z. Hui-Liang, Z. Wei, D. Xiang and F. Xian-Cai, J. Catal., 129 (1991) 426. 6. A.J. van Hengstum, J.G. van Ommen, H. Bosch and P.J. Gellings, Appl. Catal., 8 (1983) 369. 7. G. Centi and F. Trifirb, Appl. Catal. A: General, 143 (1996) 3. 8. K.R. Reddy, A.V. Ramaswamy and P. Ratnasamy, J. Catal., 143 (1993) 275. 9. S. Gontier and A. Tuel, Microporous Materials, 5 (1995) 161. 10. M.J. Climent, A. Corma, H. Garcia and J. Primo, J. Catal., 130 (1991) 138. 11. J. Miki, M. Asanuma, T. Tachibana and T. Shikada, J. Catal., 151 (1995) 323. 12. J. Miki, Y. Osada, T. Konoshi, Y. Tachibana and T. Shikada, Appl. Catal. A: General, 137 (1996) 93. 13. G. Centi, S. Perathoner, F. Trifirb, A. Aboukais, C.F. Aissi and M. Guelton, J. Phys. Chem., 96 (1992) 2617.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All fights reserved.

903

A novel selective oxidation catalyst: ultrafine complex molybdenum based oxide particles Y. Fan a, W. Kuang a , W. Zhang a'b and Yi Chen a aDepartment of Chemistry, Nanjing University, Nanjing 210093, China bNational Normal College of Inner Mongolia, Tongliao 028043, China

1. INTRODUCTION C-H activation leading to selective oxidation is one of the most challenging problems in terms of surface science and catalysis. In the past decades, many kinds of metal oxides especially Mo-based and V-based oxides have been widely used as selective oxidation catalysts, and the studies on structure and catalytic properties of these oxides have demonstrated that the nature of oxygen species in the oxides is one of the most important parameters influencing catalytic selectivity. For the oxidation of olefins and aromatics the nucleophilic lattice oxygen ions (0 2) are responsible for the selective oxidation, while electrophilic oxygen species (O2,O') may attack C=C and benzene ring because of higher electron density in these regions, which leads to C-C bond cleavage and complete oxidation[ 1-2]. In order to increase the reactivity of lattice oxygen species and to improve the catalytic selectivity, great efforts have been made to modify the state of lattice oxygen species by adding some promoters to these oxides[3]. It has been found that Bi, Fe, Sn, W and rare earth oxides are the effective promoters of the molybdenum based and vanadium based oxide catalysts for selective oxidation of toluene to benzaldehyde[4-9]. Very recently, ultrafine metal oxides have attracted much research interests in terms of materials science and heterogeneous catalysis[10-12]. These new catalytic materials are expected to have unique catalytic properties because of their nano-scale particle sizes. In this work, a novel catalyst for selective oxidation of toluene to benzaldehyde, i.e. ultrafine complex molybdenum based oxide particles, has been developed. It has been found that the reactivity of lattice oxygen ions can be improved by decreasing the oxide particle size to nano-scale and that the ultrafine oxide particles exhibit unique catalytic properties for selective oxidation. Our results have revealed that the ultrafine complex oxide particles are potentially new catalytic materials for selective oxidation reactions.

* Research project supported by the National Natural Science Foundation of China and SINOPEC

904 2. EXPERIMENTAL

2.1 Catalyst preparation The Mo-Ce(sg) oxide particle samples were prepared by a sol-gel method in which the mixed aqueous solutions containing Ce(NO3)y6H20, (NH4)6Mo70246H20 and citric acid with (Ce+Mo)/citric acid =3(tool/tool) and different atomic ratios of cerium to molybdenum were first kept in a water bath at 80~ until gelation was completed, and then the as-prepared gels were dried at 120 ~ for 4h and calcined in air at 400 ~ for 4h. The mono-component MoO3 and Ce02 oxides were prepared by same procedures but without adding cerium nitrate or ammonium molybdate to the mixed solution. For comparison, the conventional coprecipitation method was also used to prepare Mo-Ce(cp) oxide, in which cerium nitrate aqueous solution was mixed with ammonium molybdate aqueous solution, and the precipitates formed were dried and calcined. 2.2 Catalytic oxidation of toluene The toluene oxidation reaction was used as a probe to study the catalytic properties of the Mo-Ce complex oxides. The as-prepared oxides were introduced into a U-type quartz fixed bed microreactor and their catalytic properties for selective oxidation of toluene to benzaldehyde were evaluated under the reaction conditions of 0.1MPa, 400 ~ air/toluene = 9 (vol/vol), F/W =1900 ml/h.g.cat. The reaction products were analyzed by an on-line gas chromatography. Under the above reaction conditions, the main products were CO, CO2, H20 and benzaldehyde. 2.3 Characterization The morphology, particle size and structure of the oxides were determined by using transmission electron microscopy (TEM) and X-ray powder diffraction (XRD). BET surface area of the samples were measured by using Micromeritics ASAP-2000 instrument. The interaction of Ce with Mo and its effect on the nature of active oxygen species in the complex oxides were studied by using temperature-programmed reduction(TPR) and laser Raman spectroscopy (LRS).

3. RESULTS AND DISCUSSION The morphology of Mo-Ce oxide samples (Ce/Mo atomic ratio = 2/3) prepared by different methods are presented in Figures 1 and 2. It can been seen that the particle size of Mo-Ce(sg) was in the range of 20-80 nm, while that of Mo-Ce(cp) was higher than 500nm. These results have shown that the Mo-Ce oxides prepared by the sol-gel method are actually the ultrafine oxide particles(<100nm). The BET surface areas for these two samples were 19m2/g MoCe(sg) and 3m2/g Mo-Ce(cp), respectively. XRD patterns of the above two samples revealed that molybdenum oxide and cerium oxide in the Mo-Ce complex oxide samples formed a solid solution with sheelite-type structure of Ce2(MoO4)3. On increasing Ce/Mo atomic ratio to higher than 2/3, however, Ce2(MoO4)3and CeO2 were found to coexist in the complex oxide samples. Oxidation of toluene was employed to evaluate the catalytic activity and selectivity of the Mo-Ce complex oxides. The particle size, surface area and structure of ultrafine complex

905

Figure 1. Morphology of Mo-Ce(sg) oxide prepared by sol-gel method

Figure 2. Morphology of Mo-Ce(cp) oxide prepared by coprecipitation method

Table 1

The catalytic properties of complex Mo-Ce oxide catalysts (Ce/Mo atomic ratio = 1) for oxidation of toluene MoO3 Particle size(m) >200 Surface area(m2/g) 5.1 Conversion of toluene (mol%) 24.0 Benzaldehyde selectivity (%) 6.0 Benzaldehyde yield (10-6mol/m 2 s) 0.4

Ce02 10-20 41.0 54.8 0.0 0.0

Mo-Ce(cp) >200 12.1 35.5 16.0 0.6

Mo-Ce(sg) 20-40 19.0 34.0 37.0 0.9

Mo-Ce oxide particles were not changed. This suggested that the ultrafine complex oxide particles were quite stable during the catalytic oxidation. After reaction for 7h, the catalytic oxidation reached a steady state. The catalytic properties of CeO2, MoO3 and two Mo-Ce oxides catalysts (Ce/Mo atomic ratio =1) are given in Table 1. It can be seen from these results that under the same reaction conditions, the toluene conversion on the above oxides catalysts showed the order as CeO2 >Mo-Ce(cp)-~Mo-Ce(sg) >MOO3. The oxidation products on mono-component CeO2 catalyst were CO,CO2 and H20 but without benzaldehyde or other selective oxidation products, indicating that CeO2 is an active component for complete oxidation of toluene. By adding ceria to MOO3, however, the selectivity to benzaldehyde was remarkably improved, so that the complex oxides showed higher catalytic selectivity to benzaldehyde than the mono-component MoO3 catalyst. Interesting, the conversions of toluene on both complex Mo-Ce oxides were very similar, but the benzaldehyde

906

Cl

b .I

300

I

500

1

700 Temperature, ~

I

900 ~.

t

1100

Figure 3. TPR profiles of the mono-component MoO3 and the complex Mo-Ce oxide. (a) MoO3, (b) complex Mo-Ce oxide.

selectivity of the ultrafine particle catalyst was much higher than that of the larger particles. These results have revealed that ultrafine oxide particle catalysts have unique catalytic properties for the selective oxidation of toluene to benzaldehyde. In particular, the specific activity (benzaldehyde yield) of the ultrafine complex oxide particles was higher than those of the larger complex oxide particles catalyst and the mono-component MoO3 catalyst, which

907 can not be explained by the effect of mere particle size. Apparently, the differences in the nature of active species in the oxide catalysts should be taken into consideration. As pointed out by Haber[1-2], the lattice oxygen ions in molybdenum oxides are the main active species for selective oxidation of lower olefins to aldehydes. The state of lattice oxygen species in the mono-component MoO3 oxide and Mo-Ce complex oxides were studied by using TPR and LRS. As can be seen in Figure 3, which shows the TPR profiles of MoO3 and Mo-Ce complex oxide, the hydrogen consumption peaks for MoO3 appeared at 670, 755 and 990~ By adding ceria to MOO3, the hydrogen consumption peaks shift to lower temperatures: 510, 715 and 860 ~ These results suggest that due to interaction of Ce with Mo in the complex oxides, the molybdenum oxides are easier to reduce to lower valance. By using IR spectroscopy, Jonson et al.[13] showed that the vibrational frequency of Mo=O in MoO3 was 995cmq. Figure 4 shows the Raman bands of Mo=O species in the complex MoCe oxides. For the complex Mo-Ce(cp) oxide, the vibrational frequency of Mo=O was 953cmq, which is lower than that of Mo=O in MOO3. The red shift of vibrational frequency indicates a weaker bonding between Mo=O in the complex Mo-Ce oxide. This is consistent

953

1100

1000

900

850

Rarnan shift, cmq

Figure 4. Laser raman spectra of Mo=O species in the complex Mo-Ce oxide catalysts prepared by different methods. (a) Mo-Ce(cp), (b) Mo-Ce (sg).

908 with the TPR result. As the lattice oxygen species in the complex Mo-Ce oxides have higher mobility, their reactivity for selective oxidation of toluene are increased. This can account for the higher specific activity of complex Mo-Ce oxide catalysts for oxidation of toluene to benzaldehyde. Moreover, it can be seen from Figure 4 that the vibrational frequency of Mo=O species in Mo-Ce(sg) was much lower than in Mo-Ce(cp), indicating that Mo=O chemical bonding in the ultrafine oxide particles was even weaker and the lattice oxygen ions in ultrafine complex oxide particles have a higher mobility. The higher reactivity of lattice oxygen in the matrix of ultra:fine Mo-Ce complex oxides can account for the reason that ultrafine complex oxide particles exhibit unique catalytic properties for selective oxidation of toluene to benzaldehyde. Our results have clearly confirmed the reactivity of lattice oxygen ions can be improved by decreasing the oxide particle size to nano-scale. These results suggest that the ultrafine complex oxide particles are potentially new catalytic materials for selective oxidation reactions. Reference

1. J. Haber, in Proceedings of the 8th Intemational Congress on Catalysis, July, 1984, Berlin (Wes0, Vol. 1, Verlag Chemie, 1984, P. 85. 2. J. Haber, in Heterogeneous Hydrocarbon Oxidation, ACS Symposium Series 638, American Chemistry Society, 1996, P. 20. 3. Y. Moro-oka and W. Ueda, Adv. Catal., 40 (1994) 233. 4. K. Van der Wide and P. J. Van den Berg, J. Catal., 39 (1975) 473. 5. J. Buiten, J. Catal., 21 (1968) 188. 6. S. Tan, Y. Moro-oka and A. Ozaki, J. Catal., 17 (1970) 125. 7. M. Ai and T. Ikama, J. Catal., 40 (1975) 203. 8. K.A. Reddy and L. K. Doraiswamy, Chem. Eng. Sci., 24 (1969) 1415. 9. Z. Yan and S. Lars T. Andersson, J. Catal., 131 (1991) 350. 10. C. Simon, R. Bredesen, H. Gronadal, A. G. hustoit and E. Tangstad, J. Mater. Sci., 30(1995) 5554. 11. K. Maede, F. Mizukami, M. Watanabe, N. Arai, S. Niwa, M. Toba and K. Shimizu, J. Mater. Sci. Lett., 9 (1990) 522. 12. J. Y. Guo, F. Gitzhofer and M. I. Boulos, J. Mater. Sci., 30 (1995) 5589. 13. B. Jonson, R. Larsson and B. Rebenstoff, J. Catal., 102 (1986) 29.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

909

Liquid Phase Oxidation of Aikylaromatic Hydrocarbons over Titanium Silicalites Georgi N. Vayssilov a, Zdravka Popova a, Stefanka Bratinova b and Alain Tuel c aFaculty of Chemistry, University of Sofia, 1, J. Bourchier avenue., Sofia 1126, Bulgaria e-mail: [email protected] bNational Centre for Environment and Sustainable Development, Sofia 1453, Bulgaria CCNRS, Institut de Recherches sur la Catalyse, 2, avenue A.Einstein, 69626 Villeurbartr, e Cedex, France 1.INTRODUCTION Partial catalytic oxidation of alkylaromatic hydrocarbons is interesting both from the industrial and the scientific point of view. The industrial interest is due to the availability of these substances from the petrochemical industry and to a number of applications for the possible oxidation products. Conventional gas phase oxidation concerns the side chain and leads mainly to benzoic acid or even to destruction of the aromatic ring. Oxidation under mild conditions could cease the reaction at earlier stages and reduce the number of the products formed. However, the appropriate catalyst for such partial oxidation has not been found yet. A promising step in this direction is the oxidation with hydrogen peroxide over titanium silicalite molecular sieves TS-1. This catalytic system was successfully applied in the hydroxylation of phenol, epoxidation of alkenes, and partial oxidation of alkanes and other organic substances [1-4]. The advantages of the TS-1 catalyst reside in the selectivity of the reaction and the long duration of the activity of the catalytic samples [1-5]. The catalytic oxidation of toluene and ethylbenzene using H20 2 over TS-1 was mentioned in one of the initial works on titanium silicalites [ 1]. In both cases hydroxylation of the aromatic ring was observed while only the alkyl chain of the ethylbenzene was oxidised to 1-phenylethanol and acetophenone. So far, only Khouw et al. [6] and Mal et al. [7] reported the quantitative data for ethylbenzene oxidation. Under the reaction conditions used they achieved quite different turnover numbers (TON) - 52 and 8.4, respectively. The product distribution also depends considerably on the reaction conditions. Clerici [8] reported formation of acetophenone and a minor amount of ethylphenols, while the other authors found between 23 and 40 % ethylphenols [1,6,7]. Another difference was the conversion to 1-phenylethanol - from 4% to 33% [6,7]. From the other monoalkylbenzenes, only 2-propyl - and 1-butylbenzene were tested in liquid phase oxidation with H20 2 over TS-1 [8]. The former hydrocarbon was not reactive under the reaction conditions, while the side chain of the latter one was oxidised to corresponding alcohols and ketones at ot and ), positions to the phenyl group. It is clear that there are only occasional studies on the partial oxidation of alkylaromatic hydrocarbons with hydrogen peroxide in the presence of titanium silicalites. This is the first

910 work directed exclusively to the investigation of toluene, ethylbenzene, 1- and 2propylbenzene oxidation over TS-1 and TS-2 catalysts under various reaction conditions. Attention is mainly focused on the evolution of the product distribution (ring hydroxylation vs. side-chain oxidation, alcohol vs. ketone in the side chain) during the initial hours of the catalytic process. 2.EXPERIMENTAL

2.1.Catalyst Titanium silicalite-1 (TS-1) was synthesised following Example 1 of the original patent [9] and titanium silicalite-2 (TS-2), according to ref. [10]. The removal of occluded organics was performed by calcination of the sample at 550~ in air for about 6 h. Chemical analysis of the solid materials (Si/Ti= 50) was obtained by atomic absorption after solubilization of the sample in HF-HCI solution. UV-vis spectrum of the calcinated material showed an adsorption band at about 205 nm, characteristic of titanium silicalites. The absence of signal beyond 275 nm indicated that the material was free from extraframework oxide species. S.E.M. pictures revealed that the sample was in the form of very small uniform crystals of about 0.3 ~tm in size. Before each catalytic run the catalyst sample was heated at 820 K for 5 h.

2.2.Catalytic experiments Hydrocarbons used in catalytic experiments were obtained by Fluka - ethylbenzene and 1propylbenzene, and M e r c k - toluene and 2-propylbenzene. Catalytic experiments were performed in a 50-ml round-bottomed flask equipped with magnetic stirrer and condenser. The initial mixture containing 0.1 g catalyst, 5.0 ml solvent and 2.0 ml substrate was stirred and heated up to the reaction temperature. Then 1.0 ml 30% aqueous solution of hydrogen peroxide (Merck) was added. Small samples (0.1 ml) of the reaction mixture were taken after 1, 2 and 4 h, diluted in 4.0 ml methanol, dried and analysed by GC-MS (Hewlett-Packard). Veratrole was used as internal standard. Hydrogen peroxide conversion was followed by standard iodometric titration. 3.RESULTS AND DISCUSSION

3.1.General observations There are two routes to the partial oxidation of alkylaromatic hydrocarbons - oxidation of the side chain or of the aromatic ring. In both cases the oxidation could proceed at different positions (except for the methyl group of toluene) and to different extents. Titanium silicalites were found to activate the oxidation of the secondary or tertiary saturated carbon atoms, the terminal methyl groups remaining unaffected [2,6-8]. This behaviour limits the number of the

911

CH3

CH3

CH3

CH2OH

HO

CHO

HO-CHCH3 O--CCH3 + ~~/OH

C2H5.......~ .~

~o

~m m o H~~m c H o

HO'cHCH2CH3 O=CCH2CH3 OH CHEC2H5 CH2CHCH3 I

O CHECCH3 II

HO

OH H3c-C-CH3 ~ I

CH(CH3)2 CH(CH3)2 ~ ( C ~ +

" ~ HOHEC-CHCH30HC-CCH3

Scheme 1. Expectedreactionroutesto the partial oxidationof alkylaromatichydrocarbons.

912 expected products from the side-chain oxidation (Scheme 1). The aromatic ring itself could be oxidised at o-, m-, and p- positions with respect to the alkyl group, but due to the electrophilic mechanism of the reaction [4,11,12] the m-isomer can hardly be expected. Although the obtained alkylphenols should be more reactive than the initial alkylbenzenes, their further hydroxylation and oxidation to alkylbenzoquinones is hindered for steric reasons. Since hydroxyl and methyl groups have similar dimensions, the lower conversions of p-xylene and p-ethyltoluene compared to toluene and ethylbenzene [6] suggest the effect of the steric restrictions in catalytic oxidation over TS-1. Products of oxidation both in the ring and in the aliphatic chain, could also be formed in principle. The results of the catalytic experiments follow in general the above assumptions. However, some of the expected products are not observed. Conversions and selectivity in the oxidation of the alkylbenzenes studied are presented in Table 1. In all cases hydroxylation in the aromatic ring is observed, mainly to p-isomers. This p-selectivity is due not only to the substrate molecules used but also to the properties of the solvent molecules. It is known that phenol and anisole hydroxylation in methanol and ethanol exhibits substantial p-selectivity while in non-alcoholic solvents, such as acetonitrile and acetone, o-selectivity is found [5,1315]. The p/o ratio increases with the bulkiness of the alkyl group but, surprisingly, even 2propylbenzene is hydroxylated to some extent at the o-position. Table 1 Conversion and selectivity in oxidation of alkylaromatic hydrocarbon with H202 over TS-1. Alkyl Group

Conversion, %

Oxidation in the ring, %

Side-chain oxidation, %

methyl

1.4

100

0

ethyl

3.5

11

82

1-propyl

2.8

53

47

2-propyl

2.1

100

0

Reaction conditions as described in the Experimental section, solvent- ethanol. Although the intermediates formed during the oxidation of the methyl group in toluene could easily stabilize by resonance with the aromatic ring, the corresponding benzyl alcohol and benzaldehyde are not found. This implies that the simple energetics of the reaction does not determine the absence of terminal methyl group activation over titanium-containing zeolites [6]. Some other metallo-silicate molecular sieves, such as VS-1 and Sn-Sil-1, can activate oxidation of primary carbon atoms under the same reaction conditions [7,16]. Mal et al. [7] suggest that this is due to the different reaction mechanisms over these metal centers via peroxo-radical intermediates over V and Sn and via (electrophilic attack and) carbenium ion formation over Ti. Following this scheme, toluene should be oxidised in the methyl group but, since this is not the case, one should look for another explanation. Simple geometrical factors can also be excluded because Sn4+ and Ti 4+ have similar ionic radii - 71 and 68 pm,

913 respectively [17], and all three catalysts have MFI structure. A difference between these ions resides in the occupation and possibly spatial orientation of the frontal orbitals, especially dorbitals. Another reason could be the relatively higher stability of the doubly reduced Sn2+ and V 3+ metal ions compared to Ti 2+. Side-chain oxidation is observed with ethylbenzene and 1-propylbenzene, while the aliphatic part in 2-propylbenzene is not affected. The lack of side-chain oxidation at the tertiary C atom is probably due to steric hindrance by the two nonreactive methyl groups in addition to the bulky benzene ring. Only the first (benzylic) carbon atom in the ethyl and 1propyl groups undergoes oxidation. The preferential activation of this C atom in the 1-propyl group, compared to the other secondary carbon atom, can be explained by the weak C-H bond at benzylic position and/or by mesomeric stabilization of the intermediate formed (probably carbenium ion [7]). Unfortunately, our investigation cannot answer the hot question why titanium silicalites do not activate terminal methyl groups. However, it clearly shows that although the energetic factors could not explain this peculiarity, they play an important role in the discrimination of different secondary C atoms available for activation. As seen in Table 1, the lower conversion of 1-propylbenzene, compared to ethylbenzene, is due to a lower oxidation of the alkyl chain. While the ratio between the alcohol and ketone formed from 1-propylbenzene is about 1, in the case of ethylbenzene the alcohol represents more than 90 % of the side-chain oxidation product.

3.2.Oxidation of ethylbenzene Ethylbenzene was chosen for the further detailed study of the influence of various reaction conditions (amounts of oxidant, catalyst or solvent (methanol or ethanol), and the framework structure of the catalyst - TS-1 or TS-2) on the conversion and selectivity of the reaction. Ethylbenzene is convenient because the three types of selectivity could be followed (ring vs. side-chain oxidation, p/o ratio in the ring hydroxylation, alcohol/ketone ratio in the sidechain oxidation) and the conversion is the highest. Table 2 shows the reagent concentrations and the product distribution in some of the experiments after 4 h reaction. Turnover numbers found are about one-half of the result of Khouw et al. [6] - TON=52 for 12 h reaction, and higher than the TON=8.4 reported by Mal et al. [7]. The highest conversion is achieved in methanol or without solvent - 5.1 and 4.4 molar %, respectively. However, without solvent the reaction proceeds much faster (Vin= 21 h -1) and after the first hour only 0.4 % of the substratum is additionally oxidised to acetophenone or p-ethylphenol. The higher reaction rate in the beginning and the subsequent interruption of the processes are probably due to the higher temperature - 357 K, while in methanol the reaction was carried out at 342 K. The smaller TON (calculated for ethylbenzene oxidation) and the lower conversion in ethanol is due to the oxidation of the solvent itself, as already reported [18]. The main product of ethanol oxidation is diethylacetal formed by initial oxidation of the alcohol to acetaldehyde and further nucleophilic addition and substitution with other alcohol molecules. Some amount of acetic acid and acetaldehyde are also observed in the reaction mixture. However, the

914 Table 2 Reagent concentrations and product distribution for ethylbenzene oxidation with hydrogen peroxide over TS-1 and TS-2. TS-1 a

TS-1

TS-1

TS-1

TS-2

Mperoxide

mol/1

1.1

0.7

1.1

2.9

1.1

Msubstrate

mol/1

2.1

1.3

2.1

5.6

2.1

ml- 1

3.3

2.1

3.3

8.8

3.3

K

350

349

342

357

350

ethanol

ethanol b

methanol

-

ethanol

19.4

16.7

26.7

24.0

13.9

h- 1

12.2

4.6

11.3

21.1

5.3

1-Phenylethanol

%

77.6

80.3

91.5

83.6

51.6

Acetophenone

%

4.3

11.7

3.8

8.7

7.9

o-Ethylphenol

%

0.8

0.8

0.6

0.4

0.2

p-Ethylphenol

%

10.4

6.7

2.4

6.4

2.8

Others

%

6.8

0.5

1.8

0.8

37.6

Mperoxide/Ti ions Temperature Solvent TON c Initial rate Vind Product distribution

a The standard reaction condition described in the Experimental section. The molar ratio substrate/peroxide=1.9 and substrate/Ti ions=510 are the same in these experiments. b An experiment in diluted solution c TON achieved after 4 h experiment, mol ethylbenzene converted per mol Ti ions d Initial reaction rate during the first hour, mol ethylbenzene converted per mol Ti ions for 1 h. ethylbenzene oxidation in ethanol is a good example that the TS-1 catalyst activates oxidation of hydrocarbons even in excess of alcohol molecules, which are usually much more reactive. Moreover, two-fold dilution of the initial reaction mixture leads to similar TON and conversion as under the standard conditions, described in the experimental section. Of course, in ethanol, the yield with respect to the converted hydrogen peroxide is lower than in methanol (about 5 % in ethanol and 50 % in methanol). Figure 1 shows the evolution of the selectivity to oxidation in the ring and in the alkyl chain during the first hours of the reaction. A very high selectivity to side-chain oxidation is observed in methanol - up to 97 % at the 4th hour. The main product is 1-phenylethanol which amounts to 93 % of the entire conversion. The maximal conversion of ethylbenzene to 1phenylethanol achieved after 4 h reaction is 4.7 % (Fig.2). The ring hydroxylation selectivity

915

100 8O o

o

m Others

60

D Side-chain oxidation

r

40

m Ring oxidation

.,,-4

20

a

b

c

d

Solvents Figure 1. Evolution of the selectivity (1, 2 and 4 h reaction) of ring oxidation, side-chain oxidation and products of deeper oxidation over TS-1. Experiments in ethanol (a), two-fold diluted in ethanol (b), in methanol (c), and without solvent (d).

5-

40

o

-40

30

30 A

z 9 20

20 ._~

o2

.~_

10--

10

O

"~ 1 ~D

>

o0

'

1

I

2

'

I

'

3

i

4

Reaction time, h Figure 2. Time dependence of the ethylbenzene oxidation to 1-phenylethanol over TS-1 in ethanol (rhombus), two-fold diluted in ethanol (squares) and in methanol (triangles).

0

' I ' u ' u ' 1

2

3

4

0 5

Ti ions concentration, mmol/1 Figure 3. TON (open symbols) and initial reaction rate (filled symbols) for ethylbenzene oxidation in ethanol (squares) with variation of the catalyst amount. Rhombus correspond to two-fold diluted solution and triangles to methanol.

916 decreases with time and the portion of acetophenone in the side-chain oxidation product slightly increases. Ring hydroxylation leads preferably to the p-isomer with 88 % selectivity for the first hour and 81% for 4 h reaction. This decrease of the p/o ratio with time is opposite to the trend observed in phenol and anisole hydroxylation over TS-1 [13,15] both in alcoholic and non-alcoholic solvents. Qualitatively, the selectivity of ethylbenzene oxidation in ethanol is the same as in methanol. However, in ethanol the ring hydroxylation selectivity increases up to 12 % (with more than 90 % p-selectivity), and selectivity to 1-phenylethanol decreases, especially for the reaction in dilute solution. The catalytic reaction is performed by varying the peroxide and catalyst concentrations. The TON (Fig. 3) increases using a smaller amount of the catalyst and for 1.0 mmol/l Ti ions (1/4 of the standard amount) TON 35.6 is greater than in methanol (with 4.1 mmol/1 Ti ions). Ring hydroxylation selectivity in the experiments with lower catalyst concentration is lower, compared to the reaction under the standard conditions, but higher than that in methanol. In addition to the lower conversion of the substrate when the amounts of the catalyst and oxidant are reduced, the main difference in these cases is the change in the selectivity to ring oxidation vs. side-chain oxidation (Fig.4). Higher ring-oxidation selectivity is achieved at lower peroxide concentrations - up to 15 % of the product. The factor determining this increase could not be the peroxide concentration in the reaction mixture itself since in the experiment with two-fold diluted reaction mixture the ring-oxidation selectivity is much lower. Probably the important factors for ring hydroxylation are the ratio between peroxide and titanium concentrations or the changes of the solvent properties due to the smaller amount of water in the mixture. 15

,.

~2 0 0

..

.

I I I I I I 2 4 6 8 10 12 Peroxide concentration/Ti ions, 1/ml

'~ 14

Figure 4. Selectivity to ring oxidation in the experiments with variation of the peroxide concentration per Ti ion (1/ml) for 1 h reaction (rhombus) and 4 h reaction (triangles). The TS-2 catalyst shows a lower total conversion than TS-1 with the same titanium content but the selectivity to side-chain oxidation and especially to acetophenone is much higher - up to 35 % of the product after 1 h reaction. While 1-phenylethanol concentration increases with

917 time, the amount of acetophenone remain the same. This is related to the substantial increase of the products of further oxidation of acetophenone (see "others" in Table 2) - Fig.5. A small amount of such products is observed also over the TS-1 catalyst. 1.0 0.8 ~J

U 1-Pheny lethan~ ~

0.6

o

l Acetophenone

r/) ~D ""

o

0 .4

U Others

0.2 0.0 lh

2h

4h

Figure 5. Conversion of ethylbenzene to 1-phenylethanol, acetophenone and products of deeper oxidation (others) over TS-2 catalyst after 1, 2 and 4 h reaction. The deeper oxidation of ethylbenzene over TS-2 can be explained with the slower diffusion of 1-phenylethanol and acetophenone formed in the zeolite pores where they could undergo additional oxidation to acetophenone or other products, respectively. Another possible reason could be some differences in the local geometry of the titanium sites due to the different framework structure of the two titanium silicalites. In addition to the above mentioned products, traces of benzaldehyde are detected. The amount of benzaldehyde increases in the course of the reaction and is higher over TS-2 catalyst. It is probably formed after a C-C bond break in the side chain of 1-phenylethanol or acetophenone. Another by-product - ethylbenzoquinone, is observed in the experiments without solvent and in ethanol under standard conditions. The ethers derived from 1phenylethanol and the solvent- methanol or ethanol, are also found.

4.SUMMARY Titanium silicalites TS-1 and TS-2 catalyze hydroxylation in the aromatic ring of the monoalkylbenzenes studied to corresponding alkylphenols, using hydrogen peroxide as oxidant. Para-isomers are mainly formed in methanol or ethanol as solvents. In the case of ethyl- and 1-propylbenzenes, the first carbon atom of the aliphatic chain is also oxidised both to alcohols and ketones. As expected, the terminal methyl groups in all hydrocarbons are not oxidised. The probable reasons for this behaviour of titanium silicalites are discussed. The influence of the reaction conditions on the conversion and selectivity is studied by ethylbenzene oxidation. Methanol leads to higher conversion than ethanol under the same

918 reaction conditions and after 4 h the reached value is 5.1% with TON 26.7. In both solvents the reaction is selective to side chain oxidation. The highest selectivity to ring hydroxylation (15 %) is found in ethanol, at low peroxide concentration per unit titanium ion. The molar ratio of the alcohol to ketone formed is higher than 9 over TS-1 while over TS-2 this ratio is 1.5. The decrease of the amounts of oxidant and catalyst leads to an almost proportional decrease of the conversion to all products. Suggestions concerning the explanations of the observed correlations are made. Liquid phase catalytic oxidation of ethylbenzene with hydrogen peroxide over TS-1 molecular sieves is most appropriate for the production of 1-phenylethanol with high selectivity (up to 93 % of all the oxidation products in methanol) under the reaction conditions studied here. An additional increase of the 1-phenylethanol selectivity could be achieved with smaller amounts of the catalyst. The highest conversion to acetophenone is found over TS-2 zeolites but further oxidation easily takes place in this case.

Acknowledgement This work was supported in part by the Bulgarian National Science Foundation.

REFERENCES 1.U.Romano, A.Esposito, F.Maspero, C.Neri, M.G.Clerici, Chem.&lnd 72 (1990) 610 2.D.R.C.Huybrechts, L.de Bruycker, P.A.Jacobs, J.Mol. Catal. 71 (1992) 184 3.S.Gortier and A.Tuel, Appl.Catal.A 118 (1994) 173 4.B.Notari, Adv.Catal. 41 (1996) 253 5.G.N.Vayssilov, M.Yankov, Z.Popova, L.Dimitrov, Proc. 15th Conference on Cataysis in Organic Reactions, Phoenix, USA, 1994, p.443 6.C.B.Khouw, C.B.Dartt, J.A.Labinger and M.E.Davis, J. Catal. 149 (1994) 195. 7.N.K.Mal and A.V.Ramaswamy, Appl. Catal. 143 (1996) 75. 8.M.G.Clerici, Appl. Catal. 68 (1991 ) 249 9.M. Taramasso, G. Perego and B. Notari, US Pat 4 410 501 (1983). 10.J.S. Reddy and R. Kumar, Zeolites 12 (1992) 95. 11 .J.S.Reddy, S.Sivasanker, P.Ratnasamy, J.Mol. Catal. 71 (1992) 373. 12.G.Bellussi and M.S.Rigutto, Stud.SurfSci.Catal. 85 (1994) 177. 13.A.Tuel, S.M.-Khouzam, Y.Ben Taarit, C.Naccache, J.Mol. Catal. 68 (1991) 45. 14.J.A.Martents, Ph.Buskens, P.A.Jacobs, A.van der Pol, C.Ferrini, H.W.Kouwenhoven, P.J.Kooyman, H.van Bekkum, Appl. Catal. A 99 (1993) 71. 15.G.N.Vayssilov, Z.Popova and A.Tuel, Chem.Eng. Technol., in press 16.P.R.Hari Prasad Rao, A.V.Ramaswamy, P.Ratnasamy, J. Catal. 143 (1993) 275. 17.F.A.Cotton and G.Wilkinson, Advanced Inorganic Chemistry, Wiley-Interscience, New York, 1972, vol.II. 18.F.Maspero and U.Romano, J.Catal. 146 (1994)476.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

919

Coupled Vanadyl Centres in Vanadium Phosphorus Oxide Catalysts: Essential Structural Units for Effective Catalytic Performance in the Ammoxidation of Methylaromatics A. BrOckner *, A. Martin, B. LOcke and F. K. Hannour

+

Institut lhr Angewandte Chemie Berlin-Adlershof e. V., Rudower Chaussee 5, D-12484 Berlin, Germany

ABSTRACT The ammoxidation of substituted toluenes over differently prepared (NH4)2(VO)3(P207)2and VO(PO3)2-phases as well as over (VO)2P207 has been studied by catalytic and in s i t u - E S R measurements. For effective catalytic performance at least two structural properties were found to be essential: i) Closely neighbouring VO 2+ centres must be exposed at the surface which enable the simultaneous adsorption and conversion of the substrate and ii) the catalyst structure must contain building blocks of exchange-coupled VO 2+ ions e. g. in the form of chains or layers which support the electron transport during the redox process.

I. INTRODUCTION In addition to a variety of transition metal oxide catalysts also vanadium phosphorus oxides (VPO) proved to be suitable catalysts for the ammoxidation of substituted toluenes to the corresponding benzonitriles [1 ]. Although numerous papers on the mechanism of the ammoxidation reaction have been published in the past the various reaction steps, the nature of the active sites and, in particular, structural properties required for effective catalytic performance are still ambiguous. I n s i t u - i n v e s t i g a t i o n s under working conditions should lead to a better understanding of the relation between structural peculiarities and catalytic properties. However, only few methods such as electron microscopy [2], Raman [3] and FTIR spectroscopy [4] have been developed to perform such measurements so far. In this study we demonstrate how m situ - electron spin resonance (ESR) can be used to study the behaviour of unsupported VPO catalysts during the ammoxidation of substituted toluenes in a special self-constructed in situ - ESR flow reactor. From the results, conclusions on the arrangement of structural units required for effective catalytic performance were derived.

+

On leave from the Department of Chemical Engineering, University of Technology of Compi6gne, France

920 2. EXPERIMENTAL

2.1. Catalysts (VO)2P207 was prepared by calcining the precursor VOHPO4 0.5 H20 in nitrogen atmosphere for 2 hours at 753 K. The precursor was obtained by evaporating an aqueous solution ofV2Os, H3PO4 and oxalic acid as described elsewhere [5]. Pure crystalline a-(NH4)2(VO)3(P2002 (NVP~) was prepared by calcining a mixture of V205 and (Nt~)2HPO4 (V:P:N - 1:3:6) at 573 K in air [6]. An equilibrated catalyst (N ,o) consisting of ~- 75 mol-% a-(NI-hh(VO)3(P207h and = 25 mol-% of an amorphous vanadium oxide phase was obtained after treating VOHPO4 0.5 H20 under ammoxidation conditions for 40 h [e. g. 6]. Two further catalysts were prepared by mixing NVP~ mechanically with solid NH4VO3 (NVPm) and by impregnating NVP~ with an aqueous solution of NH4VO3 (N i) both in a molar ratio of NVP~/NH4VO3 = 1 followed by 3h formation under ammoxidation conditions. Two VO(PO3)2-samples were obtained by calcining the precursor VO(H2PO4)2 [7] in air at 673 K for 6 h (amorphous phase) and at 1073 K for 1 h (crystalline t-phase). Crystalline aVO(PO3)2 was prepared by heating V205 and H3PO4 at 643 K for 5 h [8]. As a reference sample, V409 was prepared by heating a finely divided stoichiometfic mixture of VzO5 and sulfur at 453 K for 65 h in an autoclave [9]. The vanadium valence state was determined by potentiometric titration [ 10]. 2.2. Methods Catalytic activities and selectivities were determined separately in a fixed-bed U-tube quartz-glass reactor (8 mm i. d.) under similar conditions. Toluene conversion and benzonitrile formation were followed by on-line GC using a FID detector. The carbon oxides were measured by non dispersive infrared photometry. ESR spectra were recorded with the cw spectrometer ERS 300 in X-band (Zentrum fttr WissenschafUichen Geratebau Berlin). For in situ - measurements, i. e., investigation of the catalyst under reaction conditions in the cavity of the spectrometer, a self-constructed ESR flow reactor equipped with a bifilar heating winding of Pt wire and connected to a gas/liquid supplying system was used [11]. In each run 400 mg catalyst particles (0.5 - 1 mm) were treated with a gas flow of 980 ml h"~ (molar ratio : aromatic hydrocarbon : air : NH3 : H20 = 1 931 94.5 924.3) at atmospheric pressure. 3. RESULTS AND DISCUSSION Suitable VPO-catalysts for the ammoxidation of toluene can be obtained on dehydrating the precursor VOHPO4 0.5 H20. When this dehydration is done in inert atmosphere crystalline (VOhPzO7 is obtained as the only product. However, when the precursor is calcined under the ammoxidation feedstock (air, NH3, hydrocarbon, HzO vapour) a catalyst is formed which contains crystalline (NH4)2(VO)3(P2OT)2 comprising about 75 % of the overall vanadium content along with an amorphous vanadium oxide phase. Both catalysts were found to catalyze the ammoxidation of toluene effectively [e. g. 1]. However, in the latter case it was not clear whether crystalline (NI-h)2(VO)3(P2OT)2 or the additional amorphous phase is the catalytically active component. Therefore, in addition to pure crystalline (NI-h)2(VO)3(P2OT)2 two further samples were tested catalytically in which NVP~ had been modified by mixing with solid NH4VO3 and by impregnating with an aqueous solution of NH4VO3, respectively, to generate an additional amorphous vanadium oxide phase as formed in NVP,o, too. In the following

921 sections the results of catalytic and ESR-spectroscopic investigations of these catalysts and three different VO(PO3)2 samples are described.

3.1. Catalytic Results The benzonitrile selectivities for the various catalysts proved to be rather similar. They varied slightly between 80 % and 100 % depending on the degree of conversion. However, marked differences were found in the activities. From the plot of the area-specific toluene conversion as a function of reaction temperature (Fig. 1) it can be seen that (NH4h(VO)3(P2OTh catalysts are most active when they contain an additional amorphous vanadium oxide phase generated either by in situ-formation during catalysis or by treatment with NH4VOa while pure crystalline (NH4)2(VO)a(P207)2 is rather inactive. This suggests that not the crystalline but the amorphous phase in (NH4)2(VO)a(P207)2-based catalysts is the catalytically active 400 componem. (VO)2P207 is a good catalyst, too, even though E 300not as active as the three e-

"~

200-

~

1000 6O0

(Nn4)2(VO)3(e2OT)2-basea catalysts. In the case of vanadyl

I

620

I

640

I

660

I

680

I

700

I

720

I

740

T/K

Fig. 1 Area specific toluene conversion, Cto~,for NVPi (I),

NVPm (#), NVPao (0), NVPsyn (A), (VO)2P207 (O), a(A), fl- ( <>) and amorphous VO(POa)2 (r'l)

polyphosphates it is interesting to note that amorphous VO(POa)2, in particular at higher temperatures, is markedly more active as the respective crystalline a- and t3VO(PO3)2. To get more information on the reasons of the different catalytic behaviour the catalysts were investigated under reaction conditions using in situ-ESR.

3.2. ESR-Investigations The ESR spectra of all samples consist of one single isotropic line. It arises from the VO 2§ centres of the structure and is narrowed due to effective spin-spin exchange interactions between them. For the same reason, the anisotropic g and his tensors are not resolved. The efficiency of this spin-spin exchange depends strongly on the arrangement of the neighbouring VO 2§centres within the catalyst structure as well as on changes of their electron density caused by the catalytic reaction. As we demonstrated recently [12], evaluation of the temperature dependence of the ESR signal intensity and the line shape leads to two parameters which can be used to characterize quantitatively the exchange efficiency: i) the exchange energy, AE, and ii) the quotient of the 4th and the square of the 2nd moment of the ESR signal, /2 Both parameters have been used in this work to analyze structural differences between the various catalysts and their behaviour under working conditions. In the crystal structure of the precursor, VOHPO4 0.5 H20, the vanadyl ions are incorporated as exchange-coupled dimers of VO6 octahedra. During the dehydration in inert or ammoxidation atmosphere, respectively, these dimers are cleaved and the VO6 octahedra rearrange to form infinite ladderlike double chains in (VO)2P207 (formed in inert atmosphere) and single chains in (NH4)2(VO)a(P207)2(formed in ammoxidation atmosphere). When this trans-

922 formation is followed by in situ-ESR [12, 13] a strong line broadening is observed at the transition temperature which is caused by the temporal collapse of the spin-spin exchange. The moment quotient, /2, passes a minimum before it increases again when new exchange pathways are established during the formation of the catalyst structures. Thus, 2 is a sensitive parameter to study structural differences in terms of the exchange efficiency. The catalytic results described above revealed that the catalytic activity of (NH4)2(VO)3(P2OT)2-based catalysts is governed by the amorphous vanadium oxide phase. This raises the question on the structure and composition of this component. From the in situ-ESR measurements under working conditions it can be deduced that the amorphous phase contains VO 2§ ions coupled by effective spin-spin exchange interactions. Catalysts containing this phase (NVP~o NVPi and NVPm) have considerably higher moment quotients than NVP~ (Fig. 2). 18 16

14 on v

12 10

4~

8

ix) v

6

NVP m

NVPi

NVPao

NVPsyn

Fig. 2 Moment quotient of the ESR signals of various (NH4)2(VO)a(P2OT)2-based catalysts under ammoxidation conditions (NH3, H20, air) in the absence (black bars) and in the presence of toluene (white bars)

140 120 100 I--

~:

8o-

uJ

60-

J,

4020-

V409

II

NVPsyn

I

I

V409 / NVPsyn I

NVF)ao

49 m i x t u r e

Fig. 3 Exchange energy, AE, derived from the temperature dependence of the ESR signal intensity

923 The mean vanadium oxidation state in equilibrated (NHa)2(VO)3(P2OT)2-based catalysts is 4.12 [6]. That means half of the vanadium ions in the amorphous phase should be pentavalent. Taking also into account the stoichiometry of the dehydration reaction the following equation can be written formally: 16 VOHPO4 0.5 H20 + 8 NH3 + 0.5 02 --) 4 (NH4)2(VO)3(P2OT)2 + V409 +12 H20

(1)

It indicates that the amorphous phase might be a mixed va+/v 5§ oxide of the type V409. We have prepared crystalline V409 containing equal amounts of vanadium in the oxidation state +4 and +5 as a reference sample. The VO 2+ ions were found to be coupled by strong spin-spin exchange interactions. Evaluation of the temperature dependence of the ESR signal intensity resulted in a coupling energy of AE = 127 cm~. This value is markedly higher than that derived for NVP~ and also for NVPao (Fig. 3). However, when as-synthesized V409 is mixed with NVP~ in a molar ratio of 1:4 as required by Equ. 1 the coupling energy of this mixture is of the same order of magnitude than that of the catalyst formed under reaction conditions (NVPao, Fig. 3). Thus, it seems likely that the amorphous phase in the latter has a similar structure and composition as V409. In accordance with previous investigations [12], the ESR experiments under working conditions show further that the spin-spin exchange in the catalyst structures is disturbed as soon as the aromatic reactant is added to the feed indicated by a decrease of the moment quotient. This effect is most pronounced for active catalysts (Fig. 2). It is caused by changes of the electron density at the surface vanadyl centres involved in the ammoxidation process having two reasons: i) their reduction and re-oxidation during the catalytic reaction and ii) the adsorption of the aromatic ~-system at the VO 2+ centres which transmits additional electron density to the latter. The extent of the exchange perturbation depends strongly on the basicity of aromatic zt-system which, for its part, can be influenced by a second substituent. In Fig. 4 the decrease of the moment quotient caused by the adsorption of the respective substituted toluene is plotted as a function of the Hammet constant of the second substituent. 100 -

~

__.JI []

A r

m v

~

t'N ~

v

8o

CH 3

m v

,A ,~ m

p-CI -

rn

7o

v

60

,..-.,...., -0.3

-0.2

.... -0.1

Hammer

, .... 0

, .... 0.1

, .... 0.2

, 0.3

constant, o

Fig. 4 Ratio of the moment quotients in the presence (/2HC) and in the absence of CH3-CrH4-R (/20) in dependence on the Hammett constant of R, for NVP~ (0), NVPi (u) and NVPm (e) under ammoxidation conditions (NH3, H20, air)

924 For all catalysts investigated a linear correlation results. The perturbation of the spin-spin exchange is lowest for electron withdrawing substituents weakening the basicity of the ~system and increases for electron donating substituents. Besides the effect of various substituents the influence of temperature on the interaction between reactant and catalyst has also been studied by m situ-ESR. In Fig. 5 the moment quotient of NVP~ during the ammoxidation of 4-methoxy-toluene is plotted as a function of temperature and reaction time. At 573 K no conversion takes place and only a slight decrease of the moment quotient is observed in the presence of the hydrocarbon indicating a weak adsorption. At 673 K the reaction is just beginning (degree of conversion: 1%). Treatment with 4-methoxy-toluene leads to a strong decrease of the moment quotient. After removing the latter from the feedstock the low level of the moment quotient is retained for about 1 h before it starts returning to the initial value. This points to a strong adsorption of the hydrocarbon which is blocking the surface. At 713 K (degree of conversion: 20 %) the perturbation of the spinspin exchange is lower again because by ammoxidizing the methyl group to a nitrile group an electron withdrawing substituent is introduced into the aromatic molecule which favours its desorption from the surface. 110 r.,-i

N H 3 + 1"!20 + 0 2

', + H n ',

100

',+ He'.:

.___.

90

O "1"

o

A O4 nn V

A t'~ nn V

A

A

v

v

Tn v

80 70 60

v

50

0

100

200

300

400

500

Reaction time / min

Fig. 5 Ratio of the moment quotients in the presence (/2c) and in the absence of 4methoxy-toluene (/20) in dependence on the reaction tine at 573 K([---I), 673 K (I) and 713 K ~ ) for NVP~ under ammoxidation conditions (NH3, H20, air) In addition to the catalysts obtained from the precursor VOHPO4" 0.5 1-120 also three vanadyl polyphosphates have been studied by examining the spin-spin exchange efficiency as a function of feed composition at a reaction temperature of 723 K (Fig. 6). For as-synthesized crystalline a- and fl-VO(PO3)2 (Fig. 6a, b) virtually no changes of the moment quotient were detectable. In contrast, a marked increase is observed for the as-synthesized amorphous VO(PO3)2- phase (Fig. 6c) as soon as the inert nitrogen atmosphere is replaced by a mixture of air and ammonia. Obviously, the surface of the catalyst is modified during the interaction with feed components. Probably, a layer is formed in which the exchange interaction between neighbouring vanadyl centres is enhanced in comparison to the untreated sample. This is in agreement with the findings for amorphous VO(PO3)2 having been used in the ammoxidation

925 r

o3

(D

o3

20"1"

."~

04 Z

/

15-

1"4"0

I

~

u.,-."l" , m +

i '

~ +

I I

/

9

10-I-

9

eO

9

'

'

i

n

!

I I

I I

A

m V A ~1" m V

9

I

, I

i

I I I I I I I

10

I

9

I I

or o i 9 t - o O ~ O

,

I

I

r

I

i r ' Z

i

J

I I

I I

0

a

,

12

r

4 . 0

i ,_-1''~u

O I

I --n w4l

t

I I

I

,

9

On

0

i

I I

I I

10 O-I I O I I I

I

4

I I I I I I I I I I

'" I I I I I I I I I

_lO

0

I I I 9

i O

9

000

9

0000

b

~ I

8 10-

9

%,.,,.-_ '._o ,. ;~ 1,,-1~,,..r 0 4 . w . = ~ 0 . 0 0 . 6 0 0 , 0 . 0 . 0 . ~-.---I

n

I I I I I

I I I I I

I

i

I

I

I I I I I

I I i I I

i

I I I I I

I

,

9

9

9-

!

8

7 ~o I i

I

1

C

I

n

u

~

I

I

I

I I !

! ! !

! ! I

I !

! !

! !

I I

I I

I !

I !

! !

! i

!

I'

2

3

'!

4

'1

5

6

Reaction time / h at 723 K

Fig. 6 Moment quotient of the ESR signals of as-synthesized (black points) and conditioned catalysts (white points): a-VO(PO3)2 (a), p-VO0~3)2 (b), amorphous VO(PO3)2 (c) in dependence on feed composition and reaction time at 723 K

926 reaction for 10 h at 723 K (Fig. 6c). In this catalyst the initial value for the moment quotient is markedly higher in comparison to the as-synthesized sample. The surface layer formed during the equilibration might be the reason for the enhanced catalytic activity of amorphous VO(PO3) in comparison to crystalline 5- and fl-VO(PO3)2 (Fig. 1) in which only a slight increase of the exchange efficiency after the equilibration process has been observed (Fig. 6a, b). While the addition of water to the feed does not cause significant changes, the exchange efficiency decreases slightly in the presence of toluene (Fig. 6). The reasons therefore are the same as discussed above for (VO)2P207 and various (NHa)2(VO)3(P207)2 - phases. It is interesting to note that for the less active crystalline a- and fl-VO(PO3)2 phases this effect is weak and occurs only in the conditioned samples (Fig. 6a, b). In the case of amorphous VO(PO3)2 showing higher activity the perturbation of the spin-spin exchange in the presence of toluene is more significant even in the flesh sample (Fig. 6c). From the results described above it becomes evident that differences in the catalytic performance of the VPO-catalysts investigated, among other reasons, might be related to the spinspin exchange behaviour of the vanadyl ions. In the following section the structure and the catalytic behaviour of the various samples are compared in relation to the reaction mechanism. In conclusion, structural units being essential for effective catalytic performance are identified. 3.3. S t r u c t u r e a n d Catalytic P e r f o r m a n c e

In analogy to oxidation reactions over VPO catalysts [14], the ammoxidation, too, is assumed to proceed with the participation of lattice oxide ions including the alternating reduction and re-oxidation of surface vanadyl ions according to a Mars-van-Krevelen mechanism [12] 9In the first step the toluene adsorbs with the aromatic vr-system at a VO 2+ centre of the surface. According to a reaction scheme published previously [4, 12] the conversion of the methyl into the nitrile group proceeds via a number of subsequent redox steps involving another VO 2+ centre in close neighbourhood to the adsorption site. To favour the simultaneous adsorption and conversion of the substrate the catalysts should contain VO 2+ centres which are neighboured closely enough to match the distance between the ~r-complex centre and the methyl carbon atom of the toluene molecule which can be estimated to be ~ 3 A from the respective C-C~g (1.4 A) and C-CH3 bond length (1.52 A) [ 15]. In addition, the above mentioned redox process at the surface vanadyl ions leads to a fluctuation of their electron density. When they are incorporated into units of exchange-coupled VO 2+ centres the alteration of the electron density can be rapidly delocalized over a certain range of the structure. We assume that in this case the electronic distortion at a discrete surface vanadyl centre is diminished and the activation energy is lowered. Thus, effective exchange pathways within the catalyst structure appear to be another requirement for effective catalytic performance. Both requirements are fulfilled for (VO)2P207 and also for V409Which is assumed to be similar to the active component in (Nn4)2(VO)3(P207)2-based catalysts either treated with NI-I4VO3 or formed under working conditions from VOHPO4 0.5 H20 (Fig. 7). In contrast, the structures of crystalline 5-(NH4)2(VO)3(P207)2, as well as of 5- and ]~-VO(PO3)2 do contain exchange-coupled single chains of VO6-octahedra, but their distance is too large to enable the adsorption of the aromatic ~r-system and the ammoxidation of the methyl group in neighbouring sites (Fig. 7). Probably, this is the reason why these catalysts show rather poor activity. We suppose that the higher catalytic performance of amorphous VO(PO3)2 is caused by a suitable arrangement of VO 2+ ions within the active layer formed during conditioning.

927

H H-C--k""(~~.,,3'

H!

!

9

! !

I ! |

(VO)2P207

V409 H

H I |

a-(N H4)2(VO)3(P207)2

! |

/~-VO(PO3)2

Fig. 7 Schematic view on the crystal structures of various catalysts and the adsorption of toluene (white units; phosphate groups, grey units: vanadyl groups)

928 4. CONCLUSIONS Comparison of the catalytic results with the spin-spin exchange behaviour investigated by in situ-ESR and the structural properties of the catalysts led to the following conclusions:

9 Obviously the catalytic activity is governed by efficient exchange pathways as provided by chains or layers of VO6 octahedra. 9 Vanadyl distances must be close enough to allow simultaneously the adsorption of the ~rsystem and the ammoxidation of the methyl group in neighbouring sites. Catalyst structures with isolated single chains of VO6 octahedra (e. g. crystalline VO(PO3)2) are markedly less active than those with layers or double chains of vanadyl centres (e. g. (VO)2P2OT). The reason is probably that the latter materials enable both the adsorption of the aromatic ~r-system and the ammoxidation of the methyl group. In contrast, catalysts with single chains do support the adsorption but not the ammoxidation due to the absence of a seeond closely neighbouring vanadyl site.

ACKNOWLEDGEMENT This work was supported by the Federal Ministry of Education, Research and Technology of the FRG and the Berlin Senate Department for Science, Research and Culture (project No. 03C 3005).

REFERENCES

1. A. Martin, B. LOcke, H. Seeboth, G. Ladwig and E. Fischer, React. Kin. Catal. Lett., 38 (1989) 33. 2. P.L. Gai and K. Kourtakis, Science, 267 (1995) 661. 3. F. Ben Abdelouahab, R. Offer, N. Guilhaume, F. Lefebvre and J. C. Volta, J. Catal., 134 (1992) 151. 4. Y. Zhang, A. Martin, H. Berndt, B. L0cke and M. Meisel, J. Mol. Catal. A, in press. 5. K. Sehlesinger, M. Meisel, O. Ladwig, B. Kubias, R. Weinberger and H. Seeboth, DD WP 256659 Al, 23.10. 1984, Zentralinstitut fi~r Anorganische Chemie, Berlin. 6. Y. Zhang, A. Martin, G.-U. Wolf, S. Rabe, H. Worzala, B. LOcke, M. Meisel and K. Witke, Chem. Mater., 8 (1996) 1135. 7 G. Ladwig, Z. Chem., 8 (1968) 307. 8. B.C. Tofield, G. R. Crane, G. A. Pasteur and R. C. Sherwood, J. Chem. Soe. Dalton Trans., 18 (1975) 1806. 9. A. Hammouche and A. Hammou, Electrochimiea Acta, 32 (1987) 1451. 10. M. Niwa and Y. Murakami, J. Catal., 76 (1982) 9. I 1. A. Br0ckner, B. Kubias and B. LOcke, Catal. Today, 32 (1996) 215. 12. A. Br~ekner, A. Martin, N. Steinfeldt, G.-U. Wolf and B. LOcke, J. Chem. Soe., Faraday Trans., 92 (1996) 4257. 13. A. Br0ckner, B. Kubias, B. L0cke and R. StOl3er, Colloids and Surfaces A: Physieochem Eng. Aspects, 115 (1996) 179. 14. E. Bordes in R. W. Joyner and R. A. van Santen (eds.), Elementary Reaction Steps in Heterogeneous Catalysis, Kluwer Academic Publishers, 1993, p. 137. 15. Handbook of Chemistry and Physics, 75th Edition, CRC Press, Boca Raton, 1995.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

929

Ammoxidation of xylenes- kinetics and selectivity Klaus Beschmann ", Stephan Fuchs b and Thomas Hahn b 9 aMerck KGaA, Darmstadt, Germany bInstitut f/ir Chemische Verfahrenstechnik der Universit~t Karlsruhe, D-76128 Karlsruhe, Germany Shape selectivity is in many cases not a result of sharp molecular sieving but of differences in intracrystalline mobility, due to steric hindrance. For example the para-isomers from alkylation followed by secondary isomerization inside zeolite crystals will preferably escape to the gas phase. In principle the direction of diffusion processes is reversible. For this reason one can imagine an educt selectivity as reversion of the product selectivity described above. In this case from a mixture of isomers the reaction of the para isomer should be favoured. The kinetics and selectivity of ammoxidation of xylenes was investigated on modified ZSM-5 crystals of different size and composition. Copper modified Na-ZSM-5 catalysts had the highest activities and selectivities. A simple kinetic scheme of first order reactions can describe the experiments in a quantitative manner. The rate parameters for the reaction of p-xylene were also valid for the reaction of an equimolar mixture of meta- and para-xylene. Educt-selectivities of 90% independent of conversion range can be observed in the latter case. 1. I N T R O D U C T I O N Shape selectivity was first described by Weisz and Frilette [1] and explained as a molecular sieve effect leading to a sharp screening of either reactants or products according to molecular size and shape. This concept was then elaborated in experiment and theory by Weisz et al. [2]. Chutoranski and Dwyer [3] described a more subtle form of shapeselectivity based on the observation that kinetics of xylene isomerization in a zeolite catalyst depends on the size of the crystals and can be represented by a model, assuming different diffusivities of the xylene isomers in the crystals affecting kinetics and product composition. In this case none of the different species are excluded entirely from entering or leaving the zeolite crystal. A concise theory of enhanced para-xylene selectivity in the methylation of toluene was presented by Wei [4] in a model representing the kinetics of simultaneous intracrystalline alkylation and subsequent isomerization of xylenes assuming linear kinetics and different diffusivities of xylene isomers in the zeolite crystal. Therefore a variation of crystal size alone can influence theproduct composition at low conversion. In a substantial experimental study Beschmann and Riekert [5] examined the influence of *correspondence author

930 crystal size and composition on product distribution for the isomerization of xylenes and for the methylation of toluene in a wide range of methanol conversion. They discussed their results with the description of Wei [4] and showed that alternative explanations in the literature for the observed product distribution can be implemented in the model of Wei as extreme cases. In principle the direction of diffusion processes is reversible. For this reason one can imagine an educt shape selectivity as reversion of the product shape selectivity described above. In this case the reaction of the pars-isomer in a mixture of xylene isomers should be favoured. In view of the fundamental interest in diffusion mediated shape selectivity, we have investigated the kinetics and selectivity of ammoxidation of xylenes to initial products of terephthalic acid. CeH4(CHs)2 + NHs + 3/2 02

CHsCeH4CN + 3 H20

(1)

CHsCeH4CN + NHs + 3/2 02

CeH4(CN)2 + 3 H20

(2)

Terephtalic acid can be obtained by hydrolysis of the Dinitrile (DN) followed by thermal cracking (Lummus process) [6]. 2. E X P E R I M E N T A L

2.1. Catalysts Three zeolites ZSM-5 with different Si/A1 ratios and crystal sizes L (table 1) were modified by ion exchange with 0.005 M metal acetate solutions (Co, Ni, Pd, Cu, Ag, Zn) for 12h at 300 K. The zeolites were washed and dried at 393K for lb. After calcination at 823K in air for 6h, the powders were pressed into pellets (1900 bar) without binder. The pellets were crushed and size fractions of 0.3 to 0.5 mm were obtained by sieving and used as catalysts.

Table 1 Zeolites (characteristic length L = Volume/outer surface) Designation Si/A1 Crystal size (pm) H-ZSM-5 24 0.1 - 0.20 H-ZSM-5 24 20- 50 Na-ZSM-5 37 0.1- 1

L (#m) 0.03 4.16 0.08

2.2. Observation of kinetics Reaction kinetics and product distribution were observed in a small isothermal quartz glass reactor (diameter 8mm) with plug flow characteristics (Fig. 1). The educt gas mixture is adjusted with thermal mass flow controllers. The N2/O~ mixture can be saturated with xylene or mixtures of xylene isomers to a certain extent. With

931

VI I

%

s

to GC ana~ysNis 2 exhaustg

KF

A

Figure 1. Schematic drawing of the experimental setup

the valves V1 to V3 the educt or product gas mixture can be analysed by a capillary gas chromatograph and in some experiments, after passing two cooling traps, CO and CO2 can be analysed by a non dispersive IR photometer. A bubble flow meter is used for measuring the total flow rate. The whole apparatus with exception of the reactor is made of stainless steel and heated to 373K resp. 543K to avoid the condensation of xylenes and products. We quantify the results by the use of the following parameters obtained by mass balancing the open system plug flow reactor: -

the modified residence time tv defined by the mass of the catalyst mcat divided by the flow rate V at ambient temperature

tv-

mcat

v

- the xylene conversion in the reactor

(3)

932

9

Xi -

~

nxylene,O -- nxylene 1;lxylene,O

(4)

- the integral reactor selectivity RS i for the formation of product i Iii " s

(5)

Rsi ----- (ftxylene,O- ftxylene) 9~xylene

- the educt selctivity for para-conversion ESp defined as conversion of p - xylene ESp = (conversion of p xylene) + (conversion of m - xylene) -

Xp-x Xp-x + Xm-x

(6)

the dimensionless normalized concentration yi corresponding to the fraction of total incoming carbon contained in the form of species i at reactor outlet

Yi = .

fli 9~i

nxylene,O " s

(7)

where fii designates the molar flow rate of species i and ~i the number of carbon atoms in that species (~xylene = 8). In the case of a product, Yi corresponds to the yield Yi. 3. R E S U L T S 3.1.

Screening

tests

Prior to the kinetic study we performed screening experiments at 673K to 723K with equimolar mixtures of meta- and para-xylene. This first set of experiments gave the following tendencies: - the educt selctivity for para-conversion ESp is almost 90% independent of temperature and conversion range. -

ESp raises slightly with decreasing particle size L. Reactor selectivity for the formation of para-products raises with decreasing particle size.

- The catalysts prepared from Na-ZSM-5 are much more stable to coking than the catalyst prepared from the hydrogen form (H-ZSM-5). - The catalyst modified with copper gave the highest values of activity and selectivity. Therefore the following study was restricted to a series of copper exchanged catalysts prepared from Na-ZSM-5 (Si/A1 = 37) with L = 0.08 microns and copper contents from 0.9 to 1.8 weight% (equal 60% to 130% apparent exchange level) by varying the time for ion exchange or repeating to the exchange procedure.

933 3.2.

Copper

modified

zeolites

To simplify the experimental procedure only p-xylene was used for the following experiments with a mole fraction of Xp-x,Feed 0.015. With increasing copper content the activity increases: for a residence time of tv = 1 16 m~,mi, and a temperature of 673K the xylene conversion increases from 25 to 43% (Fig. 2). =

9

0 . 8

.

0.7

. . . . .

X~_x,~..

0.6-

.

IVT = t

.

.

R$

.

673 K

=

0.5 . . . . . . . . . . . . . . . . . . . . . . . . . . . .

~)

.

.

.

.

.

.

.

,

eei

',

0.8 .................

'.. . . . . .

0.75 -

,. . . . . .

,

. II

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

~

,

o

0.6 0.55

I

0.8

1.2

1.4

1.6

0

o

~

~

~

,

~

.....

0

0

0

0.65 ..........................................

~

0.1-

ir,

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

:. . . . . . I

0.2

0

-It ...... '

0.7 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

~

i ........ =i ...... !. . . . . .

0

0.3

,

iJT = 673 K Xp-x = 0.45

0.85 ....................

' ......

0.4 . . . . . . . . . . . . . . . r ....... .e:

1 0.95 . . . . . . . . . . . . . . . . . . . . . . . o.s . . . . . . . . . . . . . . . . . . . . .

. . . . . .

([mg*mln]/cm

t ~

CIII--3

1.8

2

Cu-confenff/weight-% Figure 2. Conversion of p-xylene versus copper content

. . . . . . .

. . . . . .

o.s 0.8

:.

. . . . .

,

'

.

i

i

1

1.2

o

:.

i . . . . . .

o

o

. . . . .

" . . . . . .

~

.

~

i t4

o

.

.

:.

. . . . .

,

o

~

i

i

t6

1.8

- ....

2

Cu-confenf/weight-%

Figure 3. Reactor selectivity for pproduct formation versus copper content

The selectivity for para-products has a slight maximum for a copper content of approximately 1.6 weight% equivalent to an apparent exchange level of 100% (Fig. 3). For a given temperature and a given catalyst, conversion of p-xylene and yield for the para-products increases linear with oxygen content in the educt gas (xo2 = 0.04 - 0.2). By variation of the ammonia content in the educt gas we observed a maximum activity at XSns,Feed ~ 0.20 (Fig. 4). For higher ammonia contents the activity decreases and the selectivity for para-product formation remains constant (Figs. 4,5). As a conclusion of this second set of experiments we used a feed composition of X~yl~.~ = 0.015, xo2 = 0.13, XNH3 0.2 and a catalyst with 1.6 weight% copper. The reaction of p-Xylene is a first order process. The temperature dependence of xylene reaction rate in the range 523K to 673K can be described by an Arrhenius term with an apparent activation energy of 110 ka The observed selectivities for the para-produts are independent of reactor temperature. For temperatures higher than 683K it was not possible to establish an approximately isothermal reactor. Large increases in CO2 production and temperature along the reactor axis were observed. --

934

0.8

.

Xp-xylane

.

.

.

:

,

~'

9. . . . .

&

; .......

.'- . . . . . .

.:o . . . . . . .

:~ . . . . . .

0.4-

,

T

673 . K

=.

tv

=

,3.73

., . . . . . . . . . . . . . . . .

.

.

O.5 0.05

([mg*min]/cm

.

0.4

" ......

.

,. . . . . . .

3)

9. . . . . . .

I

I

0.1

0.15

0.2

0.25

,

'

--,--

'

*

-

0.1-

9 Yp

,

", . . . . . .

.j

.. . . . . .

,

,

0.05

,

,

Xp-x = .

i

0.1

; .....

,~ . . . . .

T = 673 .

i

*i*

, ......

~, . . . . . .

.......

i

0

"O'"

i ......

~Xs o

0.3

.~t.-

, . . . . :- . . . . .

o.2--

o

,

,

'

* i * +. *

:~

',. . . . . .

. I

.

0.3 .......

,. . . . . . .

I

,

,

o.5-

~

-' . . . . . . .

,

,

~t--~

0.6. . . . . . .

"'I~""

; ...............

" ..............

.

Sp

.&. ',.. . . . . . . . . . . . . . .

,

0.7-

'~

&"

,'

0.8

......

,

:

9

.

0.5 .........

i

9. . . . . . .". . . . . . . . . .

A

'

......

:,. . . . . . .

......

,

0,7 .......................

0.6 ....

Yp

.

:

""

i

0.15

,

I

0.2

0.25

0.3 •

XNHs

Figure 4. Conversion of p-xylene versus a m m o n i a content

3.3.

K

0.45

Figure 5. Reactor selectivity/yield for pproduct formation versus a m m o n i a content

Kinetic study Typical results for the ammoxidation of pure p-xylene are shown in Figs. 6 and 7.

1-

9

RS I

Yl

0.6-

~

~

;, - -

9

,~

'

:

a S~.

A:

,

t5

9

~

3

~

,

,.5

5

'lv/Cmg "min "cm -3)

Figure 6. Yield as function of modified residence time

?

..~. ..~..

0.2-

o.5

-s~.

:

S n ~

,!

0.4-

o

~

.

0-8-

o.

9

:

o

,

0.1

I

0.2

^

.

!

:

j a

%

.--. . . ...,: -. u

0.3

I

0.4

^-^

A

.:-:.-. . . -.. . i

0.5

I

0.6

u

0.7

i

9

....... -:.:. ..... I

0.8

i

0.9

1

X~x

Figure 7. Reactor-selectivity versus pxylene conversion

The yields of the mononitrile and the dinitrile are increasing with increasing residence time but the yield of byproducts (CO, CO2, toluene and others) is increasing too. It is possible to describe the kinetic behaviour of the ammoxidation of p-xylene by a simplified scheme (Fig. 8).

935

r13 ,r 12

p-Xylene

r 23

=- p-Mononi|rile

~- p-Dinifrile

1

Byproducts Figure 8. Reaction network

For the numerical description simple first order kinetics without reaction r34 is sufficient. The numeric results are represented by the lines in Figs. 6 and 7. This simple first order description is also valid for the ammoxidation of an equimolar mixture of m-xylene and p-xylene (Figs. 9,10). ~,' - .......... @.6-----

0,4

......

...........

:: i ii i .......... ::.......... i.. I v

:

:

:' _ 9

9 9

" .......... . aA

4~

0

0.5

1

1~

2

2.5

3

3.5

4

4.5

5

t,,/(,~-==.~-~) Figure 9. Yield versus modified residence time for conversion of an isomer mixture

' 0

w 0.1

I 0.2

__:v .

i' I

....

9

" ......... .. .

i~ i 0.3

! 0.,4

*s= ,s. .s.p

~"7"......

~, . . . . . . . . . . 9 ,

.

- " ~ .........

',. . . . . . . . . . , ~ .^

" a.a-= I 0.5

I 0.8

t 0.7

--I 0.8

I 0.9 X~-x

Figure 10. Reactor selectivity versus pxylene conversion for conversion of an isomer mixture

m-Xylene does not affect the kinetics of p-xylene conversion. The values of the kinetic parameters are only slightly different compared to the reaction of pure p-xylene as shown in table 2. As a consequence, almost 90% educt selectivity can be observed up to high values of p-xylene conversion for the reaction of the isomer mixture (Fig. 11). 4. C O N C L U S I O N S The results confirm the expectations for educt shape selectivity as reversion of product shape selectivity for the reaction of a mixture of xylene isomers over ZSM-5 catalysts.

936 Table 2 Kinetic parameters (cm 3 9(mg 9min) -x) for simplified reaction network reactants k12 kls k14 k23 p-xylene 0.381 0.049 0.193 0.070 m,p-xylene 0.420 0.042 0.224 0.084

k24 0.017 < 0.01

'

ESp

0.9

1

0.8

:

:

:

:

', ' A

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

:

0.7-

:

:

:

:

'~

~,

:

:

:

: '

i ,

i :

:

i

,. . . .

: . . . .

:

:

i

:

0.6

e;

i ,

....

i

i

0.4

, i

i

!

i

,

o Xrn-X

Xi -

0.8

-

0.6

-

0.4

-

0.2

'

}---

o.s-

-1

I

0.3 0.2 0.1 0

-0 0

0.5

1

1.5

2

2.5

3

3.5

4

4.5

5

fv/ [mg*min*cm-3] Figure 11. Educt-shape-selectivity for conversion of an equimolar mixture of m-/p-xylenes

Due to the chemical sink inside the zeolite crystals the reaction of para xylene to para products is favored up to high xylene conversions. From a more practical point of view it would be necessary to follow the reaction to higher mean residence times. Only for a higher yield of dinitrile with increasing residence time a practical application of this way to terephthalic acid is possible. The main advantage of the proposed route compared to the known processes would be the unification of isomer separation and reaction in one step. 5. Acknowledgements A part of this work has been supported by the Deutsche Forschungsgemeinschaft SFB 250 "Selektive Reaktionsfiihrung an festen Katalysatoren'. Fruitful discussions with L. Riekert are gratefully acknowledged by the authors.

937 REFERENCES

1. P.B. Weisz and V.J. Frilette, J. Phys. Chem. 64 (1960) 382 2. P.B. Weisz, V.J. Frilette, R.W. Maatmann and E.B. Mower, J. Catal. 1 (1962) 307 3. P.Chutoransky Jr. and F.G. Dweyer, in "Molecular Sieves"' (W.M. Meier and J.B. Uyterhoeven, Eds.), p. 540. Advances in Chemistry Series, Vol. 121, Washington DC, 1973 4. J. Wei, J. Catal. 76 (1982) 433 5. K. Beschmann and L. Riekert, J. Catal. 141 (1993) 548 6. K. Weissermel and H.-J. Arpe, Industrielle Organische Chemie, 4th ed., Verlag Chemie, Weinheim 1994, p. 430

This Page Intentionally Left Blank

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

VANADIUM-TITANIUM E.M.Al'kaeva, E.B.Burgina

939

O X I D E S Y S T E M IN [3-PICOLINE O X I D A T I O N

T.V.Andrushkevich,

G.A.Zenkovets,

G.N.Kryukova,

S.V.Tsybulya,

FRC Boreskov Institute of Catalysis, Novosibirsk, Russia

Binary vanadium-titanium oxide catalysts with various ratios of vanadium oxide and titania, as well as individual oxides of vanadium and titanium were examined in oxidation of 13-picoline. Nicotinic acid, 3-pyridinecarbaldehyde, and CO2 were the reaction products over all the catalysts. The binary catalysts and individual vanadium oxide were highly selective for nicotinic acid, the most effective in 13-picoline oxidation were the samples containing 20% and more of vanadium pentoxide. A regular stacking of crystallites of V205 and TiOz was found to be the characteristic feature of the structure of the most effective compositions.

1. INTRODUCTION The literature on vanadium-titanium oxide catalysts for selective oxidation of organic compounds is quite extensive. Some of the publications deal with the vanadium-titania system as a whole and reveal the most effective compositions [ 1-3]. Different structural features are characteristic of these compositions depending on the modification of titania (rutile or anatase). The authors of [1, 2] attribute the observed higher efficiency of the binary rutile-containing catalysts against the individual vanadium oxide to the formation of a solid solution in the system and to the ability of such a solution to stabilize a certain valence state of vanadium. As to the binary anatase-containing catalysts, their high efficiency is accounted for by dispersing vanadia through the anatase surface [3] and by the structural fit between the crystal lattices of anatase and vanadium pentoxide [4]. Formation of vanadyl bonds with the optimal bond energy of oxygen is also mentioned [5]. We showed the vanadium-titania catalysts to be also effective for oxidation of 13-picoline to nicotinic acid [6]. In the present work we studied the influence of the chemical composition of the vanadium-titania system on the phase composition, the structural arrangment and catalytic behavior in [3-picoline oxidation. 2. EXPERIMENTAL Binary samples containing 5 to 75 wt % of V205 and 95 to 25 wt % of TiO2 were prepared by mixing titania of the anatase modification, produced according to the sulfate technology [7], with vanadyl oxalate solution, drying the mixture and calcining it at 450~ Individual oxides were prepared from the corresponding raw materials in the same manner.

940 Thermal argon adsorption was ffsed to determine the specific surface area of the samples. XRD analysis of catalysts was carried out using a URD-6 diffractometer with monochromatic CuK~ radiation. IR spectra were recorded at 200 to 1300 cm"1 using a BOMEM Fourier-spectrometer. Weighed 2 mg catalyst to be studied were pelleted by pressing with CsJ. TEM studies were carried out using a JEM-4000FX microscope with resolution of 1.1 ~, and accelerating potential of 400 kV. Catalytic activity of the 0.25-0.50 mm fraction of catalysts was measured using a flowcircuit setup with a differential reactor at 250-300~ The inlet reaction mixture was 1 vol % of 13-picoline, 20 vol % of oxygen, 30 vol % of water steam, and the balance nitrogen; contact time was 0.15 to 3.00 s. The reaction products were analyzed chromatographically.

3. RESULTS 3.1. Catalyst characterization Chemical composition, specific surface area, and phase composition of the examined catalysts are summarized in the Table. The specific surface area of the catalysts is seen to decrease as the concentration of vanadium pentoxide increases; the reason is sintering of the support in the presence of V2Os.

Table Characterization of catalysts Sample V205, wt % 1 2 5 3 10 4 15 5 20 6 30 7 50 8 75 9 100

Specific surface area, m2/g 300 290 278 224 40 39 27 10 7

Phase composition TiO2* TiO2* TiO2* TiO2* V2Os+TiO2* V2Os+TiO2* V2Os+TiO2* V2Os+TiO2* V205

* anatase 3.2. XRD analysis The obtained XRD data (see the Table) showed the anatase phase alone in the samples containing less than 20% of vanadium oxide. The phases of anatase and V205 are detected in the catalysts at higher concentrations of V205.

941 3.3. IR spectroscopic studies The spectra recorded for individual TiO2 and the binary catalysts containing 5, 20, and 50% of V205 are presented in Fig. 1.

IR spectrum of individual titania and the IR spectra of the binary catalysts coincide well with the spectrum of anatase [8] at 200-1300 cm 1. However in the spectra recorded for individual TiO2 and the sample containing 5% of vanadium oxide, the maxima of broad absorption bands (a.b.) at 540 and 342 cm 1, which are characteristic of r anatase, are shifted downfield to 517 and 332 cm1, at the same time a.b. at 965 and 1050cm ~ with 1070 and 1125cm ~ shoulders, respectively, are observed, that indicates the presence of sulfate-ion admixtures. There are neither shifts of a.b. at ee~ 540 and 342cm ~ nor a.b. generated by r r admixture ions in the spectra recorded for the samples containing 20% and more of V205, these observations evidence for ordering of the anatase structure. An a.b. at 900 cm 1 produced by oscillations of the V-O-V fragment [9] appears in the spectra recorded for the binary 1200 1000 8()0 6()0 400 1 samples containing no less than 5% of V205, V, cmand an a.b. at 1020 cm1 attributed to Fig. 1. IR spectra of individual anatase (1) formation of the V=O fragment for the and binary catalysts containing 5% of V205 samples with more than 20% of VzOs. Extra (2), 20% of V2Os (3), and 50% of V2Os (4). bands at 809, 376, and 300 cm -1, which are characteristic of the V205 phase are seen in the spectra recorded for the samples with more than 50% of VzO5 [8]. Therefore, the XRD and IRS data show that the binary catalysts with the compositions ranging between 20 and 75% of V205 contain the phases of V205 and TiO2 (anatase). No other compounds were detected by these techniques. 3.4. TEM studies Fig.2 (a, b) presents micrographs of samples containing 5% (a) and 20% (b) of vanadium oxide. Identical images were obtained for the samples before and atter the reaction. Highly

dispersed crystallites of TiO2 (30-80 /~ in size) and randomly arranged associates of vanadium oxide are observed at low concentration of V205 (Fig.2a).

942

~., ~. ~

~,~::-'.:!~:~'!:i:.'.x.. "

~!;iii~i!iiii! 30A

..

Fig. 2. Micrographs of binary catalysts containing 5% ofV205 (a) and 20% ofV205 (b). 60

m 1

40 ~.-

30

~

2o

10

~

~176176176

o

100 4

80 60

o

Z

r.~ 40

~x

20

~, 5.

o~ 10 ~

5

o~~

6

i

!

20

40

!

60

80 X,%

Fig. 3. Rate of oxidation of 13-picoline (1+3) and selectivities for nicotinic acid (4), 3-pyridineversus conversion of 13-picoline. carbaldehyde (5) and CO2 (6) O - 250~ A - 270~ r--l_ 300oC.

943 An increase in the concentration of V205 results in a higher quantity of dispersed associates of the supported phase. At the V205 concentration as high as 20% (Fig.2b) the interfacial stacking between the crystallites of anatase and V205, implying the regular mutual arrangement of their crystal lattices, is observed. Such a stacking was described in [ 10]. The phase stacking occurs along the (100) axis, which is the common one for the unit cells of V205 and TiO2. The angle between (001) face of vanadium pentoxide and (110) face of the unit cell of titania is 17.4 ~. The further growth of the concentration of vanadium oxide induces no change in the nature of the interaction but increases the stacking surface area.

3.5. Catalytic properties Oxidation of 13-picoline over all the catalysts produced nicotinic acid, 3-pyridinecarbaldehyde, and CO2. The specific activity and selectivity for the reaction products against conversion of 13-picoline are shown in Fig.3 for the sample containing 20% of V205. The main reaction product is nicotinic acid; the selectivity for this product is 90% within the conversion range of 50 through 90%. The formation of nicotinic acid is intermediated by 3-pyridinealdehyde. The selectivity for the intermediate decreases abruptly when the conversion reaches 30%, the selectivity for nicotinic acid increases correspondingly at this conversion range. The byproduct, CO2, is mainly produced from 13-picoline by a parallel reaction pathway. There was observed only a little overoxidation of partial products, the aldehyde and nicotinic acid, by a consecutive reaction pathway. The high selectivity for nicotinic acid was observed over the whole temperature range. According to the type of the dependence of selectivity for the reaction products on the conversion of 13-picoline, the reaction pathway over the catalysts under study can be presented schematically as follows:

c//

O

c'~

O

C02 Variation in the catalyst composition does not change the type of the dependencies shown in Fig.3 but influences the activity and contribution of the parallel stages.

Fig.4 shows the selectivities for the reaction products and the rate of oxidation of [3-picoline against composition of the vanadium-titania system. All the binary vanadium-titania catalysts and the individual vanadium oxide reveal high selectivities for nicotinic acid. The selectivity for nicotinic acid increases from 75 to 90% and the rate of [3-picoline oxidation increases by three times as the concentration of V205 increases from 5 to 20%, no further

944

~1 ~

20

1, o 10

.20

/o--O------~----__o,.~

15

/o

10

V

5

,L

90

.9o - - o

8o

o-~176

70.

ck O - . o _ o

60-

~176

~

o

I

o-~2;~2~

o

15

o

I

o I 50

v

o-~ 3

10

o-14

5 100

----~V205,%wt Fig.4. Rate of conversion of 13-picoline (1) and selectivities for nicotinic acid (2), 3-pyridinecarbaldehyde (3) and CO2 (4) v e r s u s composition of the binary vanadium-titania system. Reaction temperature is 250~ Conversion of J3-picoline is 45%.

change being observed with the binary system. Individual vanadium oxide reveals lower activity and selectivity compared to the binary catalysts. The rate of [3-picoline oxidation and selectivity obtained with individual titania are not shown in the Figure; it exhibits a low activity and much lower selectivity against the other catalysts. The rate of ]3-picoline oxidation over this catalyst is lower by two orders of magnitude, and the selectivities for nicotinic acid, aldehyde, and CO2 are 26.2%, 49.8%, and 23.9%, respectively, even at the conversion as low as 3%. Thus, the highest activity and selectivity were observed for the binary catalysts containing 20% and more of V2Os.

4. DISCUSSION

All the binary vanadium-titania catalysts and vanadium oxide are selective for oxidation of 13-picoline to nicotinic acid. The only by-product is carbon dioxide. The selectivity for CO2 is no more than 10% even at high conversions observed for all of the binary catalysts and individual vanadium pentoxide. However we can distinguish the most effective compositions

945 among those under study; these are the binary catalysts containing 20 wt % and more of vanadium pentoxide. From the physicochemical data, the catalysts differ in XRD phase composition, microstructure, molecular structure depending on the concentration of vanadium oxide. At below 20% of vanadium oxide, the samples are the anatase phase with highly disperse XRD amorphous associates of vanadium oxide in the form of V-O-V fragments allocated over the anatase surface. The catalysts containing more than 20% of vanadium oxide is built up by the crystallites of anatase and V205. In this system vanadium is in the form of V=O and V-O-V. However one cannot attribute the higher efficiency of the binary catalysts to the emergence of the V205 phase since the individual vanadium pentoxide reveals lower activity and selectivity. A distinctive feature of these systems is the presence of coherent stacking between TiO2 (anatase) and vanadium pentoxide crystallites. The arrangement of the cells of TiO2 and V205 in the region of the stacking is presented schematically in Fig.5. The stacking makes it possible to form the V - O - Ti bonds. The states of the cations and oxygen in such a structure are expected to differ considerably from those in the individual oxides. This suggestion is supported by the data reported in [11], demonstrating a lower bond energy and a higher mobility of oxygen, also enhanced redox properties of the catalyst containing 20% of vanadium pentoxide and 80% of anatase compared to the individual oxides.

=-a

7,

0

0

b

0

0

,

o-0 O-V o-Ti

a

Fig. 5. Schematic presentation of mutual arrangement of TiO2 and V205 unit cells in the region of their stacking. We can conclude that the formation of the interface between the crystallites of anatase and vanadium pentoxide, which is likely to produce the states of vanadium and oxygen other than those in the individual oxides, can be responsible for a higher efficiency of the binary catalysts.

REFERENCES

1. A. Andersson and S.L.T. Andersson, Solid State Chemistry in Catalysis, ASC Symp. Series, 279 (1983) 121. 2. D.Kh. Sembaev, B.V. Suvorov, L.I. Saurambaeva, Kh.T. Suleimanov, Kinet. and Katal., 20 (1979) 750. 3. G.Ya. Popova, T.V. Andrushkevich, G.A. Zenkovets, Kinet. and Katal., to be published.

946 4. 5. 6. 7. 8.

M. Gasior, T. Machej, J.Catal., 83 (1983) 472. K. Mory, A. Miyamoto, Yu. Marakami, J. Phys. Chem., 88 (1984) 274. WO 95/20577, 03.08.95. Yu.D. Dolmatov, A.I. Sheinkman, Zh. Priklad. Khim, 43 (1970) 249. E.N. Yurchenko, G.N. Kustova, S.S. Basanov, Vibranal spectrum of inorganic compounds, Novosibirsk, 1981, 145. 9. B. Buska, G. Ricckard, D.S. Hewsom, J.C. Valta, J. Chem. Soc. Far. Trans., 90 (1994) 1161. 10.G.N. Kryukova, D.O. Klenov, G.A. Zenkovets, React. Kinet. Catal. Lett., 60 (1997) 179. 11.V.M. Bondareva, T.V. Andrushkevich, Yu.D. Pankratiev, React. Kinet. Catal. Lett., 61 (1997)

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

947

Selective alkene epoxidation by molecular oxygen in the presence of aldehyde and different type catalysts containing cobalt O.A. Kholdeeva a, I.V. Khavrutskii a, V.N. Romannikov a, A.V. Ykachevb and[ K.I. Zamaraev I a

Boreskov Institute of Catalysis, 5 Pr. Lavrentieva, Novosibirsk 630090, Russia

b Novosibirsk Institute of Organic Chemistry, 9 Pr. Lavrentieva, Novosibirsk 630090, Russia

Catalytic properties of different type compounds containing cobalt, namely, tetra-nalkylammonium salts of PWllCOO~9 (PW,~Co) and CoW~2060 (CoW~2) heteropolyanions (HPA), Co(NO3)2.6H20, CoNaY zeolite and Co(II) phtalocyanine (CoPc), in alkene epoxidation by dioxygen in the presence of iso-butyraldehyde (IBA) have been studied at ambient conditions. Yields of the epoxidation vary from 80 up to 99% depending on the olefin and catalyst. The reaction proceeds via radical chain mechanism with degenerate chain branching, acylperoxy radicals being the main epoxidizing species. Catalytic activity of different Co(II) compounds in alkene-IBA co-oxidation correlates ,with their activity in decomposition of perisobutyric acid (PIBA). The superior activity of Co(II) catalysts in alkene-IBA co-oxidation arises most probably from their ability to catalyze the chain branching via the homolytic decomposition of PIBA formed in the course of IBA autoxidation, and promote the chain initiation via the interaction with IBA. 1. INTRODUCTION Selective catalytic oxidation of hydrocarbons with molecular oxygen is a field of both academic and industrial interest [1-4]. Reactions of organic substrates with dioxygen in the ground (triplet) state are restricted by spin conservation, which may be overcome by two main routes: 1) via the formation of complexes with transition metals (TM), and 2) via involving 02 into a chain radical process. The former route was suggested for natural redox enzymes, such as cytochrome P-450, which are known to provide high selectivity of oxidation at ambient conditions [1-4]. At the same time, chain radical autoxidation processes are known to require quite rigid conditions and proceed with poor selectivity because of over-oxidations and side reactions [1, 5]. During the last two decades much attention was focused on the search for catalytic systems that use dioxygen and mimic the action of cytochrome P-450 monooxygenases. These systems are usually based on the combined use of 02, a reducing agent, such as NaBH4, Hz/Pt, Zn, ascorbate or aldehyde, and TM complexes [see for review refs. 1-4]. In early 1990s Mukaiyama and co-workers have found that the use of branched aliphatic aldehydes as the reducing agents provide alkene epoxidation with impressively high selectivity under mild reaction conditions [6-8]. Co-oxidation of alkenes with aldehydes

948 (mostly with acetaldehyde and benzaldehyde) has been intensively studied in 1960-70s and has been found to proceed at elevated temperature and/or oxygen pressure via radical chain mechanism producing moderate yields of the epoxidation [1, 5, 9, 10 and references cited therein]. Unusually high yields of epoxidation obtained at ambient conditions in Mukaiyama's system attracted attention of different research teams. Various TM compounds, including complexes with organic ligands [6-8, 11-14], TM-substituted heteropolytungstates [15-19], simple salts of Fe(III), Ru(III) and Cu(II) [20, 21], Ni(II)-modified clay [22, 23], Ti(IV)silicates [24] and TM-substituted NaY and NaZSM-5 zeolites [25], appeared to catalyze this reaction, cobalt-containing compounds being among the most active catalysts [ 15, 16, 18, 19, 25]. Earlier it was shown that cobalt porphyrins display higher catalytic activity as compared to other TM porphyrins in propylene epoxidation with O2 in the presence of propionaldehyde [26]. In spite of considerable number of investigations, the mechanism of dioxygen activation in O2/aldehyde/TM systems as well as the reasons for the remarkable catalytic activity of cobalt compounds still remain an issue. Here we report a comparative study of catalytic properties of different type cobaltcontaining compounds in alkene epoxidation by dioxygen in the presence of iso-butyraldehyde (IBA) and provide some data, which allow us to clarify the reaction mechanism and the nature of the catalytic action of cobalt compounds. 2. RESULTS AND DISCUSSION

We examined catalytic properties of cobalt-containing compounds having different nature, 5namely, the simple salt, Co(NO3)2.6H20, tetra-n-butylammonium salts of PW11CoO39 (PWllCO) and CoW12046o(CoW12) heteropolyanions (HPA), CoNaY zeolite and Co(II) phtalocyanine (CoPc), in alkene epoxidation by the O2/IBA system. We have found that various alkenes can be converted to the corresponding epoxides with good-to-high selectivity (80-99%) at complete alkene conversion at ambient conditions (Table 1). Neither allylic oxidation nor epoxide ring cleavage products were detected for all alkenes tested, except for cyclohexene. The nature of catalyst does not considerably affect the selectivity of the epoxidation, which depends mainly on the olefin structure. Some decrease of the selectivity was generally observed at high cobalt concentrations (about 1.10.2 M). It is widely accepted that high selectivity in oxidation processes is incompatible with a chain radical mechanism of the reaction. The yields of epoxidation are known to be much higher for alkene-aldehyde co-oxidation than for alkene autoxidation [1, 5, 9, 10]. Both of these types of alkene oxidation processes were found to proceed via chain radical mechanism, however, acylperoxy radicals formed in the presence of an aldehyde, in contrast to alkylperoxy radicals, are more prone to add to alkene double bond than to abstract an allylic hydrogen atom [10, 27, 28]. This accounts for higher selectivity of alkene-aldehyde co-oxidation. The yields of epoxides observed by different authors in alkene oxidation by Mukaiyama's type systems using branched aldehydes and TM catalysts, are superior to those obtained earlier for alkene-aldehyde co-oxidations [6-8, 11-25]. This led many authors to suppose the formation of various non-radical active species, responsible for the selective alkene epoxidation in Mukaiyama's type catalytic systems. Acylperoxy radicals [11, 12, 14] or peroxy acids [26] coordinated to TM center, high-valent metal-oxo species [15, 20, 21, 24, 26], hydroperoxo and peroxocomplexes [14], singlet oxygen-like species [14, 30], and some others were

949 considered as the possible donors of active oxygen. The activation of dioxygen and/or aldehyde via coordination on the metal center was also suggested [12, 14, 23, 26]. Table 1 Alkene epoxidation by 02 in the presence of IBA and Co(II) catalysts (a) Alkene

Catalyst

Time (h)

Alkene conversion (%)

Yield of epoxide (b) (%)

(+)-3-carene

CoNaY Co(NO3)2

2.5 3.0

100

100

99 99

ct-p inene

Co(N 03 )2

5.5

99

96

(-)-caryophyllene

CoNaY Co(NO3)2

2.0 2.5

100

100

99 (c) 98 (c)

cyclohexene

Co(NO3)2

4.0

97

85 (d)

trans-stilbene ~e)

PWl 1Co

2.5 4.0 10.0 3.0 3.0

100 100 100 100 100

90 (f~ 82 81 83 84

COW12 CoPc Co(NO3)2 CoNaY

(a) Reaction conditions: alkene 0.30 mmol, IBA 2.28 mmol, catalyst ,6.10 .3 mmol (CoNaY 11 mg), MeCN 3 ml, air, 24~ (b) GLC yield based on alkene consumed; (c) only monoepoxide is formed; (d) 2-cyclohexenone is the main co-product; (e) IBA 1.14 mmol; (f) only trans-epoxide is formed together with benzaldehyde. On the other hand, the reaction under discussion was found to be inhibited by small additives of radical scavengers, such as ionol and benzo- or hydroquinone (HQ) [14, 17, 19, 23, 25]. This fact indicates unequivocally the radical chain mechanism of the reaction. We observed the inhibition of the epoxidation by HQ and ionol for all the cobalt catalysts studied, except for CoPc. Fig. 1 demonstrates the inhibition effect of HQ for trans-stilbene-IBA cooxidation in the presence of Co(NO3)2. We recently observed a similar picture for PWllCO and CoNaY catalysts [ 19, 25]. The lack of inhibition in the case of CoPc, however, does not eliminate the radical chain component in the reaction with this catalyst, because scavengers may have no effect in reactions involving short radical chains [2]. The extreme sensitivity of the alkene-IBA co-oxidation to microimpurities of catalysts and inhibitors, which was discussed previously [19, 25], also indicates the chain reaction mechanism. Fig. 2 shows the dependence of the epoxidation rate estimated from the fast-rate portions of kinetic curves, which usually have S-type character, on the amount of the Co(NO3)2 catalyst. Similar dependencies of the oxidation rate on the Co(II) concentration were observed previously [19, 29]. Such a dependence is consistent with the radical chain mechanism and may be attributed to the participation of the catalyst in chain termination [1, 5, 29] via the reaction

C o I + RCO~ -..

~ Co II

RCO3

950 Figure 1. Kinetic curves of alkene consumption (11) and epoxide accumulation (~) for aerobic oxidation of trans-stilbene (0.3 mmol) in the presence of IBA (1.14 mmol), Co(NO3)2.6H20 (6.10 .3 mmol) and hydroquinone (3.10 .3 mmol) in acetonitrile (3 ml), at 25 ~ Hydroquinone was added 45 min. after the beginning of the reaction

100 o 80 E 60

/

o 40 -~ 20

~ -

0

i

,

-

,

m

3 4 Reaction time, h

5

Note that even "critical phenomena" for the catalyst-inhibitor conversion were observed in alkene epoxidation by O2/IBA system in the presence of PWllTi and PWllV catalysts [19]. All these data provide evidence in favor of the radical chain mechanism of the epoxidation in most O2/IBA/Co(II) systems studied where acylperoxyradicals are most probably the main epoxidizing species 9 "7 r~

=._ 20 Figure 2. Dependence of the rate of trans-stilbene epoxidation on the concentration of Co(NO3)2 (trans-stilbene 0.3 mmol, IBA 1.14 mmol, acetonitrile 3 ml, air, 25 ~

O

-.....

_ •

10

o

e ..~ A

~

5

A

v

........

,

........

,

........

,

,

-2 Concentration of Co(NO3)2, tool.1-1

10 -s

10 -4

1 0 -3

10

Non-stereospecific character of the epoxidation in the Mukaiyama's type catalytic systems was shown previously [6, 17-21, 25]. At the same time, in the absence of metal catalyst the epoxidation of cis-alkenes was found to proceed with the retention of alkene configuration [31 ]. We have studied the stereochemistry of the epoxidation of cis-stilbene by O2/IBA/Co(II) systems. The distribution of cis- and trans-epoxides for the reactions with different Co(II) catalysts is given in Table 2. The trans-epoxide is presumably formed in all Cases. The nonstereospecific character of the epoxidation excludes any known concerted mechanism for oxygen transfer to alkene molecule and therefore, eliminates peroxy acid (coordinated or not) as the main epoxidizing species. The loss of alkene stereochemistry supposes that the oxygen atom transfer step proceeds via free-rotating intermediate, which may be formed by the addition of a radical species, e.g., acylperoxy radical, or metal-oxo species to alkene double bond [ 1-4, 10, 18]. Note that the loss of stereochemistry is more for CoPe as compared to the other homogeneous Co(II) catalysts studied (Table 2). We believe that species, different from acylperoxyradicals, might be involved in the epoxidation process in the presence of CoPc catalyst. Remember that the reaction in the presence of CoPe is not inhibited by radical scavengers (Fig. 1).

951 Table 2 Epoxidation of cis-stilbene by 02 in the presence of IBA and homogeneous Co(II) catalysts (a) Catalyst

PW1 ICo

COW12

Co(N03)2

CoPc

Trans-/cis-epoxide ratio (b)

3.8

3.4

2.3

13

(a) Reaction conditions: cis-stilbene 0.30 mmol, IBA 1.14 mmol, catalyst 6.10 .3 mmol, MeCN 3 ml, air, 24~ (b) determined by 1H NMR by comparison of peaks with 8 4.15 (cis-epoxide) and 8 3.66 (trans-epoxide). Finally, we address the question of the reasons for the superiority of Co(II) compounds as the catalysts for the alkene epoxidation by O2/aldehyde system. We have studied the catalytic activity of Co(II) compounds using trans-stilbene as a model substrate. Kinetic curves for the trans-stilbene epoxide accumulation are given in Fig. 3. The kinetic curves show autocatalytic character. The time of the complete alkene conversion depends on both the induction time and the rate of the reaction after the completion of the induction period. It should be noted that the induction time increases considerably with decreasing aldehyde concentration and goes through a maximum with increasing Co(II) concentration (in the range 10-5+10.2 M). One can see from Fig. 3 that the rate of the epoxidation lowers in the order: PW~ICO, CoW~2 > CoNaY, Co(N03)2 > CoPc.

100

Figure 3. Kinetic curves of trans-stilbene epoxide accumulation for aerobic oxidation of trans-stilbene (0.3 mmol) in the presence of IBA (1.14 mmol) and Co(II) catalysts (6.10 .3 mmol): COW12 (A), PWllCO (O), Co(NO3)2 (11), CoPc ((3) and CoNaY (~) (11 mg)

80 60 O

40

20 0 0 1 2 3 4 5 6 7 8 9 10 11 Reaction time, h

It was assumed that high activity of Co(II) compounds in alkene-aldehyde co-oxidation is due to the ability of cobalt to form complexes with dioxygen, thus providing its activation [ 14, 26]. Indeed, numerous Co(II) chelates are known to form stable superoxo and pperoxocomplexes by reaction with dioxygen [ 1-4]. Note that among the catalysts studied, only CoPc is known to coordinate 02 [1, 3 and references therein]. Nevertheless, it showed less catalytic activity as compared to the other Co(II) catalysts employed (Fig. 3). By analogy with manganese and iron complexes, the formation of active CoIV--o or CoV=O species by the twoelectron oxidation of Co(II) or Co(III) ions with peroxy acid was suggested by few authors [15, 26]. It is well known that appropriate ligands are needed to stabilize higher oxidation states of transition metals [1-4]. Alkene epoxidation by PhIO in the presence of PWllCO and PWllMn was proposed to proceed most likely via the formation of the high valent metal-oxo species [2, 18, 32]. The formation of such species may in principle be expected for CoPc, but

952 it seems to be unlikely for Co(N03)2, CoNaY and impossible at all for CoWl2, in which Co(II) occupies the center of the HPA and hence is not capable to form oxo-species. Remember that CoWl2, contrary to PWllCO, was found to be inactive in the alkene epoxidation by PhIO [18]. UV-Vis data confirm the stability of the Keggin structure of the HPA during the reaction with O2/IBA [33]. The fact that Co(NO3)2, CoNaY and CoWl2 appeared to be good catalysts for the epoxidation with OJIBA indicate that the ability of Co(II) to form complexes with dioxygen or to form metal-oxo species is not the determinant reason of the high catalytic activity of cobalt in Mukaiyama's system. Moreover, the pronounced activity of CoWl2, the well-known reagent of the outer-sphere electron transfer processes [34], shows that preliminary coordination of dioxygen and/or aldehyde is not necessary for the reaction to proceed successfully. Table 3 PIBA decomposition in the presence of Co(II) catalysts (a) Catalyst

PWllCO

COW12

Co(NO3)2 CoNaY (c)

CoPc

PIBA conversion(b) (%)

68

73

39

23

40

(a) Reaction conditions: PIBA 0.28 mmol, catalyst 6.10 .3 mmol, MeCN 6 ml, 24~ Cb)determined by iodometric titration after 2 h; (c) CoNaY 11 mg (Co 3.29%). To clarify the mechanism of the catalytic action of Co(II) compounds we have studied their activity in decomposition of perisobutyric acid (PIBA), formation of which was detected by IH NMR during alkene-IBA co-oxidation [16, 17, 19]. The ability of Co(II) compounds to mediate homolytic decomposition of peroxy acids was mentioned earlier [19, 25, 26, 35, 36]. The data on the activity of Co(II) compounds in the PIBA decomposition are summarized in Table 3. PWllCO and CoWl2 display the highest activity, while CoPc does the lowest one. The catalytic activity of Co(II) compounds in the PIBA decomposition seems to correlate with the rates of the epoxidation in O2/IBA/Co(II) systems. We have recently found an analogous correlation for O2/IBA/PWllM systems, where M = Co(II), Mn(II), Cu(II), Pd(II), Ti(IV), Ru(IV) and V(V) [19]. Taking into account the radical chain mechanism of the alkenealdehyde co-oxidation, we concluded that superior catalytic activity of cobalt compounds, at least in part, arises from their ability to catalyze the chain branching via the homolytic decomposition of peroxy acid formed during aldehyde autoxidation. Moreover, cobalt most probably takes part in the chain initiation via the reaction with aldehyde. The initiation step was proposed to proceed via the formation of the oxygenated adduct, like CoIII-O-O~ which reacts with aldehyde in the rate-determining step [26]. However, the fact that the induction time is more for CoPc as compared to CoWl2 (the latter is unable to form any adducts), shows that species different from oxygenated cobalt adducts may promote the chain initiation. We believe that for most cobalt catalysts studied these species are Co(III) forms of the catalyst produced by one-electron oxidation of Co(II) initial forms with peroxy acid, which in turn is produced in the course of aldehyde autoxidation. Usually, the end of the induction period coincides with the change of the reaction mixture colour expected for Co(III) appearance. The redox potentials (E) for the majority of the studied Co(II) catalysts are unknown for MeCN medium. In aqueous solution E is known to be higher for PWllCO as compared to CoWI2 (1, 39 and 1.0 V vs SHE, respectively) [37, 38]. Among the catalysts studied, CoPc

953 most likely has the lowest E (0.61 V vs SHE in DMF [39]). Note, that the induction time falls in the order: PWllCO > CoW~2 > CoPc. Thus, we may assume that the greater E is, the higher is the rate of the chain initiation. On the other hand, the chain termination via the reaction of RCO; radicals with Co(II) (vide supra) seems to be more favorable for compounds with low E. Coordinated acylperoxy radicals were proven to act as the epoxidizing species when Mn(III) Salen complexes were used as the catalysts [11 ]. We suppose that similar species may be involved in the epoxidation process in the case of CoPc or some other chelate complexes with low E. In this case one might expect some effects of a catalyst on the stereoselectivity of the epoxidation as it was observed in [11, 14] and in this work (Table 2). At the same time, acylperoxy radicals are expected to be the main epoxidizing species for cobalt compounds having relatively high values of E. Electrochemical and UV-Vis studies, which are in progress now, should provide further understanding of the mechanism of the alkene-aldehyde catalytic co-oxidation. The results obtained allow us to propose the reaction mechanism comprising the following elementary steps of the chain radical process leading to epoxide and isobutyric acid formation: RCHO+Co

3+

RCO + 02 ~

~RCO+Co

2+ H+ +

RCO 3

(1)

B

RCO~ + -- ~

RCOf/

(6)

(2)

RCO3+ RCHO ---~ RCO3H+ RCO

(3)

RCO~/ - - " ---~ , / ~

RCO3H+ Co22---~ RCO)+ Co3+ + OH"

(4)

RCO2 + RCHO ~

RCO~H + Co3+ ----~ RCO ~+ C2++ H+

(5)

2RCO~ ~

+ RCO~

(7)

RCO 2H + RCO

(8)

termination

(9)

3. CONCLUSIONS The results obtained in this investigation prove that alkene epoxidation by 02 in the presence of IBA and cobalt catalysts proceeds via radical chain mechanism. Acylperoxy radicals most likely act as the main epoxidizing species although some other species, e.g., coordinated to the metal center acylperoxy radicals, may contribute into the epoxidation process when catalysts with low redox potentials are used. Superior catalytic activity of cobalt compounds in alkene epoxidation by O2/IBA system is due to the high ability of cobalt to catalyze the chain branching and promote the chain initiation rather than the ability of cobalt to activate dioxygen via its coordination. The presence of chelating organic ligands is not necessary to provide the efficient alkene epoxidation in O2/IBA/Co(II) systems but such ligands are needed when the effect on the epoxidation stereochemistry is desired.

4. EXPERIMENTAL Catalysts. Co(NO3)2.6H20 was of pure grade. CoNaY zeolite and tetra-n-butylammonium salts of PWllCO and CoWl2 heteropolyanions were obtained as described in [19] and [25], respectively. The content of cobalt in the CoNaY zeolite was 3.29 % wt. The formation of the Keggin structure and the purity of CoW~2 and PWllCO were confirmed by 170 and 31p NMR, respectively. Materials. Trans-stilbene (Fluka AG) and (+)-o~-pinene (Aldrich Chemical Company) were used as received. Cis-stilbene and perisobutyric acid (PIBA) were prepared as described

954 in [19]. (-)-Caryophyllene (>99%) was isolated from the oil of Eugenia caryopyllata by vacuum rectification. (+)-3-Carene (95%) was prepared by rectification of the Pinus sylvestris turpentine. Oxidation procedure. Alkene oxidation was carried out in a thermostated 20 ml Pyrexglass reactor equipped with a stirring bar and a reflux condenser. Isobutyraldehyde was added to a solution of alkene (0.1 M) in a solvent (3 ml) containing a catalyst, and the reaction mixture was vigorously stirred. Product analysis. The oxidation process was monitored by GLC ("Tsvet-500", 2mx3mm Carbowax 20M on Chromaton N-AW-HMDS for stilbenes and 15mx0.3mm SE-30 for other alkenes, Ar, FID). The reaction mixture was percolated through alumina (/=2 cm, Q=I cm), concentrated at reduced pressure and the crude product was analyzed by ~H and ~3C NMR on a Bruker AM 400 instrument. The reaction products were identified from their NMR and GCMS spectra. The yields of epoxides were determined as described in [25]. ACKNOWLEDGMENTS

This work was supported by Russian Basic Research Foundation (Grant N 96-03-34215). We thank Prof. M. A. Fedotov and Dr. A. V. Golovin for NMR measurements. A generous gift of cobalt phtalocyanine from Prof. E. N. Savinov is highly appreciated. REFERENCES

1. R.A. Sheldon and J.K. Kochi, Metal-Catalyzed Oxidations of Organic Compounds, Academic Press, New York, 1981. 2. C.L. Hill. (ed.), Activation and Functionalization of Alkenes, Wiley, New York, 1989. 3. K.A. Jorgensen, Chem. Rev. 89 (1989) 431. 4. B. Meunier, Chem. Rev. 92 (1992) 1411. 5. N. M. Emanuel, E. T. Denisov and Z. K. Maizus, Chain Reactions of Hydrocarbon Oxidation in Liquid Phase, Moscow, Nauka, 1965. 6. T. Yamada, T. Takai, O. Rhode and T. Mukaiyama, Bull. Chem. Soc. Jpn., 64 (1991) 2109. 7. T. Takai, E. Hata, T. Yamada and T. Mukaiyama, Bull. Chem. Soc. Jpn., 64 (1991) 2513. 8. T. Mukaiyama and T. Yamada, Bull. Chem. Soc. Jpn., 68 (1995) 17. 9. R.V. Kucher and I.A. Opeida, Co-oxidation of Organic Compounds in Liquid Phase, Kiev, Naukova Dumka, 1989. 10. T. V. Filippova and E. A. Blyumberg, Uspekhi Khimii, 51 (1982) 1017 (in Russian). 11. Y. Katsuki, Coord. Chem. Reviews 140 (1995) 189. 12. S. Bhatia, Y. Punniyamurthy, B. Bhatia and J. Iqbal, Tetrahedron, 49 (1993) 6101. 13. P. Mastrorilli and C.F. Nobile, J. Mol. Catal., 94 (1994) 19. 14. P. Mastrorilli, C.F. Nobile, G.P. Suranna and L. Lopez, Tetrahedron, 51 (1995) 7943. 15. N. Mizuno, T. Hirose, M. Iwamoto, in V.C. Corberan and S.V. Bellon (Eds.), New Developments in Selective Oxidation II, Elsevier Science B.V., 1994, p. 593. 16. N. Mizuno, T. Hirose, M. Tateishi and M. Iwamoto, Chem. Lett. (1993) 1839.

955 17. M. Hamamoto, K. Nakayama, Y. Nishiyama and Y. Ishii, J. Org. Chem., 58(1993) 6421. 18. O.A. Kholdeeva, V.A. Grigoriev, G.M.Maksimov and K.I. Zamaraev, Topic in Catalysis, 3 (1996) 313. 19. O.A. Kholdeeva, V.A. Grigoriev, G.M. Maksimov, M.A. Fedotov, A.V. Golovin and K.I. Zamaraev, J. Mol. Catal. A, 114(1-3) (1996), 123. 20. S.-I. Mirahashi, Y. Oda and T. Naota, J. Am. Chem. Soc., 114 (1992) 7913. 21. S.-I. Mirahashi, Y. Oda, T. Naota and N. Komiya, J. Chem. Soc. Chem. Commun., (1993) 139. 22. E. Bouhlel, P. Laszlo, M. Levart, M.-T. Montaufier and G.P. Singh, Tetr. Lett., 34 (1993) 1123. 23. P. Laszlo and M. Levart, Yetr. Lett., 34 (1993) 1127. 24. A. Atlamsani, E. Pedraza, C. Potvin, E. Duprey, O. Mohammedi and J.-M. Bregeault, C.R. Acad. Sci. Paris, ser. II, 317 (1993) 757. 25. O.A. Kholdeeva, A.V. Ykachev, V.N. Romannikov, I.V. Khavrutskii and K.I. Zamaraev, Stud. Surf. Sci. Catal., 1997, in press. 26. J. Haber, Y. Mlodnicka and J. Poltowicz, J. Mol. Catal. 54 (1989) 451. 27. P. I. Valov, E. A. Blyumberg and N. M. Emanuel, (a) Bull. Acad. Sc. USSR (1966) 1283; (b) ibid. (1969) 718. 28. A.D. Vreugdenhil and H. Reit, Recl. Trav. Chim. Pays-Bas, 91 (1972) 237. 29. F. Marta, E. Boga and M. Matok, Disc. Far. Soc., (1968) 173. 30. Y. Nishida, T. Fujimoto and N. Tamake, Chem. Lett., (1992) 1291. 31. K. Kaneda, S. Haruna, T. Imanaka, M. Hamamoto, Y. Nishiyama and Y. Ishii, Tetr. Lett., 33 (1992) 6827. 32. C.L. Hill and R.B. Brown, J. Am. Chem. Soc., 108 (1986) 536. 33. Y. Shimura and R. Tsuchida, Bull. Chem. Soc. Jpn., 30 (1957) 502. 34. L. Eberson and L.-G. Wistrand, Acta Chem. Scand. B34 (1980) 349. 35. T. I. Ikawa, T. Fukushima, M. Muto and T. Yanagihara, Can. J. Chem., 44 (1966) 1817. 36. M.A. Brook, J.R. Landsay Smith, R. Higgins and D. Lester, J. Chem. Soc. Perkin Trans., II (1985) 1049. 37. B.C. Rong and M.T. Pope, J. Am. Chem. Soc., 114 (1992) 2932. 38. Z. Amjad, J.-C. Brodovitch and A. McAuley, Can J. Chem., 55 (1977) 3581. 39. M.R. Tarasevich, K.A. Radushkina, Catalysis and Electrocatalysis by Metalloporphyrines, Nauka, Moskow (1982) 35.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

957

Epoxidation o f olefins over thermally stable polyimide-supported M o ( V I ) complexes J.H. Ahn a, J.C. Kim a,

S.K. Ihm b and D.C. Sherrington~

"Department of Chemical Engineering and RECAPT, Gyeongsang National University, 900, Kajwa-dong, Chinju 660-701, Korea* bDepartment of Chemical Engineering, Korea Advanced Institute of Science and Technology, 373-1, Kusung-dong, Taejeon 305-701, Korea* ~Department of Pure and Applied Chemistry, University of Strathclyde, 295 Cathedral Street, Glasgow G1 1XL, United Kingdom

Polyimide particulates carrying a functional group have been prepared by non-aqueous suspension polycondensation. Molybdenum(VI) complex has been supported on a functional polyimide bead and used as a catalyst in the liquid-phase epoxidation of cyclohexene with tert-butylhydroperoxide (TBHP), as oxygen source. The polyimidesupported Mo catalyst was highly active and selective, and has been recycled 10 times with no detectable loss of Mo from the support.

I. INTRODUCTION Polystyrene-based resins have been used widely as supports for metal complex catalysts and other reactive species. These polymers, however, have a drawback in their limited thermo-oxidative stability [1,2]. The scope for application is therefore restricted, particularly in polymer-supported transition metal complex oxidation catalysts [3]. Consequently there is a need for the development of polymer supports with a much higher intrinsic thermo-oxidative stability. Polybenzimidazoles and polyimides are likely candidates in this respect. Polybenzimidazole resin in a porous bead has been applied as a polymer support for homogeneous metal complexes [4-6]. The polybenzimidazole-supported Mo(VI) species showed retention of high activity in the epoxidation of propylene, but progressive loss of activity on recycling in the epoxidation of cyclohexene.

*This work was carried out by G-7 Environmental Technology Development Program and the financial support is gratefully acknowledged.

958 Unlike polybenzimidazole-based thermo-oxidatively stable supports, polyimides can be prepared under relatively mild conditions from starting materials of only low or modest cost [7,8]. Polyimide particulates were prepared in a bead form without functional [9,10]. We now report the synthesis of functional polyimide beads and their use as epoxidation catalyst supports. The presence of the functional group in the polyimides allows further chemical exploitation, particularly as a catalyst support capable of operating under rather severe oxidative conditions. In this work, polyimide-supported Mo(VI) complexes were prepared and employed as heterogeneous catalysts in the epoxidation of cyclohexene using t-butylhydroperoxide (TBHP), as the oxidant.

2. EXPERIMENTAL

2-1. Materials N,N'-Dimethylacetamide (DMAc) (Aldrich, HPLC grade) was used without further purification. Acetic anhydride (BDH) was pre-dried over anhydrous sodium acetate. Pyridine (Aldrich, anhydrous) was distilled from KOH prior to use. Poly(maleic anhydride-co-octadec-l-ene)(l:l) (Polysciences), as a polymeric stabilizer, was used as supplied. Paraffin oil (A.J. Beveridge Ltd., liquid paraffin 5LT) was used as a suspending medium. Pyromellitic dianhydride (1,2,4,5-benzenetetracarboxylic dianhydride) (Aldrich) was recrystallized from butan-2-one before use. p-Phenylenediamine (Aldrich) and 2,6-diaminopyridine (Aldrich) were recrystallized from ethanol. 3,5-Diaminobenzoic acid (Aldrich) and 2,5-diaminobenzene sulfonic acid (Aldrich) were recrystallized from water and heated at l l0~ under vacuum to remove water. 3,5-Diamino-l,2,4-triazole (Aldrich) and tris(2-aminoethyl)amine (Tokyo Kasei) were without further purification. 2.2. Suspension Polycondensation A procedure similar to that which we have already reported was employed [9,10]. This involves the preparation of a pre-polymer poly(amic acid) (PAA) solution in DMAc, followed by imidization in suspension in paraffin oil. A typical procedure for the preparation of linear functionalized spherical polyimide particulates was as follows. A round-bottomed 3-necked flask was flushed with N2 and charged with a diamine in DMAc. The diarnine was completely dissolved in DMAc. While solution was mechanically stirred, finely ground pyromellitic dianhydride (PMDA) was added to the mixture on an ice bath in small portions, and then stirring continued overnight at room temperature. Paraffin oil with poly(maleic anhydride-co-octadec-l-ene)(l:l) (0.5wt% in oil) as a suspension stabilizer was added to the flask. The PAA solution was suspended for 2hr at 60~ at the speed of 400rpm. After that, imidization was initiated by dropwise addition of a mixture of acetic anhydride (4.0 molar excess of PMDA used) and pyfidine (3.5 molar excess of PMDA used). After 24hr, the polyimide particulates were filtered, washed with dichloromethane and then dried at 80 ~ in a vacuum oven. To obtain crosslinked polyimide particulates tris(2-aminoethyl)amine as a crosslinking agent was added to the PAA solution suspended in paraffin oil. After 24hr, a mixture of acetic anhydride and pyridine was added to give crosslinked polyimide.

959

2.3. Preparation of polyimide-supported Mo complex The polyimide bead bearing a triazole residue was used to immobilize a stable and active Mo(VI) epoxidation catalyst. The polyimide was refluxed with molybdenyl acetylacetonate in ethanol for 3days. Upon completion, a polyimide-Mo complex catalyst was filtered and extracted with ethanol in a Soxhlet apparatus for 3 days. The supported complex was dried thoroughly under vacuum. The molybdenum content was measured by inductively coupled plasma (ICP) to be 1.08mmolg~ for PI-DAT.Mo and 1.10mmolg~ for CPI-DAT.Mo. 2.4. Catalytic Epoxidation Catalyst (0.08g), cyclohexene (7.5ml), and bromobenzene (0.5ml) were placed in a three-necked thermostated reaction vessel equipped with condenser, septum cap, and stirrer, and left to thermally equilibrate at 60~ for 20min. Anhydrous TBHP solution (2ml, 5mmol TBHP) was added. Samples were withdrawn by syringe, and the concentration of cyclohexene oxide was monitored by gas chromatography (HP5890 Series II plus) with a capillary column (Ultra 2). 2.5. Analytic Methods Particle size distribution of functional polyimide particulates was determined by sieving: mesh 38/~m, 75/~m, 106/~m, 212/~m and 425/~m. Each size fraction was represented by wt%. FTIR spectra were recorded on a Nicolet SX20B instrument with KBr discs.

3. RESULTS AND DISCUSSION The suspension polycondensation methodology adopted has already been reported [9,10]. In this instance the pre-polymer poly(amic acid) solution in N,N'-dimethylacetamide was prepared from pyromellitic dianhydride and the functional diamines, 3,5-diamino- 1,2,4-triazole; 2,5-diaminobenzoic acid; 2,5-diarninobenzene sulfonic acid or 2,6-diaminopyridine. Each pre-polymer solution was then dispersed as droplets in paraffin oil containing poly(maleic anhydride-co-octadec-l-ene)(l:l) as a suspension stabilizer, and imidization induced at 60~ by addition of a mixture of acetic anhydride and pyridine (Figure 1). For comparison a nonfunctional polyimide (PI) was prepared using p-phenylene diamine and a crosslinked analogue of this (CPI) also produced by inclusion of tris(2-amino ethyl)amine. Typically 90---100% of mainly spherical polyimide particulates (---20g) were obtained after washing and drying (Figure 2). The results of elemental analyses and particle size distributions were shown in Table 1 and 2. For all species the H content found is higher than expected, probably reflecting trapping of solvent, fragments from the dehydrating agents, and moisture. Figure 3 shows the FTIR spectra of the triazole-containing polyimide bead (PI-DAT). The characteristic absorption bands were obtained at 1780, 1720(heterocyclic carbonyl vibration), 1348(C-N stretch vibration) and 720cm~(imide ring deformation). The thermogravimetric analysis curves for the polyimide beads are shown in Figure 4. The progressive loss below ~-300~ almost certainly corresponds to the physically trapped components mentioned above.

960 Serious degradation of all the functional polyimide particulates in air does not happen until --400~ All the polyimides therefore show good prospects for high temperature application as supports, certainly in reaction up to 300~ The homogeneous Mo complex, MoO2(acac)2, was supported on the PI-DAT bearing a triazole group. The FTIR spectra of PI-DAT.Mo (Figure 3) showed a band attributable to Mo=O stretching mode at 960cm1, indicating the presence of oxomolybdenum centers. There was no evidence in the IR spectrum of PI-DAT.Mo for an Mo-O-Mo bridge and so the structure of the Mo center in this polymer catalyst remains unclear. However, it is considered that all Mo derivatives on the supported catalysts probably contain oxomolybdenum centers from the initial rate data with no induction period (Figure 5). One of the authors has already reported on the possible structures of polymer-supported Mo complexes [11].

0~~

-t-

O

.o o

H2N-Ar-NH 2

A&.

-'~ N H - - ~ ~ O OH O PAA solution

Paraffin Liquid with a suspension stabilizer

PAA droplets

in oil

Acetic anhydride / Pyridine - H20 PI particulates

H PI--DAT

PI

COOH PI-COOH

SO3H PI-SO3H

PI-Py

Figure 1. Schematic synthesis of polyimide particulates. Table 1 Elemental analysis of polyimide particles CODE

Calculated(%)

Found(%)

C

H

PI

66.21

2.07

N 9.66

C 64.72

H 3.10

N 8.97

PI-DAT

58.85

1.59

10.16

45.79

2.93

12.63

CPI-DAT

49.56

3.05

17.15

52.86

2.99

16.80

PI-COOH

61.08

1.80

8.38

58.21

3.26

8.85

PI-SO3H

51.89

1.62

7.51

52.50

3.20

8.06

PI-Py

61.86

1.73

14.43

49.71

3.59

9.75

961

Table 2 Particle size distribution of polyimide particles Particle size fraction(wt%) a CODE A

B

C

D

E

F

PI

7.7

2.3

7.1

38.2

44.7

0

PI-DAT

0

3.0

5.6

34.8

41.4

15.2

CPI-DAT

0

0

0.3

1.5

80.4

17.8

PI-COOH

0

7.4

20.0

33.4

39.2

0

PI-SO3H

0

1.5

2.3

3.8

16.0

76.4

PI-Py

2.4

10.8

7.0

6.9

22.3

50.6

aA; < 38/~ m, B; 38-75/~ m, C; 75-105/1 m, D; 106-212/~ m, E; 212-425/~ m, F; >425/~ m

Figure 2. Optical Photograph of PI-DAT beads.

962

PI-DAT.Mo

PI-DAT

I

I

2000

I

1600

I

1200

I

800

400

Wave number(cm")

Figure 3. FTIR spectra of PI-DAT and PI-DAT.Mo.

100~

80

o~' v

60

rr

9

40

9

PI 9 PI-DAT PI-COOH

9 PI-SO3H 9 PI-Py

20

0

0

i

1

200

400

600

800

Temperarure(C) o

Figure 4. TGA analysis of functional polyimide particulates.

963 It is well known that the soluble species MoO2(acac)2 is a potent catalyst in the epoxidation of olefinic compounds. Figure 5 shows conversion curves for the epoxidation of cyclohexene by TBHP catalyzed by homogeneous MoOz(acac)2 and heterogeneous CPI-DAT.Mo. The initial activity of homogeneous MoO2(acac)2 was higher than that of CPI-DAT.Mo. However, the activity of CPI-DAT.Mo should be comparable with that of the homogeneous analogue after 30min.

CH3 I + H3C--C--OOH I CH3

Mo(VI)

,~-

[ ~

CH3 I O + H3C--C--OH I CH3

100

o-e,

~

60

0

o

40

~-

2o

o

0

20

40

60

80

1 O0

120

Time(min)

Figure 5. Epoxidation of cyclohexene with TBHP using CPI-DAT.Mo and homogeneous Mo. The prolonged activity of polymer-supported heterogenized catalysts is probably the most important factor in their performance. The deactivation of the catalyst by either degradation of the polymer-supported itself or by leaching of active species from the catalyst is unfavorable. Figure 6 shows the yield of cyclohexene oxide after 120min. The triazole-containing polyimide-supported Mo catalysts were recovered at the end of each run and used repeatedly under identical conditions. The catalyst shows substantial retention of activity over 10 recycles unlike an earlier polybenzimidazole-supported Mo complex, where the latter displayed rapid deactivation on recycling. The presently reported retention of activity is most encouraging, and suggests that catalysts based on functional polyimide particulates might form the basis of a range of stable polymer-

964

\ 60

\

/

o o

"6

9 9 9

40

"o

~--

C P I - D A T . M o 700C C P I - D A T . M o 60~ PI-DAT.Mo 60"C

20

I

i

I

I

I

I

I

I

i

1

2

3

4

5

6

7

8

9

10

R e c y c l e number

Figure 6. Recycling of polyimide-supported Mo catalysts in the epoxidation of cyclohexene by TBHP at 120min supported metal complex catalysts, where the support is readily synthesized and is highly cost-effective. Application on the both a laboratory and a technical scale also seems feasible. REFERENCES

1. P. Hodge and D.C. Sherrington (eds.), Polymer-supported Reactions in Organic Synthesis, Wiley-Interscience, Chichester, 1980. 2. Y.I. Yermakov, B.N. Kuznetsov and V.A. Zakharov, Catalysis by Supported Complex, Elsevier, Amsterdam, 1981. 3. D.C. Sherrington, Pure Appl. Chem., 60 (1988) 401. 4. H.G. Tang and D.C. Sherrington, J. Catal., 142 (1933) 540. 5. M.M. Miller and D.C. Sherrington, J. Catal. 152 (1995) 377. 6. M.M. Miller, D.C. Sherrington and S. Simpson, J. Chem. Soc. Perkin Trans. 2 (1994) 2091. 7. D. Wilson, H.D. Stenzenberger and P.M. Hergenrother (eds.), Polyimides, Chapman and Hall, New York, 1990. 8. L.H. Lee (ed.), Adhesives, Sealants and Coatings for Space and Harsh Environments, Polymer Science and Techology, Plenum, New York, 1988. 9. T. Brock and D.C. Sherrington, J. Mater. Chem., 1 (1991) 151. 10. T. Brock, D.C. Sherrington, and J. Swindell, J. Mater. Chem., 4 (1994) 229. 11. M.M. Miller and D.C. Sherrington, J. Catal., 152 (1995) 368.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

965

Selective Partial O x i d a t i o n of P r o p y l e n e to P r o p y l e n e Oxide on A u / T i - M C M Catalysts in the Presence of H y d r o g e n and O x y g e n Yuri A. Kalvachev* ", Toshio Hayashi b, Susumu Tsubota a and Masatake Haruta" aOsaka National Research Institute, AIST, Midorigaoka 1-8-31, Ikeda 563, Japan bNippon Shokubai Co., Ltd, Research Division, Nishi-Otabi 5-8, Suita 564, Japan 1. INTRODUCTION The epoxidation of alkenes is attracting increasing attention of researchers, from both academia and industry. The epoxides are one of the most useful intermediates in organic synthesis as they are versatile products that easily undergo ring-opening reactions to form bifunctional compounds. In the chemical industry, propylene oxide (PO) is mainly used for producing resins. Major conventional manufacturing methods for the synthesis of PO require a two-stage processes, using chlorhydrin or hydroperoxides. The direct synthesis of PO, by the use of oxygen, has long been considered desirable and is one of the most important reactions still not solved by catalysis [ 1]. The direct vapour-phase oxidation of propylene to PO in the presence of oxygen and hydrogen is an environmentally friendly process, or as J. Thomas once stated : "The name of the game is to get mild conditions and cheap oxidants that are environmentally friendly" [2]. Since direct oxidation is difficult to achieve, several approaches have been made to attempt to produce PO by oxidation with hydrogen peroxide over titalaosilicate catalyst [3,4]. In liquid phase Pd-Fe zeolite [5] and Pd-titanosilicate [6] are known to oxidize alkanes and alkenes to oxygenates by hydrogen peroxide generated in situ. New materials consisting of amorphous silica with regular pore structure, therefore called mesoporous molecular sieves, have recently been described [7]. Isomorphous substitution of Si by Ti has been attempted by performing the synthesis in the presence of titanium compounds. Ti-MCM have been tested for oxidation of hydrocarbons in liquid phase, using H202 or hydroperoxides as oxidants [8-10]. Because gold has long been regarded as being catalytically less active than platinum group metals it has attracted little attention in the development of heterogeneous catalysis. The basic reason is that gold catalysts are highly sensitive towards the preparation methods and it is normally impossible to prepare active gold catalysts with classical impregnation methods. However, when gold is dispersed as fine particles over suitable support by coprecipitation or deposition-precipitation methods, it has been found that the supported gold exhibits exceptionally high catalytic activity for such reactions as : CO oxidation [ 11-14], CO~ and CO hydrogenation [ 15], hydrocarbon combustion [ 16], the reduction of NO to N 2 [ 17]~ and the water-gas shift reaction [ 18,19]. Quite often reactions occur at low temperature. In previous work [20] the first evidence of the direct vapour-phase oxidation of propylene to propylene oxide was presented, using a catalyst, comprised of gold deposited on TiO 2.

*permanent address : Institute of Catalysis, Bulgarian Academy of Sciences, 1113 Sofia, Bulgaria

966 In continuation with our work, we now investigate the catalytic activity of gold deposited on titanium-containing MCM-41 for the partial oxidation of propylene to PO, in the presence of oxygen and hydrogen. In the present report, we have studied the influence of the amount of titanium in Ti-MCM-41 and gold loading over this reaction.

2. EXPERIMENTAL 2.1. Sample preparation Ti-MCM-41 samples were synthesized by using dodecyl trimethyl ammonium chloride as a template following the procedure described in [8]. The MCM supported gold catalysts were prepared by deposition-precipitation [21] on:pure-silica MCM-41 ; Ti-MCM-41 with ratio Ti/Si=l/100 ; Ti/Si=3/100 ; Ti/Si=6/100 by using HAuC14, followed by washing, drying and calcination in air at 673 K for 4 h. TEM micrographs showed that gold particles are homogeneously dispersed on the support, with an average diameters around 2 nm. The particle size distributions were obtained by TEM (Hitachi H-9000NA). Au-TiOJTi-MCM-41 were prepared by the same method by replacing Ti-MCM with TiOJTi-MCM, which was prepared by impregnating Ti-MCM with titanyl acetyl acetonate, followed by drying and calcination ha air at 773 K. The same procedures were also applied to Au-TiO2/SiO 2 catalysts. 2.2. Catalytic activity The catalytic activity of the samples was measured in a flow reactor at atmospheric pressure over the temperature range of 323-393 K. A gaseous mixture of propylene, oxygen, hydrogen and argon as diluent was passed through a fixed bed containing catalyst sample (Ar:O::H2:Pr = 7:1:1:1). Samples were sieved between 70 and 120 meshes and weighed to 0.5 g. Catalysts were pretreated in mixture of Ar:O 2 (7:1) at 573 K for 1 hour. The space velocity was 4000 h1. Reaction products were analyzed by gas-chromatography. 3. RESULTS AND DISCUSSION Table 1 lists the measured values of selectivity to PO, propylene conversion and hydrogen conversion for the reaction of propylene with hydrogen and oxygen over AuFFi-MCM catalysts at 323 and 373 K. The conversion of propylene is about 1-2 %, but the selectivity to PO is very high, in most cases above 90 %. Carbon dioxide is the main by-product. The most active catalyst in this reaction was Au/Ti-MCM with a Ti/Si ratio of 3/100. A sample with the same gold loading on silica MCM-41, which did not contain titanium, is catalytically inactive. Ti-MCM (31100) sample without gold loading is also inactive under these reaction conditions. This means that both components - gold and titanium are necessary for the oxidation of propylene, and that there is most probably a specific interaction between gold and titanium. A synergistic behavior of gold deposited on metal oxides is observed for lowtemperature oxidation of CO [22] and low-temperature water-gas shift reaction [ 19]. Fig.1 (a,b) shows the results for propylene oxidation over 8 wt%AtffTi-MCM with Ti/Si ratio 1/100 and 3/100, respectively. In Fig. 2 (a,b) the results from oxidation of propylene at 373 K on Ti-MCM (Ti/Si=3/100) with 4 and 12 wt % gold loading are presented. The PO yield is the highest on 8 wt % Au/Ti-MCM with aPO selectivity of 95 % (Fig.lb). On 12 wt % Au/Ti-MCM conversion of propylene is higher, but PO selectivity is 80 % . Thus the yield of propylene oxide is lower than that observed on 8 wt % Au sample. The catalytic activity increased with time over 2-3 hours (at 323 K) and 70-90 min (at 373 K) followed by a period of relative stability. Naito et al. [23] have suggested that oxygen molecules on small gold particles behave as a peroxo-like adsorbed species, which enhances the dissociation of hydrogen molecules. Moreover, the products are similar to those obtained

967 TABLE 1 Oxidation of propylene over Au/Ti-MCM catalysts Ti/Si ratio

T, K

gold

conversion

PO

loading

selectivity % propylene

323

1/100 3/100

9

18

4%

96

0.65

9

96

0.74

9

1.21

28

2%

91

0.36

11

4%

84

1.01

27

8%

91

0.36

13

8%

97

1.34

30

4%

94

0.96

27

8%

95

1.75

36

12%

80

1.89

50

6/100

a

0.45

90

373

1.20

97

8%

3/100

1.40

8%

12% 6/100

1/100

2%

80

0.88

47

4%

69

1.20

55

8%

80

1.13

26

9 9

9

uJ

1.80 -1.60 1.40 :

9 9 9

1.00 o _1 0.80

9

9 ua

* 9

>" 0.60 O

hydrogen

~

~.oo

b 000 9149149149149 O0

9

0.80

o 0.60

o. 0.40

0.40

0.20

0.20 +O

0.00

0.00 0

50

I O0 TIME

(MIN)

150

200

0

50

I O0 TIME

150

(MIN)

Figure 1. PO Yield in the oxidation of propylene at 373 K over 8 wt % Au/Ti-MCM as a function of time a) Ti/Si=l/100 ; b) Ti/Si=3/100.

200

968 in the oxidation of hydrocarbons by H202 on Ti-MCM [8]. We can thus speculate that the active species are a similar hydroperoxo-species, located at the boundary between gold and titanium. 1.00 T

1.40

a

0.90 -

9

0.80

9

0.70

1.20

9

9

9

uJ >" 0.50 o 0.40 " 0.30

9

1.00

9

o 0.60

41,

9

9 1 4 9 1 4 9c~ 0.80 u.i ~ 0.60 o

9

0

O~'OO0 0 9

0

~ 0.40

0.20

0.10

9

9

0.20

o.oo

0.oo o

50

~o 9 TIME

~5o

z0o

o

~

~

~

~

5o

~o 9

~5o

zoo

TIME

(MIN)

(MIN)

z50

Figure 2. P 9 yield in the reaction of propylene at 373 K over Au/Ti-MCM (Ti/Si=3/100) as a function of time a) 4 wt%Au ; b) 12 wt %Au. Fig.3 shows the P 9 yield as a function of time at different ratio of hydrogen : propylene in the gaseous mixture. When the concentration of H 2 is lower (H2:Pr=-5/10), the catalytic activity decreases and the shape of the curve with time is different- the activity increases over 6 hours. This may suggest that hydrogen is involved in forming the active species.

~.ZOT a

9

1.00

-o~ ~

OOOOOO0

0.80 "~ 0.60 o "

o

0.20 0.00

0.50 0.40

~176

0.30

***

0.40

b

"

o o.zo -

9

,,~

0.10

9

o

0.00 50

_O~

100

150

TIME (MIN)

200

250

f

0

100

200 TIME (MIN)

Figure 3. Oxidation of propyleneat 323 K over 12 wt % Au/Ti-MCM (Ti/Si=3/100) with a) H2:Pr=10/10 ; b)H~:Pr=5/10 vol.%.

300

400

969 The results presented in Fig.4 show that at the beginning of the reaction, the consumption of hydrogen is higher, whereas PO yield is low. Moreover, at low temperature (323 K) an induction period has been observed, during which there is not a formation of PO. We speculate that during this period a formation of hydroperoxide species takes place. 20.0

-

,r.,,

18.0 16.0 N -1-

14.0

"

12.0

0

9 10.0

8.0 bO

z O u

6.0 4.0

2.0 0.0 0

I

I

I

I

I

100

200

300

400

500

TIME

(MIN)

Figure 4. The consumption of hydrogen in the reaction of propylene with oxygen and hydrogen at 323 K over 2%Au/Ti-MCM (Ti/Si=6/100) In order to confirm this possibility, the following experiment was carried out - introducing only hydrogen and oxygen over the catalyst for the first 80 minutes. It turns out that before adding propylene, the hydrogen is oxidized 100 %. As it is seen in Fig.5 after the addition of propylene, the consumption of hydrogen decreases and PO yield increases with time. There is a competition between hydrogen and propylene in the oxidation reaction. The gradual increase in PO yield over a period of 70 min indicates that in the absence of propylene, hydroperoxidelike species are not formed or decomposed very rapidly. 0.60 T

I00.0 90.0

0,0

0.50

80.0 e,i 70.0 "-r" u. 60.0 0 50.0 40.0 30.0 Z

0.40 r~ _I

ua 0.30 O

~. 0.20

0

0.10

-4,0,,,,,, __ b

9

00000,0

20.0 10.0 o.o

0.00 5o

~o o TIME

(MIN)

~s o

zoo

o

so

~o o TIME

~s o

zoo

(MIN)

Figure 5. PO yield (a) and hydrogen consumption (b) at 373 K on 4%Au/Ti-MCM (Ti/Si=6/100).

970 Fig.6 shows PO yield over 8 wt % Au/TiOJSiO 2 as a function of time. The catalytic activity of Au/TiOJSiO 2 catalyst is not stable. Water is continuously formed during the oxidation of propylene and the oxygenated intermediates may block the active sites and depress the adsorption of propylene on the surface of the catalyst. MCM materials have hydrophobic character and Ti-MCM preferentially adsorbs less polar olefin molecules. This decreases the competition from water and probably avoid the accumulation of the oxygenated intermediates to lead to more stable catalytic activity. 2.00

9 9

1.80

Oo

1.60 1.4o

00

000 O0

o, 1.20 "' 1.00

O0 0

0.80 0.60 0.40 0.20 0.00 0

I

I

I

I

so

lOO

1 so

zoo

TIME

(MIN)

Figure 6. PO yield over 8 wt%Au/TiOJSiO 2 at 373 K. Table 2. Oxidation of propylene on Au/1 wt%Ti02/Ti-MCM-41 catalysts Ti/Si ratio*

T, K

gold loading

PO selectivity %

conversion % propylene

silica MCM

323

8%

hydrogen

inactive

1/100

8%

90

1.10

22

3/100

8%

90

0.65

16

8%

87

0.93

17

1/100

8%

74

2.45

53

3/100

8%

84

1.33

36

silica MCM

373

*this ratio is for Ti-MCM support

971 In Table 2 the results for the reaction of propylene with hydrogen and oxygen over AtffTiO2/Ti-MCM are presented. The most active sample was Ti/Si with a ratio of 1/100. From obtained data it can be concluded that for oxidation of propylene there exists an optimum amount of titanium in the catalyst. This may suggest that the density of the active sites influences the catalytic activity and PO selectivity. The results show that the optimum temperature for selective pal~ial oxidation of propylene to propylene oxide is about 100~ With increasing temperature, the conversion of propylene increased but PO selectivity was then lower. The conversion of propylene on these catalysts is about 1-2 %, while conversion of hydrogen, in all cases, is high (Tables 1 and 2). Probably the oxidation of propylene occurs at the perimeter interface between gold and the support, whereas hydrogen is activated on the surface of the gold particles. 4. C O N C L U S I O N S There exist typical characteristics features in the reaction of propylene with hydrogen and oxygen over Au/Ti-MCM catalysts: - an increase of PO yield over 2-3 hours (at 323 K) and 70-90 min (at 373 K), followed by a period of relative stability ; - increasing PO selectivity and decreasing hydrogen consumption as a function of time ; - PO selectivity decreases with increasing the amount of titanium in the catalysts. - in all cases the conversion of hydrogen is high (10-50%). The results point to the formation of hydroperoxide species that are the active oxidant agents. The reaction occurs at the interface perimeter of the support around the gold particles. The activity of Au/Ti-MCM-41 in the reaction of propylene with hydrogen and oxygen can be attlibuted to synergetic effects between gold and titanium in these catalysts. The hydrophobic character of MCM molecular sieves is the probable reason for the stability of the activity with time.

A C K N O W L E D G M E N T

Yu.A.K. gratefully acknowledges financial support by the Science and Technology Agency of Japan.

R E F E R E N C E S

1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11.

J.F.Roth, Chemtech, (1991) 357. M. Freemantle, C&EN, 74(31) (1996) 47. M.Clerici, G.Bellussi, and U.Romano, J.Catal., 129 (1991) 159. B.Notari, Adv. in Catal., 41 (1996) 253. N.Herron and C.Tolman, J.Am.Chem.Soc., 109 (1987) 2837. T.Tatsumi, K.Yuasa, and H.Tominaga, J.Chem.Soc., Chem.Commun., (1992) 1446. C.Kresge, M.Leonovicz, W.Roth, J.Vartuli, and J.Beck, Nature, 359 (1992) 710. A.Corma, M.Navarro, and J.Perez-Pariente, J.Chem.Soc.,Chem.Commun., (1994) 147. P.Tanev, M.Chibwe, and T.Pinnavaia, Nature, 368 (1994) 321. T.Blasco, A.Corma, M.Navarro, and J.Perez-Pariente, J.Catal., 156 (1995) 65. M.Haruta, S.Tsubota, T.Kobayashi, H.Kageyama, M.Genet, and B.Delmon, J.Catal., 144 (1993) 175. 12. M.Bollinger and M.Vannice, Appl.Catal., B:Environmental, 8 (1996) 417. 13. S.Gardner, G.Hoflund, B.Upchurch, B.Schryer, E.Kielin, and J.Schryer, J.Catal., 129 (1991) 114.

972 14. 15. 16. 17. 18. 19. 20. 21. 22. 23.

S.Tanielyan and R.Augustine, Appl.Catal. A:General, 85 (1992) 73. H.Sakurai and M.Haruta, Appl.Catal. A:General, 127 (1995) 93. M.Haruta, A.Ueda, S.Tsubota, and M.R.Torres Sanchez, Catal.Today, 29 (1996) 443. A.Ueda, T.Oshima, and M.Haruta, Appl.Catal. B:Environmental, in press. D.Andreeva, V.Idakiev, T.Tabakova, A.Andreev, and R.Giovanoli, Appl.Catal. A:General, 134 (1996) 275. D.Andreeva, V.Idakiev, T.Tabakova, and A.Andreev, J.Catal., 158 (1996) 354. T.Hayashi, K.Tanaka, and M.Haruta, Preprints of Symposia on Heterogeneous Oxidation, 21 lth National meeting of Amer.Chem.Soc., New Orleans, 1996, pp. 71-74. S.Tsubota, D.Cunningham, Y.Bando, and M.Haruta, "Preparation of Catalysts VI", G.Poncelet, J.Martens, B.Delmon, P.A.Jacobs and P.Grange (eds.), Elsevier, (1995) 227. M.Haruta, S.Tsubota, A.Ueda, H.Sakurai, "New Aspects of Spillover Effect in Catalysis", T.Inui, K.Fujimoto, T.Uchijima, M.Masai (eds), Elsevier (1993) 45. S. Naito and M. Tanimoto, J.Chem.Soc., Chem.Commun., (1988) 832.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

973

Immobilization of t r i a z a c y c l o n o n a n e - t y p e metal complexes on inorganic supports via covalent linking: spectroscopy and catalytic a c t i v i t y in olefin o x i d a t i o n Y.V. Subba Rao, D.E. De Vos,* B. Wouters, P.J. Grobet and P.A. Jacobs Center for Surface Chemistry and Catalysis, Katholieke Universiteit Leuven, Kardinaal Mercierlaan 92, 3001 Heverlee (Belgium) Different approaches are tested in the covalent linking of the triazacyclononane (tacn) macrocycle to amorphous or mesoporous siliceous supports. The best catalytic results for the epoxidation of olefins are obtained with a tacn, attached via a 3-oxypropyl-2-hydroxypropyl spacer to the support. The organic structures on the surface are studied with TGA, TPD-MS, 13C-NMR and sorption measurements. ESR is used to probe the details of the metal binding on these surfaces. 1. I N T R O D U C T I O N Catalytic mono-oxygen transfer from first row transition metals to nucleophilic substrates has been the subject of intensive studies since the late seventies [1-2]. The classic procedures of porphyrin-catalyzed oxidations have however obvious disadvantages [3-6]. Chlorinated solvents are often used, either in a two phase system or as co-solvents to dissolve the porphyrin. The reaction mixtures are heavily colored. Catalyst recuperation is not obvious, and often the porphyrin doesn't even survive a single catalytic run. Several groups have attempted with varying success to inlprove the usability of porphyrins by diverse heterogenization techniques [7-10]. As an alternative to porphyrin and phthalocyanine catalysts, complexes of Mn and the cyclic triamine 1,4,7-trimethyl-l,4,7-triazacyclononane (tmtacn) clearly deserve more attention [11]. In acetone and at subambient temperature, the activity of Mn-tmtacn matches that of the more active porphyrins, with 1,000 turnovers within a few hours in the styrene epoxidation [12]. Moreover, Mntmtacn is colorless after reaction, and because of its relatively moderate price, it has even been commercialized for a short while in laundry powders [13]. A heterogeneous version of Mn-tmtacn would obviously offer even more advantages. We have proposed an immobilization of Mn-tmtacn based on zeolite We acknowledge support from K.[LL~dven (YVSR) and F.W.O. (DEDV and PJG). This work was performed in the frame of an interuniversitary attraction pole (I.U.A.P.) program Supramolecular Catalysis. E-mail: [email protected]

974 Y [14]. A major problem is however that this hydrophilic matrix attracts H202. This is a drawback from the peroxide efficiency viewpoint. Therefore a purely siliceous matrix seems more attractive. As pure SiO2 lacks ion exchange capacity, one has to revert to other immobilization strategies, such as the covalent route. The present paper investigates various routes for covalent a t t a c h m e n t of the tacn macrocycle to a pre-formed support matrix. Two different spacers are used to link the surface and the tacn: propyl (P), and glycidoxypropyl (GP). The affinity of the modified surface for metals is probed with the test ions Cu 2§ and Mn 2§ and styrene is the test substrate for the selective hydrocarbon oxidation. A preliminary note on this work has appeared [15]. 2. E X P E R I M E N T A L

MCM-41 was prepared following an existing procedure [16]. The quality of the synthesis was evaluated based on the diffractogram and the N2 sorption isotherm. The material was calcined at 823 K and stored in a desiccator to avoid rehydration. Silica was purchased from Fluka (70-230 mesh) and pretreated under vacuum. 3-Chloropropylsilica was from Aldrich. For the anchoring of the organosilane on the Si matrix, 4.5 mmol of (3glycidyloxypropyl)trimethoxysilane was reacted during 10 h with 3 g of dry support material in 25 ml pre-dried and refluxing toluene. Excess silane was removed by toluene soxhlet extraction. For the reaction of tacn with GP- or Pbearing materials, 80 mg tacn was reacted overnight with 1 g of the vacuumdried support at 323 K in 50 ml toluene, followed by another toluene extraction. Eventually the remaining secondary amine groups on the bound tacn were alkylated with an estimated 5-fold excess of propylene oxide (ethanol, 293 K, 24 h). An overview is given in Scheme 1. H I

S c h e m e 1.

~o

,",,,~o

H i

H

H,,~'-~~H

/'~,/NCl

H i

~OH

H I

..

.

HOL

975 For TGA, a S e t a r a m TGA-DTA 92 a p p a r a t u s was used. Alternatively, a homebuilt a p p a r a t u s was employed for TPD-MS. GC analysis was on a Chrompack CP-Sil-5 column, eventually coupled to a Fisons mass spectrometer. ESR spectra were recorded with a Bruker ESP-300 and a TEl04 cavity at t e m p e r a t u r e s between 130 and 300 K. N~ sorption experiments were performed with an Omnisorp-100 i n s t r u m e n t . The t-plot method was applied for the analysis of the pore volume. Solid state 13C NMR spectra were recorded using a B r u k e r MSL 400 spectrometer at a resonance frequency of 100.61 MHz. Cross polarization was optimized with glycine as a reference. For the m e a s u r e m e n t of liquid samples, a B r u k e r AMX 300 system was used, operating at 300.13 and 75.47 MHz for 1H and ~3C, respectively. 3. R E S U L T S

Characterization of the functionalized surfaces T h e r m o g r a v i m e t r i c analysis is a basic technique for quantifying surface loading with organic groups. Samples were heated at 5 K per minute up to 1073 K in a He/O~ atmosphere. The weight loss (in %) above 453 K is given for all samples in Table 1. The organic weight increase after the reaction with tacn shows t h a t the linking is successful both for 3-chloropropyl (P) and for glycidyloxypropyl (GP) residues. Tacn surface concentrations are highest with MCM-GP and lowest for the commercial Sil-P. The TGA profiles are highly similar for tacn-containing samples (Sil-P-tacn, Sil-GP-tacn, MCM-GP-tacn) on one h a n d and the tacn-free precursors (Sil-P, Sil-GP, MCM-GP) on the other h a n d (Figure 1). With the latter materials, one main and sharp exotherm is observed around 473 K. With all tacn-containing samples, the combustion occurs over a much broader t e m p e r a t u r e interval (473-823 K).

~

,

Exo

_

-5

-5

Exo I

a

a

-15

2o0

40o

600 (c~

-35 ~,,

200 I

400 I

600 I

(el

F i g u r e 1. Weight loss (%, a) and heat flow (b) for MCM-GP (left) and MCM-GPtacn (right).

976 T a b l e 1. Surface loading (wt. % or tacn concentration) as d e t e r m i n e d by TGA.

Sample

% wt. loss

Sil-P Sil-P-tacn Sil-GP Sil-GP-tacn MCM-GP MCM-GP-tacn

[tacn] (retool / g) 0.20 0.28 0.40

5.9 8.5 12.3 15.9 12.2 17.3

W h e n this t h e r m a l decomposition is assessed by mass spectroscopy, typical f r a g m e n t s of decomposition of e.g. GP are detected. For instance, m/z = 57 is probably due to a 2,3-epoxypropyl group. 13C-NMR can be applied to check the intact nature of the surface groups after the anchoring and the extractions. As an example, we discuss the d a t a for the glycidylated MCM-41 (MCM-GP). *

a

~t

i

,!!

'i

I,,~~i,

!

!

,

i

t I :<,,,',~/4,,j ,,,~~\A,4,f,.,.,v/~,Wr (7, ' ' );i:~

<<~0

i

~i,

ii

W v'W~

r

,

l

T

r

--=

4c}

....

(}

l

F i g u r e 2. 13C-NMR spectra of MCM-GP (a) 1H-decoupled, (b) CP with 1H. Table

2. Comparison of 13C-NMR shifts for immobilized GP and its silane

precursor.

6 (ppm)

assignment

6 (ppm)

5.26

C2

9 (C2)

(liquid)

(liquid)

(solid)

......

0

6

[/~'... 5

7

~-

3

22.85

C3

23

(C3)

4"y~

44.28

C7

44

(C7)

~ i

50.52, 50.88 71.43, 73.55

C6 and C1 C4 and C5

50 (C6) 73.4 (C4,C5)

(CH30)3/

1

:2

977 In the proton-decoupled spectra, the sharp bands of the residual, liquid-like toluene are d o m i n a n t (marked with asterisks in Figure 2a). The signals of the immobilized species only become well-observable when cross-polarized spectra are recorded (Figure 2b). For comparison, the chemical shift values and the a s s i g n m e n t s for the GP silane precursor are also t a b u l a t e d in Table 2, and compared to the shifts of the anchored GP group. The only significant change (within the spectral resolution) is observed for C~, close to the anchoring site. Based on the signals of C6 and C7, the epoxide group is clearly intact after the anchoring. Finally coating of the surface with large fragments such as GP s u b s t a n t i a l l y influences the sorption properties of the material. Figure 3 compares the sorption isotherms of a single batch of MCM-41 before and after GP anchoring. From these plots, the total pore volume with a radius below 2 nm was shown to decrease from 1.8 ml.g -1 to 0.62 ml.g -] Moreover the m a x i m u m in the pore size distribution plot decreases from a pore radius r - 1.7 nm to r - 1.1 nm. V (ml/g)

V (ml/g) 1000

1000 500

500

0

0.5

P/Po

1.0

0

z

0

0.5

P/Po

1.0

F i g u r e 3. N~ sorption isotherms for calcined MCM-41 (left), MCM-GP (right).

Metal b i n d i n g on l i g a n d - f u n c t i o n a l i z e d

surfaces

To probe the metal-binding characteristics of the functional surface, Cu 2+ or Mn ~§ were introduced into the tacn-containing m a t e r i a l by stirring for 2 h in a solution of the corresponding sulfates in m e t h a n o l (90 %) - w a t e r (10%). As this process is f u n d a m e n t a l l y different from ion exchange, ESR was used to confirm the m e t a l u p t a k e and the selective binding of the m e t a l by the N ligand. For Cu ~§ N-coordination (as opposed to T a b l e 3. ESR p a r a m e t e r s for Cu ~§ systems an all oxygen coordination) is (X-band, 130 K). evidenced by a decrease of gll (from -2.40 to <2.30) and an increase of gll All (cm-1) All. The p a r a m e t e r s of Table 3 Cu-MCM-4i 2.409 0.0131 leave no doubt t h a t Cu ~§ binds Cu-Sil-P-tacn 2.289 0.0159 solely on the ligands and not on 2.223 0.0181 the silica or MCM-41 surface. Cu-MCM-GP-tacn 2.229 0.0176 Moreover comparison with [Cu(tacn)] ~§ 2.278 0.0158 literature values for [Cu(tacn)x] ~§ [Cu(tacn)~] ~§ 2.229 0.0177 complexes proves t h a t at low Cu ~§

978 concentrations, MCM-GP-tacn contains exclusively bis complexes, while also m o n o complexes are formed with the P spacer on silica [17]. Thus the length of the spacer seems to determine whether g=2 interaction between neighboring anchored ligands can occur. For Mn 2§ the ESR analysis requires a sample manipulation under N2 to avoid oxidation of the N-coordinated Mn. Nitrogen binding causes a decrease of the originally pseudo-octahedral symmetry of [Mn(H20)6] 2+. For a high-spin d 5 ion such as Mn 2§ this results in zero-field splitting. The non-zero values of the D and E parameters lead to new absorptions in the F i g u r e 4. X-band ESR for spectrum, at apparent g values different from 2 Mn-Sil-GP-tacn (293 K). ( ' n o n - c e n t r a l lines') [18]. We have recently applied such an analysis in the study of the Mn siting in zeolite A [19]. In ligand-free Mn-MCM-41 non-central lines are absent [20]. As an example, Figure 4 shows the X-band spectrum of Mn-Sil-GP-tacn, with non-central absorptions at 260 G, 1400 G and 2450 G (indicated by arrows). This proves that Mn binds selectively on the covalently attached ligand and not on the siliceous surface. Catalytic

activity

of the new

materials

Catalytic oxidations were performed with 1 mmol of substrate, 2 mmol of 35 % aqueous H202, 1 ml of solvent and 20 mg of metallated, functionalized support (containing 4-8 pmol Mn) at 273 K for 1 h. To improve the catalyst performance, 60 24 !ii!ii!~!iiiilililili!!iiiilit

T.O. 30 ~i~:~T~T~I 72 i:~:i:~i~!:?;:i?i~iii~ii~i~l

iiiiiiiiiii!iiiiiiii!iii!)i!iii

I ~1

.................

1

2

3

4

-

PO

-

PO

Sil-P-tacn

Sil-GP-tacn

87 87

5

62

liiiiiil

.................................

6

7

8

PO

PO

PO

MCM-GP-tacn

F i g u r e 5. Turnover numbers in the styrene oxidation with covalently anchored Mn-tacn. The solvent is acetone, except in 7 (CH3OH) and 8 (CH3CN). Numbers above bars are epoxide selectivities.

979 the effect of reacting the anchored tacn with two molecules of propylene oxide (PO) was tested. Turnover numbers in the oxidation of styrene are given in Figure 5 for the various systems. The post-modification with propylene oxide gives the best catalytic results, both with Sil-GP-tacn and MCM-GP-tacn. For the latter system, methanol seems a superior solvent in comparison with acetone and acetonitrile. The epoxide selectivity in reaction 7 is limited, but this is due to secondary reactions of initially formed epoxide to phenylacetaldehyde and the diol. When these secondary products are taken into account, the selectivity for the epoxide and its derived products increases to 76 %. 4. D I S C U S S I O N

The physicochemical characterization of the different intermediates and the final functionalized surfaces demonstrates that the glycidylation approach (Scheme 1) is a valuable alternative to the more familiar 3-chloropropyl method [21]. NMR proves that the oxirane group in the glycidyl residue is intact after the silane-surface c o u p l i n g . The reactivity of this oxirane group allows highly selective subsequent surface reactions under mild conditions. TGA and sorption experiments show the presence of considerable concentrations of GP or GP-tacn groups on the surface. ESR proves the specificity of the metal binding on these materials. Finally the catalytic experiments show that a tacn-type reactivity, with selective mono-oxygen transfer, subsists in the final material. While there is a far-reaching analogy between the catalytic activities of Mnporphyrins and Mn-tacn type complexes, the strategies for their covalent attachment are necessarily very different. The aromaticity of the porphyrin dictates that substitutions can only be made at the periphery of the ligand, which makes the influence of the substituent on the metal activity low. By contrast, the N-substituents in the tacn macrocycle are in the immediate proximity of the metal ion, and might even offer coordinating atoms to the metal center. We have previously studied the catalytic effects of changing the substituents on the tacn macrocycle [22]. With R=H (tacn), the system is almost completely inactive. With R=CH3 (tmtacn), the catalyst has a high activity, but only when acetone is used as a solvent. With hexadentate ligands (R=2-hydroxyalkyl or acetato) activities are somewhat lower, but still considerable when methanol is used as the solvent. For the present heterogeneous system (Mn-support-GP-tacn + PO), the structure and the solvent effects in the catalytic experiments resemble most those of the latter, hexadentate complexes. For such hexadentate complexes, a temporary removal of one of the pendant arms is necessary to create a coordinative vacancy on the metal. The particular role of methanol might be to assist in the temporary dehgation via hydrogen bond formation with 2-OH-alkyl groups. The system is unique in that the covalent link to the surface can participate in the metal coordination via the 2-hydroxy group, as indicated by the arrow in Scheme 1.

980 To our knowledge, this work is the first example of a covalently anchored, catalytically active non-heme Mn compound. Attempts are under way in our laboratory to develop similar catalysts with an even higher activity by immobilization of partially methylated tacn rings.

R E F E R E N C E S AND NOTES 1. B. Meunier, Chem. Rev. 92 (1992) 1411. 2. R. Holm, Chem. Rev. 87 (1987) 1401. 3. P. Battioni, J.P. Renaud, J.F. Bartoli, M. Reina-Artiles, M. Fort and D. Mansuy, J. Am. Chem. Soc. 110 (1988) 8462. 4. B. Meunier, E. Guilmet, M.E. de Carvalho and R. Poilblanc, J. Am. Chem. Soc. 106 (1984) 6668. 5. J.T. Groves and T. Nemo, J. Am. Chem. Soc. 105 (1983) 6243. 6. A.J. Castellino and T.C. Bruice, J. Am. Chem. Soc. 110 (1988) 158. 7. P. Battioni, E. Cardin, M. Louloudi, B. Sch611horn, G.A. Spyroulias, D. Mansuy and T.G. Traylor, Chem. Commun. (1996) 2037. 8. K. Miki and Y. Sato, Bull. Chem. Soc. Jpn. 66 (1993) 2385. 9. A. Sorokin and B. Meunier, J. Chem. Soc. Chem. Commun. (1994) 1799. 10. J.R. Lindsay Smith and R.J. Lower, J. Chem. Soc. Perkin Trans. 2 (1992) 2187. 11. D.E. De Vos and T. Bein, Chem. Commun. (1996) 917. 12. While acetone is less harmful than chlorinated solvents, one should beware of the organic peroxides that result from prolonged contact between acetone and H~O~. These peroxides are potentially explosive (e.g. in distillations). 13. R. Hage, J. Iburg, J. Kerschner, J. Koek, E. Lempers, R. Martens, U.S. Racherla, S.W. Russell, T~ Swarthoff, M. Van Vliet, J. Warnaar, L. van der Wolf and B. Krijnen, Nature 369 (1994) 637. 14. D.E. De Vos, J.L. Meinershagen and T. Bein, Angewandte Chemie, Int. Ed. Engl. 35 (1996) 2211. 15. Y.V.S. Rao, D.E. De Vos, T. Bein and P.A. Jacobs, Chem. Commun. (1997) in press. 16. J.S. Beck, J.C. Vartuli, W. Roth, M. Leonowicz, C.T. Kresge, K.D. Schmitt, C. Chu, D.H. Olson, E.W. Sheppard, S. McCullen, J.B. Higgins and J.L. Schenker, J. Am. Chem. Soc. 114 (1992) 10834. 17. R.D. Bereman, M.R. Churchill, P.M. Schraber and M.E. Winkler, Inorg. Chem. 18 (1979) 3122; P. Chaudhuri, K. Oder, K. Wieghardt, J. Weiss, J. Reedijk, W. Hinrichs, J. Wood, A. Ozarowski, H. Stratemeier and D. Reinen, Inorg. Chem. 25 (1986) 2951. 18. B. Bleaney and D.J.E. Ingram, Proc. Royal Soc. (London) A205 (1951) 336. 19. D.E. De Vos and T. Bein, J. Am. Chem. Soc. 118 (1996) 9615. 20. D. Zhao and D. Goldfarb, Stud. Surf. Sci. Catal. 97 (1995) 181. 21. D. Brunel, A. Cauvel, F. Fajula and F. DiRenzo, Stud. Surf. Sci. Catal. 97 (1995) 173. 22. D.E. De Vos and T. Bein, J. Organometallic Chem. 520 (1996) 195.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

981

Simultaneous Determination of Reaction Kinetics and Oxygen Activity during Selective Oxidation of an Aldehyde over an Oxidic Multicomponent Catalyst M. Estenfelder and H.-G. Lintz, Institut ftir Chemische Verfahrenstechnik, Universit~it Karlsruhe (TH), 76128 Karlsruhe, Germany

The partial oxidation of an aldehyde to a carboxylic acid has been studied over an oxidic multicomponent catalyst, mainly based on Mo, V and Cu. An experimental set up for the simultaneous determination of reaction kinetics and the oxygen activity of the catalyst is introduced. The kinetic measurements were performed by monitoring the gas phase composition along the length of a fixed bed of catalyst. The reactor was treated as an isothermal plug flow system. The reaction kinetics can be described with a simple triangle network consisting of the main reaction (aldehyde to carboxylic acid), a consecutive reaction (carboxylic acid to byproducts) and a parallel reaction (aldehyde to by-products). For the simultaneous measurement of the catalyst oxygen activity a so-called Solid Electrolyte Potentiometry (SEP) cell was additionally connected to the apparatus. The cell consisted of an oxygen-ion-conducting solid electrolyte (ZrO2 + 8.5 wt. % Y203) which was coated with a platinum reference electrode on one side and with a catalyst measuring electrode on the other side. The latter electrode was made of the same catalyst material used in the tubular reactor. While the measuring electrode was in contact with the gas phase to be analyzed, the reference electrode was flushed with air. The measured potential difference between both electrodes was correlated to the oxygen activity of the catalyst. From the determined axial profile of oxygen activity in the catalyst it can be concluded that the catalyst under working conditions is always in a reduced state, with the aldehyde acting as a stronger reducing agent than the carboxylic acid. It is also shown that there are two different oxidation states in the catalyst, each existing over a wide range of gas phase composition.

1. I N T R O D U C T I O N Multicomponent oxidic catalysts are commonly used in the partial oxidation of hydrocarbons and other organic compounds. They may show - at reasonably high activities extremely good selectivities for the products aimed at [ 1]. The efficiency of such catalysts and the synergy effects observed in the presence of one or more well defined oxides poses the question of their origin. Different explanations of these synergetic effects, the enhancement of both activity and selectivity have been proposed [2]. They cover the change of morphology of

982 the catalyst under the conditions of the catalytic reaction [3], the remote control process [4], which is due to spillover of oxygen from one to another oxide, and the interactions between phases [5]. Anyhow the availability and transfer of oxygen is a function of the phase composition of the catalyst, its oxidation state. Under operating conditions this oxidation state is not only determined by the oxygen partial pressure in the gas phase but results from the rates of oxygen transfer to and from the solid. There is a mutual interaction of the gas phase and the catalyzing solid. One may wonder how the latter is modified when the composition of the gas phase is totally changed from feed to product stream. One may even suspect that the use of multicomponent catalysts, often not justified by any topographic reason, is determined by the necessity to ,,buffer" the oxidation state independently of the gas phase composition. To characterize the oxidation state of a solid Carl Wagner [6] proposed the measurement of its oxygen activity by a potentiometric method. It is based on the use of an ion conducting solid as electrolyte. Two porous electrodes of a galvanic cell are separated gas-tight by the solid electrolyte. One electrode is exposed to the reacting gas mixture and acts simultaneously as the catalyst, the second electrode is in contact with a reference oxygen partial pressure. The measurement of the potential difference between these two electrodes leads to the value of oxygen activity in the catalyst which is operationally defined by this method, named Solid Electrolyte Potentiometry - SEP. In the following we present an experimental set up which allows the simultaneous determination of the gas phase composition and the oxygen activity of the solid catalyst under operating conditions along a tubular fixed bed reactor. Preliminary results of the partial oxidation of acrolein to acrylic acid at an oxidic catalyst illustrate the possibilities of the procedure.

2. E X P E R I M E N T A L 2.1. Kinetic measurements

The experimental set up (Fig.1.) can be divided into two parts, the fixed bed reactor for kinetic measurements and a second device for the Solid Electrolyte Potentiometry. The fixed bed reactor has been described in detail elsewhere [1,7]. The reaction mixture is admitted through thermal mass flow controllers, the concentrations of acrolein and water being fixed by loading a stream of nitrogen or air via two saturation-condensation systems operated at constant pressure and temperature. The kinetic experiments were carried out in a tubular reactor (1500 mm length, 15 mm i.d.) at a pressure of 1.4 bar and a temperature of 300~ A bed of catalyst consisting of 200 g spherical egg-shell catalysts was employed. The feed stream contained 4 mol % acrolein, 5.8 mol % oxygen, 5 mol % water and balance nitrogen at volume flow rates of 60 or 80 Nml/s. The reactor was divided into nine segments with separate electrical heating and individual temperature control to maintain isothermal operation. Heated capillaries were located at the inlet and at the outlet of each segment leading to the analytical section. Thus the gas phase composition along the length of the fixed bed could be monitored by distributed local sampling. In contrast to the set ups used before [1,7] each side stream selected by use of a multiposition valve entered a SEP cell prior to analysis by gas chromatography (organic components), infrared photometers (CO, CO2) and a magnetomechanic device (02).

983

Air

Nitrogen --~ Air Acrolein

__

~'~~ 9 ,

*e o

..-I

TO. t--

~-Waste Gas

~MI

~

Water

SEPI

TOR: Total Oxidation Reactor FBR: Fixed Bed Reactor MPV: Multi Position Valve

Analysis

Fig. 1. Schematic diagram of the experimental set up

2.2. Potentiometric measurements

The principles of the SEP measurements have been described in detail elsewhere [8]. The underlying assumptions for its use in the case of oxidic catalysts have equally been discussed [9]. A schematic drawing of the galvanic cell is shown in Fig.2.

ZrO2

Measuring Side

+

8.5 Gew.%Y20~

Oxidic Catalyst~

PBy

Platinum

'-'rP o d u c f s

20z---~-- 02+4e-

Reference Side

-- 0,79 bar I I

i

02+4e-~ 202-

AE Fig.2: The electrochemical cell: schematic diagram of the measuring arrangement

It consisted of a solid electrolyte disk (thickness 2 mm) made of yttria (8.5 wt. %) stabilized zirconia, which was coated with the same active component (cf. 2.3.) used in the reactor on the measuring side, and with a porous platinum electrode on the reference side. The measuring

984 electrode was in contact with the gas phase to be analysed and so simultaneously works as a catalyst. The reference electrode was flushed with air. Both electrodes were connected with a high ohmic volt meter. Temperature and pressure in the cell were the same as in the fixed bed reactor (T = 300 ~ p = 1.4 bar). It is the aim of the potentiometric measurements to determine the distribution of oxygen activity in the catalyst along the fixed bed in parallel to the concentration profile in the gas phase. In order to avoid the use of one galvanic cell at each sampling port the side streams collected passed sequentially over the cell shown in Fig.2. The catalyst electrode has been prepared in a way that the acrolein conversion in the cell was kept below 10 % in any case. Thus the measured potential difference between the electrodes is characteristic for the catalyst under the reaction conditions at the corresponding sample port of the tubular reactor.

2.3. Catalyst and electrode preparation A bed of catalyst consisting of 200 g spherical egg-shell catalysts was employed in the fixed bed reactor. The catalyst bed was diluted by shattered steatite particles (0.9 mm < d < 1.6 mm) in a mass ratio 1:1 to obtain a plug flow system. The catalyst used throughout the study was prepared by coating spherical steatite particles of 4-5 mm diameter with a porous oxidic layer. The egg-shell catalyst contained 20 weight % active component, the thickness of the shell being 215 lam. The oxidic catalyst consisted mainly of Mo, V and Cu, its preparation has been described elsewhere [10]. The reference electrode was prepared by coating the electrolyte with a thin layer (80 lttm) of a commercial platin paste and sintering the film for two hours at a temperature of 800~ The measuring electrode employed the same catalyst powder [ 10] which was used to prepare the egg-shell catalyst. It was mixed with an organic binder giving a viscous paste. The solid electrolyte was coated with this paste up to a film thickness of 80 lam and sintered for one hour at 320~ leading to a sufficient adhearence of the porous electrode on the electrolyte.

3. R E S U L T S AND DISCUSSION

3.1. Data evaluation It has been shown earlier that the kinetics of selective acrolein oxidation can be described by using a simple network of chemical reactions [7]. CO2, CO and acetic acid generated in small amounts can be lumped together into one pseudo-species (,,by-products"). The network consists of the main reaction from acrolein to acrylic acid, a parallel reaction of acrolein to byproducts and the consecutive reaction of acrylic acid to by-products.

985

kin1,2

Acrolein

Acrylic Acid

By-Products (CO, CO2, Acetic Acid ...) Fig. 3. Network of acrolein oxidation

The rate of the individual reactions in this system can be represented by the rate equations

kml,2 "Cacr

(1)

rm~'2 = 1 + b Cacr

krnl'3 "Cacr rn~l'3 = 1 + b- Cacr rm2,3 --

(2)

km2,3"Cacrylic acid

(3)

where rmi,j is in mol g-~s-~. In order to obtain the coefficients kmi,j and b a set of three simultaneous differential equations is solved numerically by use of the Runge-Kutta method and the kinetic constants are chosen after curve fitting of the calculated concentration profiles to the experimentally determined values. As stated above the oxygen activity ao 2 in the catalyst is operationally defined by the potentiometric measurement. Its value is related to the experimentally determined potential difference AlE via NERNST equation 2

A E = ~ . RT " In ~ a ~ 4-f Po2,Ref

(4)

where PO2,Ref designates the oxygen partial pressure at the reference side and F is the FARADAY constant (F = 96486 C mo1-1 ). It is based on the potential determining reaction which takes place at the three phase boundary line gas-electrode-electrolyte at both electrodes.

02,g Jr-4e-(Electrode) ~

202-(Electrolyte)

(R1)

986 If the adsorption equilibrium between the solid catalyst and the oxygen in the gas phase is attained, the oxygen activity in the solid is equal to the oxygen partial pressure in the gas phase. 2

(5)

a o - Po~

3.2. Discussion Typical results of the combined kinetic and potentiometric measurements are shown in figures 4 and 5. The experimentally determined values of the gas phase concentration and the oxygen activity are plotted against a modified residence time related to the mass of the catalytically active component. Fig. 4 shows the concentration profiles, yi representing a normalized dimensionless concentration of the three components. The solid lines are the results of the model calculations, obtained by the procedure described in the preceeding paragraph. The agreement of the simulation with the experimental measures justifies the description by a simplified network. It is worth remarking that for the temperature under consideration (T =300~ the value of kml,3 is equal to zero. This means that the main reaction of acrolein to acrylic acid is much faster than the parallel reaction of acrolein to by-products.

-50

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cm3/(g

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km2,3 = 0 . 3 4

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0 0

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0.15

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0.35

Modified g's

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Residence -3

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Figure 4: Product distribution and the natural logarithmn of (ao/Po2,Ref) 2 as a function of the modified residence time during selective acrolein oxidation. Symbols: experimental data, lines: calculated after curve fitting by use of rate constants given

987

The oxygen activity is plotted in a reduced form related to the value of the oxygen pressure on the reference side. A steady profile along the length of the catalyst bed was obtained. For acrolein conversions less than about 70 %, the oxygen activity in the catalyst is at a constant, low level independent of the composition of the gas phase. In the range of conversions between 70 and 100 %, the oxygen activity is increasing, reaching a second constant level for complete acrolein conversion. The determined axial profile of oxygen activity shows that the catalyst is always in a reduced state, with acrolein acting as a stronger reducing agent than the product acrylic acid.

1

k_.

mmm:n

.s >-. ,"L_

0.01

9

9

9

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1E-4 . . . . . . . . . . . . . . . . . . . . . . . . . . .

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:. . . . . . . . . . .

9

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IE-10

Po2

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,

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Catalysf:

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-.io

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..........................................

2

1E-12 1E-14 1E-16 1E-18

. . . . . .

.

.

.

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.

.

.

.

.

.

.

.

.

.

.

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9

T -

300~

p -

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X C3H40,in

--

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x 02,i,.,

=

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X H20,in

--

5%

! . . . . . . . . . . . . . . . . . . . . . . . . . . .

.

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. . . . . .

Mo-V-Cu-Ox

.

.

.

.

.

.

.

.

.

.

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.

.

.

.

.

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.

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.

.

.

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1E-20 0

0.05

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0.15

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0.4

0.45

Residence

0.5

Time

-3

Fig.5 9Oxygen activity corresponding to equation 4 and the measured oxygen partial pressure as a function of the modified residence time This is illustrated by Fig.5 where absolute values of the oxygen activity are plotted as a function of the residence time. They are orders of magnitude lower than the oxygen partial pressure of the corresponding gas phase over the catalyst, clearly indicating that oxygen adsorption is far away from equilibrium. It is interesting to note that the activity values show that acrylic acid is a weaker oxygen acceptor than acrolein. This is in good agreement with the fact that the rate constant of acrolein partial oxidation, kml,2 is three orders of magnitude higher than the rate constant km2,3 of the consecutive total oxidation of acrylic acid. As the latter is the partially oxidized product aimed at, this illustrates the quality of the catalyst used.

988 4. CONCLUSIONS AND OUTLOOK Gas phase composition and the oxidation state of the catalyst, represented by its oxygen activity, during selective oxidation of acrolein over an oxidic multicomponent catalyst were both monitored by simultaneous kinetic and potentiometric measurements. To our knowledge it is the first time that the oxygen activity profile along a bed of an oxidic catalyst during partial oxidation of an aldehyde could be determined by use of the Solid Electrolyte Potentiometry. The results show that in the case of acrolein partial oxidation there exist two domains of a nearly constant catalyst composition. This indicates that two values of the oxidation state of the catalyst are ,,buffered" in the system under consideration even if the composition of the oxygen acceptors (acrolein, acrylic acid) in the gas phase are dramatically changed. In the case when several of the phases of a multicomponent catalyst are identified and can be synthesized separatedly it may be promising to determine the oxygen activity profile of the single phases as well as of the multicomponent catalyst itself by the same experimental approach as described here. We can expect that the combination of the kinetic and potentiometric measurements will give a strong hint concerning the role of the different phases within the reaction network and may lead to practical correlations between the oxidation state of a catalyst and its kinetic behaviour (selectivity and activity) during oxidation.

REFERENCES [1] R. Recknagel and L. Riekert, Chem. Technik 46 (1994) 324 [2] J. Haber, Stud. Surf. Sci. Catal. 7__22(1992) 279

[3] J.M.M. Millet, H. Pouceblanc, G. Goudurier, J.M. Herrmann, J.C. Vedrine, J. Catal. 142 (1993) 381 [4] B. Delmon, L.T. Weng, Appl.Catal. A81 (1992) 141 [5] O. Legendre, Ph. J~iger, J.P. Brunelle, Stud. Surf. Sci. Catal. 7__22(1992) 387 [6] C. Wagner, Adv. Catal. 21 (1970) 323 [7] R. Recknagel, Thesis Karlsruhe (1994) [8] H.-G. Lintz, C. Vayenas, Angew. Chem. Int. Ed. Engl. 28 (1989) 708 [9] H.-H. Hildenbrand, H.-G. Lintz, Ber. Bunsenges. Phys. Chem. 95 (1991) 1191 [ 10] R. Krabetz et al., EP 17000, (1980)

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

989

On the m e c h a n i s m of the selective oxy-dehydrogenation of n-butenes to 1,3-butadiene on m a g n e s i u m ferrite: an FT-IR study. E. Finocchio, G. Busca, G. Ramis and V. Lorenzelli Istituto di Chimica, Facolt/~ di Ingegneria, Universit/~ di Genova, I-16129 Genova, Italy.

The adsorption of the linear C4 olefins (1-butene, cis-2-butene and trans-2-butene) and of 1,3-butadiene on MgFe204, and the evolution of the surface and gas-phase species upon heating has been followed by FT-IR spectroscopy. For comparison, the behavior of several C4 oxygenates has also been investigated. It has been found that the three linear butenes give rise by interaction with the oxidized catalyst sites to but-3-en-2-oxide species that later decompose to butadiene. The main pathway giving rise to total oxidation likely goes through but-2-en-3one that breaks to give different carboxylate species. 1. INTRODUCTION. 1,3-butadiene is a major intermediate in petrochemistry, due to the large production of rubbers based on its homopolymers and copolymers [1,2]. The synthesis of butadiene is mainly based on the pure or oxidative dehydrogenation of n-butenes from the C4 oil cracking fraction [1,2]. Among different processes industrially available for butadiene production, the n-butene oxy-dehydrogenation over ferrite catalysts like Mg and Zn ferrite, attracted much interest [35]. The oxy-dehydrogenation Of n-butenes to 1,3-butadiene belongs to the "allylic oxidation" reaction class, and should have, from the mechanistic point of view, features in common with other allylic oxidation reactions such as the synthesis of acrolein from propene [6]. Nevertheless, while ferrites are perform quite well for butadiene synthesis, they are not good catalysts for acrolein synthesis. We previously undertook a study of hydrocarbon activation over different transition metal based oxide catalysts mainly in relation to their total oxidation [7-9]. We proposed that hydrocarbons are activated at their weakest C-H bond on high-oxidation-state transition metal cationic centers with the formation of alkoxy species [7,8]. In the case of propene and 1butene we suggested that the primary surface intermediates are allyl-oxy species [7,8]. We applied our FT-IR technique also to selective oxidation catalysts [ 10]. We present here our results on an FT-IR study of the interaction and oxidative conversion of n-butenes over MgFe204, i.e. an active catalyst for butene oxy-dehydrogenation. The aim of the present work was to have information on the activation of olefins over allylic oxidation catalysts and to find information on the masons why ferrites are good catalysts for allylic oxy-dehydrogenation of butenes but they are are not good catalysts for allylic oxidation of propene. 2. E X P E R I M E N T A L

The preparation and characterization of the MgFe204 aerogel has been reported previously [11]. This material was characterized to be a crystallographically pure partly inverted spinel (MgFe204, magnesioferrite), with the surface area of 130 m2/g.

990 The IR spectra were recorded by a Nicolet Magna 750 Fourier transform instrument. The adsorption and oxidation experiments were performed using pressed disks of the pure powders, activated by outgassing at 300-1070 K into the IR cell. 3. RESULTS 3.1. The interaction of n-butenes on MgFe204. The FT-IR spectra of the species arising from the contact of the MgFe204 catalyst with 1butene are shown in Fig. 1, where the spectrum of 1-butene liquefied at 150 K over a KBr disk is also shown for comparison. The spectrum of the surface species formed at r.t. certainly contains features due to molecularly adsorbed 1-butene (Fig. 1,b), whose positions are summarized in Table I where they are also compared with the bands of 1-butene liquid (Fig. 1,a) and molecularly adsorbed on other oxide catalysts. Some modes are slightly shifted by the surface interaction to lower frequencies (1635 vs. 1641 cm -1, C=C stretching) or to higher frequencies (1001 vs. 994 cm -~ =CH2 twisting; 924,912 vs. 909 cm -~, =CH2 wagging, and the corresponding overtone at 1840, 1827 vs. 1825 cm-1). These perturbations are typical of olefin species when they interact through the re-orbital of their C=C olefinic double bond fiat on the surface, with electron withdrawing centres [12]. All the features that can be assigned to molecularly adsorbed 1-butene disappear almost completely by simple outgassing at r.t.. Additional strong bands appear in the range 1200-1000 cm -1, namely at 1189, 1153, 1089 and 1056 cm ~, and do not disappear upon outgassing at room temperature (Fig. 1,c). In these conditions the C=C stretching is possibly still present, very weak and partially masked by a broad adsorption due to carboxylate species or carbonate impurities (maximum near 1580 cm-~). In the C-H stretching region peaks after outgassing are found at 3082, 2973, 2935, 2905 and 2882 (shoulder) cm q, where the band at 3082 cm -~ confirms the persistence of an olefinic entity (=CH2 asymmetric stretching). Bands in the range range 1200-1000 cm -~ are typical of C-C and C-O stretchings of alkoxide-species, i.e. species involving a C-O single bond. Heating at 200 ~ and above causes the progressive disappearance of these bands too, while very strong absorptions grow near 1605, 1590, 1435, 1390, 1370 and 1355 cm -~. These bands are typically due to asymmetric stretchings (1605, 1590 cm-1), symmetric COO stretchings (1435 and 1370 cm -~) and C-H deformations (1355 and 1390 cm -~) of acetate (1605, 1435 and 1355 cm -1) and formate species (1590, 1390 and 1370 cm-1). The presence of a pronounced shoulder at 1455 cm ~ suggests that acrylate and/or propanoate species could also be present. In Fig. 2 the spectra of the surface species arising from adsorption of trans-2-butene and of but-3-en-2-ol on the same catalyst, are shown. The position of the bands of the species that disappear upon outgassing (Fig. 2,a), assigned to molecularly adsorbed trans-2-butene, are reported in Table II. Also in this case these species show perturbations typical of re-type interactions [12]. Additional strong bands are observed at r.t. (Fig. 2,b) that do not correspond to those of the molecular olefins and that do not disappear by outgassing at r.t.. These bands are located in the region 1200-1000 cm ~, where the typical strong absorptions of C-O and C-C stretchings of alkoxide species fall. By increasing the time of contact of n-butenes gas with the catalyst surface, these bands further increase in intensity. A similar result is obtained whan cis2-butene is adsorbed. The spectrum (Fig. 3,a) contains the features of molecular species (desorbed by outgassing at r.t.) and of alkoxide species (strong bands in the 1200-1000 cm q region), resistant to outgassing at r.t.. The spectra of the alkoxide species formed from the three linear butenes are almost superimposable, with only weak differences in the relative

991

intensities of some bands (compare Figs. 1,c, 2,b and 3,a), indicating that the same species, or two species with slightly different amounts ratios, are formed from the three butene isomers. " i .....................................

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/~

.... i

0.,8-:

"-',.,,

~1

\

!t

o.o'i

11,

ooo~-'-~ ~

!1

.

oo~,~

_

,, a

II

11

.

h,~.

.

J't,h

.

.

It

/ t .

.

9

s

i

t.i

",-..-,

,,

1

...... "'......-. .......--. ...................

"\ ~ , )

................................

[i

.... ~

il

,,o~

...J!

C

rVt,.~ Wavenumbers

................................................................

Y (cm-!)

"

\

..........................................

Fig. 1. FT-IR spectra of (a) liquid 1-butene film; (b) 1-butene reversibly adsorbed as such on MgFe204; (c) products of the reactive adsorption of 1-butene on MgFe204 at r.t. 0,12-'." 0,10~ 0,08~

0,04-"."

0,02!

o.oo5 a -

0

~ ,

0

M~~,~.~___~

2

,~) v ~

/~ . , /

v

-0,04 -0,06~

ooo!

oo83,

b

v

A/A A ~

/~

:

-0,14-"

~r~rT~n~1~1~1~1~`~.~~.f.~.r~r~r~rn~v~1~1.1~.1~1~.~.~,~.~.~r ",','," "1" "v'vn'T'l'~"l','l"~'l"l"l'l'l'l'l'l'l',"""'|',','r'r'r'l"rn'n'T'l"l'1""l" 1900 1850 1800 1750 1700 1650 1600 1550 1500 1450 1400 1350 1300 1250 1200 1150 1100 1050 1000 950 900 850

.......................................................................................................

..W..a.y.e..n..u..m..b..er.s...(.c .m.-1) ...................................................................................

Fig. 2. FT-IR spectra of (a) trans-2-butene reversibly adsorbed as such on MgFe204; (b) products of the reactive adsorption of trans-2-butene on MgFe204 at r.t.; (c) products of the irreversible adsorption of but-3-en-2-ol (methyl-allyl alcohol) on MgFe204 at r.t..

992

T a b l e I. Position (cm 1) o f the IR bands of 1-butene in the gas, in the liquid and in the a d s o r b e d state on o x y - d e h y d r o g e n a t i o n catalysts and on TiO2 as a reference. Vibr. vl v2 v3 v4 v20 v5 v21 v6

2 v26 v7

Ass.

sym

v~=CH2 v=CHvs=CH2 ranCH3 Va~CH3 vsCH3 v~CH2 vsCH2

a' a' a' a' a" a' a" a'

2 ~5~CH3 2 8~CH3 x 2 wCH2 .

a'

vC=C

.

IR + R gas s-cis ] gauche 3090 3019 3008 2982 2978 29521 2948 2936 2888

. a'

-. . . . . .

2873 2848,2845 1828

2876 2860 1840

1647

1641

1469 1460 1463

1463 1458

1643

--

IR IR IR IR liquid . MgFe204. MgCr204. FeCrO3 . 3080,3070 3075 3074 I 3006 3000 2995 2973 2970 2965 2935 2937 2917 2925 2894,2888 2905 2895

v22 v8

~Sa~CH3 ~Sa~CH3

. a" . a'

v9

~5-CH2-

a'

1450

vl0 vll v12

~'I42= 8sCH3 w-CH2-

a' a' a'

1426 1421 1380 1342 1318

1444

1439 1434 1414 1385,1373 1345 1320 1290

2877

IR TiO2 3075 3000 2995 2975 2971 2945 2925 2904 2884

1870

1875

1636

1907 1835 1634

1632

1462

1462

1462

1442

1441

1440

1417 1378 1346

1418 1377 1313

1418 1375

1625 1610 1467 1462 1455 1449 1443 1418 1378 1315

v13 8-CH= a' 1306 1296 1296 1295 v23 x-CH2a" 1264 1268 1264 v24 rCH3 a" 1177 1162 1173 v14 rCH3 a' 1128 1127 1131 v15 r=CH2 a' 1071 1079 1071 1077 v16 vCC a' 988 1020 1018 1019 v25 wCH=CH a" 999 993 994, 973 1001 v26 w=CH2 a" 915 912 909 v17 vCC a' 836 854 852,836 912 v27 rCH2 a" 784 14 12 Refs. ' ' i3 ' this work ' 'I 15 ' T h e l o w - f r e q u e n c y regions for adsorbed species are non-available due to the a d s o r b a n t skeletal absorptions.

Table 11. Position (cm-') of the IR bands of cis-2-butene (left) and trans-2-butene (right) as such and in the adsorbed state. Ass.

Vibr.

I

I sym. I

~ 2 2 (v=CH-

I &

2949 2932 __ 2890 -

vXH3 2 S,CHs 2 S,CH3 KH~+wCH

v4

2871 2000

comb comb

a,

vc=c

a,

1786 1699 ._

1481 1461 1447

W W

994 1,2

A

1,1-,."

t, l i

i

i~.!

r

/

0,9~

.,"

--

~

"",,.

it""

,;I/

"-

,:

.

,

}-v " '"v""

L,.

~,

\ ......

A ,,..

,Jilt

N.,J

t

J .

! ! i! II

.9

'w"

9 ,,t

:! ~

-.."

x, .

I

[

I I

"

[ "., i i

/~

l~ ..

i

I ~

ii

",

i'

i! :!

-x..,

".-

,,.%_.................................. - ',.A., -,,, . . . . . .

o.s~

O,4~:

:~

o.34

o2~ 'i

:.....

o,.,i o,oi

i

i t

|l

C

-w~,

~, ........

i\

s t

_,.j

/

d

r-~--'i'- -'l--l-.1--1-.!-.t-l.-.i.--t-.i.-.i.-t.-.i,.-~'-'.r 1900 1850 1800 1750 1700

.

/

/ ~""

',

\

, 1

t,

,,..

"-'"-J"

,,._J "J

!i

_

,=,~

"11I

Itit,

sty.\ " ,

ff

'''",.

~,4

"r-"r.'l""t"'r"*" !-. v.l-.-,-.l--~.-.t-t--+-,-+-+.-l" ,,'.-r162162 1650 1600 1550 1500 1450 1400

....

,

;'t

~;

"-'---.. . . . . .

"

I-'"

A

.,

)

1[ ~

ilk,

t l ii

i{ \.. j ilt

11 :i II

"

i i

I i ", ! = ",

ii

i!,

]iii 9 V~

/

/ t\

\

X\

T :r-":~':-'-~-."l--T-~'~" " ~ < 7 " - l " g - 1 " 4 " r 4 ~ ":;':e"l-"r--t"--r. "r.T -~--'l--'r.v. 1" "*"I--Y. !-. Y+'+'+'+'-I'" +'"+"r ;" t '~ yi?''t''ll'1350 1300 1250 1200 11 SO 1100 1050 1000 g50 900 850

Wavenumbers

~ c m - 1~)

Fig. 3. FT-IR spectra of (a) surface species arising from cis-2-butene adsorption on MgFe204 at r.t.; (b) liquid cis-2-butene; (c) surface species arising from 1,3-butadiene adsorption on MgFe204; (d) liquid 1,3-butadiene. The bars denote the species arising from reactive adsorption of cis-2-butene, i.e. but-3-en-2-oxides, and but-2-en-1-oxides.

3.2. Interaction

of but-3-en-2-ol

(methyl-ailyl

alcohol)

on MgFe204.

In order to define the structure of thesurface alkoxides produced by the reactive adsorption of the three linear butenes on MgFe204, we investigated also the interaction of the catalyst with several C4 oxygenate compounds. In Fig. 2,c the spectrum of the irreversibly adsorbed species arising from but-3-en-2ol adsorbed at r.t. is shown. It is dominated by the features of the corresponding alcoholate species, but-3-en-2-oxide. Evidence is provided by the almost complete disappearance of the bands associated to vOH and 8OH vibrations. The strongest bands of but-3-en-2-oxides are virtually identical to the strongest bands of the irreversibly adsorbed species arising from the three n-butenes on the same surface. The positions of the absorption bands of but-3-en-2-oxide species formed from adsorption of but-3-en-2-ol on MgFe204 are compared in Table III with those of the spectra of the alkoxide species arising from linear butenes. For comparison, the position of the IR bands of prop-2-en-1-oxide species arising from allyl alcohol (prop-2-en-l-ol) and of alkoxide species arising from propene adsorption on MgFe204 are also reported [17]. It is evident that the alkoxide species formed on Mg-ferrite by interaction of butenes and propenes are the same as those produced by the corresponding allylic alcohols, i.e. they are (methyl)allyl-alcoholate species. It is not excluded that the weak differences in the relative intensities of some bands, and the presence of some additional shoulder in the spectra of the species arising from the olefins, is due to the copresence of the isomeric form but-2-en-1-oxides.

995 Table III. Wavenumber (cm -1) of the IR bands of the CH2=CH-CH2-O- allyloxy- species and CH2=CH-(CH3)CH-O- methyl-allyloxy- species. Ass.

vC=C ~-CH2~5~CH3 ~r-7-/2= ~-CH= asCH3 w-CH2~CH-

Allyl alcohol allyloxy- species on methyl methyl-allylateson MgFe204, from C3HsOH M ,1 cis I gauche CCla from from alcohol - C4HTOH 1-C4I-I8 c-2-C4H8 sol C3HsOH propene C4H7OH t-2-Can8 I 1654 1646 1641 1646 1638 1453 1455 1452 I 1451 1428 1414

1421 1410

1372

1373

j

1321

13391 1314 1290 1230 1199

1348

1370 1320

|

1191

1132 vC-O ;1038

1114 1110

tCH2= ' 995 wCH2= ! 919 I 930 Refs. 18

1120 1026' 992 921

|

1420 1385 1368 1317

1325 1252

J

o

x-CH2vC-C / rCH2=

1423 1410

i

i

~COH

1420

1287 1247

1100 1050

1140 1045

1152

1230 1175 1154

1102 1064

1115 1074

991 i 1009 922 17 and this work

1189 1153 1130, 1115 1089 1056 1004

1188 1154 1130, 1115 1089 1058 1000

The spectra of the alkoxide species produced by adsorbing other alcohols, in particular of 2butoxide species formed by 2-butanol, strongly differ from those arising from the C4 olefins. This excludes that alkoxides are formed by electrophilic addition of acidic OH's on the C=C double bond (like found on other surfaces [ 12]) and supports the conclusion that such species are formed by oxidation of the olefins allylic position. In Fig. 2,c bands at 1700 and 1230 cm -1, which are due to C=O stretching and C-C-C asymmetric streching of a ketone, are also found. In effect, these bands correspond to those of adsorbed methyl-vinyl ketone, that can be easily produced by further oxidative dehydrogenation of but-3-en-2oxides [15]. Further heating causes the progressive growth of bands assignable to carboxylates, that are closely similar to those observed upon heating of the species arising from adsorption of n-butenes. Again, they are apparently mostly composed by bands of formates, acetates and acrylates.

3.3. Adsorption and oxidation of 1,3-butadiene on MgFe204. The spectra obtained upon adsorption of 1,3-butadiene over the Mg-ferrite catalyst are reported in Fig. 3. The spectrum obtained is similar to that previously reported for the adsorption of 1,3-butadiene over ot-Fe203 [19]. The detection of the symmetric C=C stretching mode at 1636 cm -1 (IR inactive in the gas) shows that the molecule is either in the noncentrosymmetric s-cis configuration, or its symmetry is strongly lowered by interaction with

996

the surface. On the other hand, the shifts undergone by the most intense and sensitive bands do not bypass 3 cm -1, indicating that vibrational perturbation is weak. Outgassing at r.t. causes the complete disappearance of 1,3-butadiene adsorbed as such, but weak bands due to species produced by reactive adsorption still are at the surface. By increasing temperature, weak bands become detectable that can be associated to carboxylate species. 4. D I S C U S S I O N The data reported above show that well characterized molecular adsorbed species of the three n-butene isomers are formed on the surface of MgFe204 n-butene oxy-dehydrogenation catalyst. Their vibrational perturbation indicates that a n-bonding occurs between the olefinic C=C double bond and Fe 3+ surface cationic centers. The results described above show that methyl-allyl alkoxides are also formed. Such species can also be produced by adsorption of but-3-en-2-01 (methyl-allyl alcohol) and can easily decompose to give butadiene gas. Our previous studies concerning hydrocarbons oxidation over transition metal oxides allowed us to conclude that most hydrocarbon molecules are activated at their weakest C-H bond by oxidation occurring at the expense of high-oxidation-state transition metal cations exposed at the catalyst surface, giving rise to surface alkoxides [8]. It is likely that this also occurs on the surface Fe 3+ cations on Mg-ferrite surfaces, with a mechanism that can be described as in the following scheme. H

!

#,

H2C~--CH'

0 2-

.... C .... ' C H 3

-

H

I H2C--CH

..... C .... ' C H 3

i

0 2-

O-

_

Fe 3+

F-e 3+

Fe 2 +

OH Fe 2 +

It seems reasonable to propose that the primary interaction inducing the proposed C-H activation mechanism consists in the interaction of the C-H bond with fenic ions, which are the oxidizing agents, as already proposed for the oxidation of different hydrocarbons at the surface chromate species of oxidized MgCr204 [8]. This mechanism implies that the two electrons of the C-H bond are assumed by the catalyst surface, where two ferric ions can become reduced to ferrous ions. But-3-en-2-oxide decomposes easily by elimination (as most alkoxides do) to the corresponding olefin, 1,3-butadiene in this case, and a surface OH group. Two OH's condense to allow water to desorb and oxygen reoxidizes the reduced iron cations. This is a very reasonable catalytic cycle for 1,3-butadiene production. The behavior of butenes on MgFe204 is in line with that previously observed in the case of propene adsorption and oxidation, that goes through allyl-oxy species [17]. We previously [8,14] tentatively assigned to (methyl-)allyloxy- species bands observed after adsorption of both propene and n-butenes on MgCr204 (which is also an active catalyst for butene oxydehydrogenation [5]). On another good catalyst for this reaction, the mixed oxide FeCrO3, we observed methyl-vinyl ketone (but-3-en-2-one) to be formed from linear butenes, i.e. a species that can be produced by further oxidative dehydrogenation of but-3-en-2-oxide species [15]. This shows that the surface chemistry of this family of oxy-dehydrogenation catalysts works with the same or closely similar mechanisms.

997 The predominant production of but-3-en-2-oxides, likely together with variable but minoritary amounts of but-2-en-l-oxides (as deduced by of the presence of additional shoulders in the spectra we have assigned to but-3-en-2-oxides), from all three n-butene isomers imply that these two species should interconvert. This could occur through a dissociated form which has the character of a methyl-allyl carbenium ion. Previous studies demonstrated that allyl species intermediates in propene oxidation to acrolein must be "symmetric" before bonding irreversibly with oxygen atoms [6]. Nevertheless, according to recent theoretical chemistry investigations [20] and to our previous proposal [21~, alkoxide species have a more or less pronounced carbocationic nature, that becomes predominant at higher temperatures. In other (and more precise) words, alkoxides are the adsorbed intermediates while carbenium ions are the "U'ansition states" they produce [21]. It seems consequently very reasonable to propose that allyl-alcoholate species we observe at room temperature by olefin allylic oxidation can react as carbenium ions ionically interacting with oxide anions. In the case of the species arising from propene, the resulting carbenium ionic species are symmetric, and the bond with oxygen can indifferently occur at 1 or at 3 position. This picture strongly supports the idea that allyls play the role of symmetric cationic transition states (allyl carbenium ions) more than anionic intermediates in nature in allylic oxidation. H H I

._

H 2 C : C H ' ....C .... ~I

I

-"

A

H

/C~

"c, + ,c

O-

H

O=

.H

H

H T

H

H 2 C : C H .... C ....CH3 O-

--,

/C

"C +~C"

CH 3

c/CH=CH-CH3

H O- H

O

An evolution of alkoxides alternative to elimination, consists in a second oxidative dehydrogenation step to give the corresponding carbonyl compounds. This is what necessarly occurs with the C3 allyl alkoxide in the way producing acrolein from propene, because there are not hydrogens available for the elimination reaction [6]. We have observed spectroscopically this reaction at the surface of quite unselective catalysts like MgCr204 both from propene and butenes [ 17], and in the case of FeCrO3 for butenes [15]. It is consequently evident that two competitive ways can occur, in our case: H2C=CH-CH=CH 2 -OH

T H2CzCH'

d '+

i n+ M

~+

.... C .... 'CH 3

\

O-

n+ M

"~ H2C---CH ~C /

CH3

11 O M

(n- 1 )+

+2 H + M

(n_ 1 )+

998 On our catalysts, the oxy-dehydrogenation of butenes to butadiene via the above elimination way is def'mitely faster than the evolution via the dehydrogenation way. The methyl-vinyl ketone, formed from but-3-en-2-oxides is further easily oxidized because of its ability to give enolization [20], and finally gives rise to total combustion to carbon oxides. Carboxylate species are likely intermediates in this way. The production of 1,3-butadiene needs catalysts that cannot adsorb butadiene strongly. Actually, this occurs on Mg ferrite and distinguishes the behavior of this material from that of vanadia-based catalysts, that allow the oxidation of butenes to maleic anhydride and actually adsorb butadiene strongly [12]. Our data suggest that the predominant pathway to total oxidation is competitive with respect to the key alkoxide elemination step, more than successive to it, over MgFe204. REFERENCES

1. 2. 3. 4. 5. 6. 7.

K. Weissermel and H.J. Arpe, Industrial Organic Chemistry, VCH, Weinheim, 1993 J. Schulze and M. Homann, C4-hydrocarbons and derivatives, Springer, Berlin, 1989. R.J. Rennard and W.L. Kehl, J. Catal. 21,282 (1971) M.A. Gibson and J.W. Hightower, J. Catal. 41 (1976) 420. H.H. Kung and M. C. Kung, Advan Catal. Relat. Subj. 33 (1985) 159. R.K. Grasselli and J.D. Burrington, Advan. Catal. 30, 133 (1981) E. Finocchio, G. Busca, V. Lorenzelli and R.J. Willey, J. Chem. Soc. Faraday Trans., 90, 3347 (1994) 8. E. Finocchio, G. Busca, V. Lorenzelli and R.J. Willey, J. Catal., 151,204 (1995) 9. E. Finocchio, G. Busca, V. Lorenzelli and V. Sanchez Escribano, J. Chem. Soc. Faraday Trans., 92, 1587 (1996) 10. G. Busca, F. Cavani, E. Etienne, E. Finocchio, A. Galli, G. Selleri and F. Trifirb, J. Mol. Catal., A, Chemical, 114 (1996) 343 11. G. Busca, M. Daturi, E. Kotur, G. 0lived and R.J. WiUey, in Preparation of Catalysts VI, G. Poncelet et al. eds., Elsevier, Amsterdam, 1995, p. 667-676. 12. G. Busca, V. Lorenzelli, G. Ramis and V. S. Escribano, Mat. Chem.Phys.29 (1991) 175. 13. J.R. Durig and D.A.C. Compton, J. Phys. Chem. 84 (1980) 773. 14. E. Finocchio, G. Ramis, G. Busca, V. Lorenzelli and R.J. Willey, Catal.Today, 28 (1996) 381. 15. G. Busca and V. Lorenzelli, J. Chem.Soc., Faraday Trans., 88 (1992) 2783 16. D.C. McKean, M.W. Mackenzie, A.R. Morrisson, J.C. Lavalley, A. Janin, V. Fawcett and H.G.M. Edwards, Spectrochim. Acta 41A (1985) 435. 17. G. Busca, E. Finocchio, V. Lorenzelli, M. Trombetta and S. A. Rossini, J. Chem.Soc., Faraday Trans., 92 (1996) 4687. 18. B. Silvi and J.P. Perchard, Spectrochim. Acta 32 A, 11 (1976) 19. G. Busca, L. Marchetti, T. Zerlia, A. Girelli, M. Sorlino and V. Lorenzelli, Proc. 8th Int. Congr. Catal., Verlag Chemie, Weinheim, 1984, Vol. III, p. 299-310. 20. V.B. Kazansky, M.V. Frash and R.A. vanSanten, l lth Int. Congr. Catal., Elsevier, Amsterdam, 1996, p. 1233. 21. G. Busca, E. Finocchio, G. Ramis and G. Ricchiardi, Catal. Today, 32 (1996) 133.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

999

CYCLOHEXENE OXIDATION CATALYZED BY TITANIUM MODIFIED HEXAGONAL Y TYPE ZEOLITES Kenneth J. Balkus, Jr.*, Alia K. Khanmamedova and Jimin Shi

Department of Chemistry, University of Texas at Dallas, Richardson, TX 75083-0688, United States

SUMMARY Hexagonal NaY (Si/Al=5.1) and the dealuminated hexagonal NaY (Si/Al=63) type zeolites were modified with titanium by reacting sin-face silanol groups with titanocene chloride. After calcination to remove organics the resulting titanium containing hexagonal Y zeolites were found to be effective catalysts for the hydrogen peroxide based oxidation of cyclohexene. Preliminary results for XRD, UV-Vis and FT-IR characterization of the zeolite catalysts are presented. 1. INTRODUCTION The success of the titanium silicate molecular sieve TS-1 [1] as a commercial oxidation catalyst [2-4] has prompted a flurry of activity in the area of titanium zeolites. Titanosilicates and other transitional metal modified zeolites continue to show promise in many areas of oxidation catalysis [5]. Much effort has been expended on the characterization of TS-1 in order to relate the catalytic properties with the presence of isolated titanium ions incorporated into the zeolite framework [6,7]. This has certainly encouraged the synthesis of other titanosilicate molecular sieves such as the large pore Ti-UTD-1 [8,9], Ti-beta [10] and Ti-MCM-41 [11] for example. Although, titanium may be incorporated into the oxide framework of many zeolites by direct synthesis, a simpler approach involves the post synthesis reaction of surface silanols with reactive titanium organometallics. This strategy is exemplified by the work of Maschmeyer et al [12] which involved the modification of mesoporous MCM-41 with titanocene dichloride followed by thermal decomposition to yield an active oxidation catalyst. Related studies of molecular sieve surface modification with reactive metal species also appear promising [ 13-17]. We have employed this method in the modification of dealuminated hexagonal Y zeolites with titanium The structure of hexagonal Y has the EMT topology and is characterized by a three dimensional large pore (12 membered ring, 7.4 x 6.5A.) system which would have a size advantage over the medium pore TS-1 catalyst. Upon dealumination by acid treatment the structure of hexagonal Y is retained but the presence of defect sites is evident. We have reacted titanocene dichloride with the template free hexagonal NaY as well as the dealuminated zeolite. After template removal a si~ificant amount of titanium is retained. Subsequent, calcination heals the defect sites and

1000 produces a titanium aluminosilicate that is an effective epoxidation catalyst. Preliminary results for the characterization of titanium modified hexagonal NaY as well as results for the oxidation of cyclohexene using H202 are presented below. 2. EXPERIMENTAL Hexagonal NaY (Si/Al=5.1) was prepared and characterized as previously described [18]. Samples 1-4 were calcined 650~ to remove the crown ether template but not dealuminated while sample 5 was stirred at pH 1 (HC1) for 4 hours and then calcined at 650~ Titanocene dichloride, CpETiC12 (Aldrich) was used as the titanium source for modifying the surface of hexagonal NaY following a procedure similar to that previously described [ 12]. In a typical reaction 0.5 g of calcined hexagonal NaY was stirred in 20 mL of chloroform at room temperature followed by the addition of 0.2 g of titanocene dichloride dissolved in an additional 20 ml of chloroform The mixture was stirred for 0.5 hours and then 5 g of triethylamine were added with stirring for two more hours to scavenge HC1. During that period the color of the suspension changed from red to yellow. The resulting zeolite was suction filtered, washed with chloroform and dried at 80~ The organometallic modified hexagonal NaY was then calcined at 650~ in flowing dry oxygen for 4 hours. FT-IR spectra were recorded as KBr pellets using a Mattson 2025 FT-IR spectrophotometer. XRD patterns were collected using a Scintag XDS 2000 diffractometer. CaF2 was used as an internal standard. UV-Vis spectra were collected from samples prepared as nujol mulls between quartz plates using an Hitachi U-2000 IYV-Vis spectrophotometer. Elemental analyses were performed by Galbraith Laboratories, Knoxville, TN. The catalytic oxidations of cyclohexene were carded out as batch reactions in sealed glass vials (15 mL) under nitrogen at -~55~ In a typical experiment the reactor was loaded with 0.1 gram of catalyst, 8 mmol of cyclohexene, 2.5 mmol of H202 (30 wt.%) and 4 mL of acetonitrile solvent. All reactions were sampled by syringe through a rubber septum~ Products were analyzed by gas chromatography using an HP 5840 A capillary GC equipped with a 15 m AT-1 capillary column and a flame ionization detector. Products were verified by known standards. The data in Table 2 represent an average of multiple analyses in order to verfy the significance of the changes. 3. RESULTS AND DISCUSSION 3.1. Modification with Titanium

The organometallic Cp2TiC12 is a reactive 16 electron complex that is readily hydrolyzed. Reaction with silanols is expected to result in Ti-O linkages with concomitant production of HC1. This type of complexation has already been shown to occur with the silanols ofMCM-41 resulting in immobilized titanocene species [ 12]. One might also envision the partial exchange of sodium ions for Cp2TiCI~2n§ ions but this is probably not a si~ificant process in chloroform. In the present study Cp2TiC12 was reacted with calcined hexagonal Y and a dealuminated sample. Table 1 shows the results for elemental analysis of calcined hexagonal NaY treated with different amounts of titanocene dichloride (Samples 1-4) and the dealuminated titanium zeolite (sample 5). For the calcined only samples the amount of

1001 titanium incorporated appears to track with the amount of organometallic used. One might expect that dealumination would generate defect sites or silanol nests that would increase the Table 1. Results for elemental analysis of hexagonal NaY samples after calcination. Samples

Cp2TiC12

Titanium

Si/Al

(grams)

(weight %)

1

0.020

0.68

5.1

2

0.033

1.33

5.1

3

0.044

1.97

5.1

4

0.200

7.22

5.1

5

0.200

11.44

63.3

amount of organometallic that would react with the surface. Additionally, the more siliceous zeolite should become more hydrophobic which is a desirable feature titanosilicate catalysts. Table 1 shows that when hexagonal NaY is dealuminated more titanium is incorporated. For sample 5 the same concentration of Cp2TiC12 as in sample 4 was reacted but there was a -~60% increase in the resulting titanium loading. This is consistent with the greater number of silanols in sample 5. Additionally, if the titanium modification in samples 1-4 was predominately simple impregnation or deposition of titania in the channels, then one would not expect such a difference between samples 4 and 5. Interestingly the amount of titanium incorporated in sample 5 is almost an equimolar amount in relation to the quantity of aluminum removed. It should also be noted that HC1 treatment of hexagonal NaY did not result in any si~ificant loss in crystaninity as evidenced by XRD results and scanning electron microscopy. The high titanium loadings achieved with samples 4 and 5 seems to validate this approach to modifying zeolite uaTaces with reactive metal species. However, from a catalysis perspective one would prefer to have isolated titanium centers which will be difficult at such high loadings. One might expect formation of Ti-O-Ti oligomers or bulk TiO2 at such loadings. UV-VIS spectral analysis reveals that samples 4 and 5 have absorption edges above 325 nm indicating the possible presence of anatase [ 19]. Therefore, the nature of the surface titanium in these samples will have to be considered when evaluating catalytic reactivity. Infrared spectroscopy has proven to be a valuable tool in characterization of titanosilicates. In the case of TS-1 and other Ti-substituted molecular sieves the infamous band at -~960 cm-1 was held to be evidence of titanium incorporation into the framework [20,21]. However, the same band was found with titanium free silicalite [22] and not observed when silicalite was modified from aqueous fluorotitanate solutions [23]. Most would agree now that this band is an Si-O stretch derived from defect sites in the silicate yet there may be some overlap and contribution from Si-O-Ti stretches. It seems in some cases

1002 the intensity of the band at 960 cm1 correlates with the amount of titanium but only to the point where titanium oligomers begin to form [ 17].

(f) (e) (d)

o t..)

(c) (b)

(a) 2000 1500 1000 Wavenumbers, cm-1

500

Figure 1. FT-IR spectra of (a) the as synthesized hexagonal NaY, (b) calcined hexagonal NaY, (c) hexagonal NaY after HC1 treatment, (d) hexagonal NaY after HC1 treatment and calcination, (e) dealuminated hexagonal NaY after reaction with Cp2TiC12 and (f) dealuminated hexagonal NaY after reaction with Cp2TiC12 followed by calcination. The FT-IR spectra in figure 1 reveal that the as-synthesized hexagonal NaY has a broad shoulder between 950 and 850 cm-1. However, after calcination and removal of the crown ether template this region of the spectrum is clear and the bands associated with the asymmetric stretches become better defined. After HC1 treatment there is clear evidence of dealumination as shown in figure l c. The band assigned to the asymmetric Si-O stretching mode associated with internal tetrahedra shifts from 1032 to 1092 c m "1 w h i c h is consistent with an order of magnitude increase in Si/A1. The weaker bands associated with internal symmetric stretches and double ring vibrations also decrease in intensity, however, the zeolite structure is preserved according to XRD results. Dealumination also results in formation of a band at 954 cm1 which is relatively sharp compared to partially amorphous zeolites. Interestingly, this band at 954 cm1 disappears upon calcination with very little change in the rest of the spectrum (figure l d). In most cases, the zeolites that show such bands do not exhibit this type of behavior upon calcination but rather the band intensifies. The loss of the

1003 954cm-1 band suggests some healing process associated with the defect sites that does not involve reinsertion of T atoms since none of the other bands change. One might expect these defect sites to be reactive towards Cp2TiC12 which might affect this band. Figure l e indicates that modification with the organometallic reduces the intensity and slightly shifts the band to --~947cm"1. Any IR bands that might be associated with the Cp rings are not evident, possibly because they are masked by the zeolite bands or dissociated [12]. Nevertheless, we know from the results in Table 1 that a si~ificant amount of titanittm has been included at this stage. Calcination of this sample (figure If) results in loss of the 947cm-1 band and no shifts in the other bands. The reaction of Cp2TiC12 with the calcined hexagonal NaY in figure lb does not generate a 947cm 1 band nor is there any other si~ificant change in the resulting spectra (not shown). One would have to conclude that in all these cases the titanium species are being grafted to the surface of the zeolite pores. There is no evidence that would suggest that titanium is substituted for T atoms in the framework. 3.2 Oxidation of Cyclohexene The oxidation of cyclohexene using hydrogen peroxide was chosen as a test reaction for the catalytic evaluation of the titanium modified hexagonal NaY samples. Scheme I illustrates some of the typical products of cyclohexene oxidation. The epoxide and the diol which is a hydrolysis product of the epoxide, generally reflect a concerted process. In contrast the allylic alcohol and ketone are often ascribed to an autoxidation or radical process. We anticipated that some homolytic decomposition of the peroxide may be observed with these acidic zeolites. In fact, there was-~74% conversion of H202 over calcined hexagonal NaY after heating at 55~ for 24 hours. This resulted in only a 1% conversion of

[•0

+ H20 2 Cyclohexene

~Cyclohexene oxide

2-Cyclohexene-l-ol

,.

OH trans..1,2-Cyclohexanediol

7-Oxabicyclo[4.1.0]heptan-2-oi

__. d>O 2-Cyclohexene-l-one

7-Oxabicyclo[4.1.0]heptan-2-one

Scheme 1

1004 cyclohexene during this period to a 3:1 mixture of 2-cyclohexen-1-one and 2-cyclohexen-1ol. No epoxide was formed under these conditions. Table 2 shows the results for the oxidation of cyclohexene catalyzed by a series of titanium modified hexagonal NaY zeolites. The epoxide is produced in all cases but under the present conditions is transformed to the diol as the major product. Additionally, there are smaller amounts of autoxidation products (K + A). This product distribution was to be expected given the acidity of these zeolites and has certainly been noted before with other titanium modified zeolites [24,25]. After aluminum removal from the zeolites, the side reactions and selectivity improves [26]. Most of the catalyst activity occurs in the first few hours since after one day there is only a small increase in the conversion of substrate or peroxide. There is an apparent improvement in selectivity for the epoxide after a day but the rate of reaction is quite low. This would suggest clogging of the zeolite pores which retards the reaction chemistry. The spent catalyst is typically pale yellow in color while the starting zeolite is white. The MCM-41 modified with titanium in a similar fashion to our study also turns yellow in a peroxide based cyclohexene oxidation but deactivates after only 90 minutes [ 12]. The used catalyst may be calcined at 500~ in flowing oxygen to remove the color and restore the original catalytic activity. Controlling tile catalyst acidity as well as evaluating the influence of solvent will be important factors to consider in improving catalyst lifetime. Table 2. Results for cyclohexene oxidation catalyzed by titanium modified hexagonal NaY.

% Conversion Sample

1 2 3 4 5

% Selectivitya'b

Hours

CY

H202

CYO

diol

K

A

% Efficiencyc

5 24 5 24 5 24 5 24 5 24

5.7 6.4 8.2 10.5 10.8 12.8 7.2 9.5 4.5 6.4

50 54 69 75 56 65 36 50 30 47

5.0 4.8 6.1 5.5 5.7 4.3 4.3 4.0 6.3 5.7

58.1 60.0 57.9 50.2 69.3 53.1 59.1 68.1 52.5 51.8

24.2 22.9 31.1 31.0 15.4 25.8 27.5 19.9 32.0 31.9

12.6 12.2 16.8 13.2 9.3 1.6.7 8.9 7.8 9.2 10.6

36 37 38 45 59 61 64 62 54 53

CY = cyclohexene; CYO = cyclohexene oxide; diol = 1,2-cyclohexanediol; K - 2-cyclohexen- 1-one; A = 2-cyclohexen- 1-ol. b Selectivity = (mmol product / mmol total products) x 100. c Efficiency = (retool of converted cyclohexene /mmol of converted H202) x 100 a

1005 Table 2 indicates that the conversion of cyclohexene as well as the peroxide efficiency increases as the titanium content increases for the first three samples. Recall samples 4 and 5 showed evidence of bulk TiO2 occlusion which must have adverse effects on the reaction including pore blockage. Sample 5 is the dealuminated zeolite which should have fewer but stronger acid sites. In spite of having the highest titanium loading this sample exhibits the lowest conversion of both substrate and peroxide. The peroxide efficiency is comparable to the better catalysts which may reflect the lower acid site density. A slightly lower level of epoxide hydrolysis may also represent a more hydrophobic environment at the active site. These result certainly warrant further investigation of the high silica hexagonal Y type zeolites as support materials for isolated titanium species. 4. CONCLUSIONS We have grafted titanocene derived species onto the surface of hexagonal NaY before and after dealumination resulting in effective epoxidation catalysts. The organometallic was intended to reduce the possibility of forming oligomers and bulk titania but it is clear that at high enough loadings this is unavoidable. However, this method of zeolite surface modification has resulted in epoxidation activity one would associate with isolated titanium centers which then represents a viable alternative to framework titanosilicate catalysts. Clearly more work is needed to better define the optimum zeolite composition and reaction conditions that will stabilize the system and maximize selectivity. ACKNOWLEDGMENTS The support of the Robert A. Welch Foundation and the National Science Foundation are gratefully acknowledged. REFERENCES

.

4. 5. .

7. 8.

M. Taramasso, G. Perego and B. Notari, Preparation of porous crystalline synthetic material comprised of silicon and titanium oxides, U. S. Patent No. 4,410,501 (1983). M. Taramasso, G. Manara, V. Fattore and B. Notari, Silica based synthetic material containing titanium in the crystal lattice and process for its preparation, U.S. Patent No. 4,666,692 (1987). B. Notari, Catal. Today, 18 (1993) 163. B. Notari, Stud. Surf. Sci. Catal., 37 (1987)413. D.E. De Vos, P.L. Buskens. D.L. Vanoppen, P.P. Knops-Gerrits and P.A. Jacobs in Comprehensive Supramolecular Chemistry 7 (1996) 647 and references there in. 1L Millini and G. Perego, Gaz. Chim Ital., 126 (1996) 133. B. Notari, Adv. Catal., 41 (1996) 253. K~J. Balkus, Jr., A.G. Gabrielov and S.I. Zones, Stud. Surs Sci. Catal., 97 (1995) 519. I~J. Balkus, Jr., A.A. Khanmamedova, A. Gabrielov and S.I. Zones, Stud. Surf Sci. Catal., 101 (1996) 1341.

1006 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26.

A. Corma, M.T. Navarro and J. Perez-Pariente, J. Chem. Soc., Chem. Commun., (1994) 147. P.T.Tanev, M. Chibwe and T.J. Pinnavaia, Nature, 368 (1994) 321. T. Maschmeyer, F. Rey, G. Sancar and J.M. Thomas, Nature, 378 (1995) 159. P.G. Pries de Oliveira, J.G. Eon and J.C. Volta, J. Catal., 137 (1992) 257. A. Corma, A. Fuerte, M. Iglesias and F. Sanchez, J. Mol. Catal., 107 (1996) 225. N. Ichikuni, M. Shirai and Y. Iwasawa, Catalysis Today, 28 (1996) 49. IL Butch, N. Cruise, D. Gleeson and S.C. Tsang, Chem Commtm., (1996) 951. S. Schwartz, D.1L Corbin and A.J. Vega, Mater. Res. Soc. Symp. Proc., 431 (1996) 137. M.W. Anderson, K.S. Pachis, F. Prebin, S.W. Cart, O. Terasaki, T. Ohsuna and V. Alfredsson, J. Chem Soc., Chem_ Commun., (1991) 1660. J. Klaas, IC Kulawik, G. Schulz-Ekloff and N.I. Jaeger, Stud. Surf. Sci. Catal., 84 (1994) 2261. M.1L Boccuti, I~M. Rao, A. Zecchina, G. Leofami and G. Petrini, Stud. Surf. Sci. Catal., 48 (1989) 133. J.S. Reddy, 1L Kumar and P. Ratnasamy, Appl. Catal., 58 (1990) L1. A. Zecchina, G. Spoto, S. Bordiga, A. Ferrero, G. Petrink G. Leofanti and M. Padovan, Stud. Surf. Sci. Catal., 69 (1991) 251. G.W. Skeels and E.M. Flanigen, ACS Symp. Set., 398 (1989) 420. C.B. Dam and M.E. Davis, Appl. Catal., 143 (1996) 53. A Corma, P. Esteve, A. Martinez and S. Valencia, J. Catal., 152 (1995) 18. M.A. Camblor, M. Constantini, A Corma,L. Gilbert, P. Esteve, A. Martinez and S. Valencia, Chem Commun. (1996) 1339.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1007

OXIDATIONS CATALYZED BY ZEOLITE Ti-UTD-1 Kenneth J. Balkus, Jr.* and Alia K. Khanmamedova

University of Texas at Dallas, Department of Chemistry, Richardson, TX 75083-0688, United States

SUMMARY Titanosilicates have been synthesized which have the large pore structure of UTD-1. These molecular sieves were prepared using the metal complex Cp*2Co§ (bis(pentamethylcyclopentadienyl)cobalt(lll) ion) as the template. Ti-UTD-1 has been found to be an effective catalyst for the oxidation of alkanes, alkenes and phenols using hydrogen peroxide as well as the larger oxidant t-butylhydroperoxide. The channel structure defined by 14 membered rings in Ti-UTD-1 also allows the conversion of larger substrates such as 2, 6- di-t-butylphenol. 1. INTRODUCTION UTD-1 (University of Texas at Dallas) is a novel large pore high silica zeolite [1-4] with many promising applications in the area of catalysis. The structure ofUTD- 1 is based on one dimensional channels nmning in parallel [5]. The elliptical shaped channels are defined by 14 tetrahedrally coordinated silicon atoms with pore opening of 10 x 7.5 A as shown in Figure 1. Silicon in the UTD-1 framework may be replaced (using direct synthesis) by many

Figure 1. View ofUTD-1 along the [010] direction transition metals, including titanium [3,6,7]. The Ti-UTD-1 preparation involves the novel use of a metal complex template such as bis(pentamethylcyclopentadienyl) cobak(m) hydroxide as well as titanium ethoxide as the source of titanium As much as 3.5% titanium by weight has been incorporated into Ti-UTD-1 during synthesis. It was anticipated that Ti-

1008 UTD-1 may have catalytic properties such as those observed with the commercially sucsessful titanium silicalite TS-1 catalyst which is effective for alkane and alkene oxidation as well as phenol hydroxylation in the presence of hydrogen peroxide [8]. The large pore nature of Ti-UTD-1 should allow the reaction of large substrates such as 2,6-di-tertbutylphenol as well as the use of oxidants such as tert-butylhydroperoxide (t-BHP) which are too large for the medium pore TS-1 zeolite. Ti-UTD-1 offers an opportunity to examine reactivity in pore space greater than Ti-beta but less than the mesoporous Ti-MCM-41 type molecular sieves. In the present study results for the peroxide based oxidation of cyclohexane, cyclohexene and 2, 6- di-tert-butylphenol will be presented. 2. EXPERIMENTAL Ti-UTD-1 (3.5 % Ti by weight) was prepared and characterized as previously described [3]. The catalytic oxidations were carried out as batch reactions in sealed glass vials (15 ml) at 25, 50 and 60~ In a typical reaction the reactor was charged with 0.1 grams of Ti-UTD-1 catalyst, 6 mmol of substrate, 3 mmol (or 12 mmol) of 90% aqueous tert-butyl hydroperoxide (t-BHP) and 20 mmol (or 27 retool) of acetone or tert-butanol as a solvent. In the case of cyclohexene oxidation using hydrogen peroxide (H202), 8 mmol of olefin and 2 mmol of H202 in t-butanol as well as acetonitrile solvent (80 mmol) were used. All reactions were sampled by syringe through a rubber septum and analyzed by gas chromatography using an HP 5840A capillary GC equipped with a 15 m AT-1 capillary column and a flame ionization detector. Products were verified by known standards and/or GC-MS. 3. RESULTS AND DISCUSSION

3.1 Cyclohexane oxidation The oxidation reactions involving hydrocarbons are frequently important components in many industrial processes yet can be among the most difficult of reactions to accomplish selectively. Cyclohexane serves as a good model substrate but itself is commercially important as a precursor to adipic acid. We have previously reported that Ti-UTD-1 is effective catalyst for the oxidation of cyclohexane using t-butylhydroperoxide [7]. H202 which can donate 47% of its oxygen in comparison with only 17.8% for t-BI-IP [9], is the preferred oxidant. However, the results for cyclohexane oxidation using H202/Ti-UTD-1 are disappointing. Although, Ti-UTD-1 should provide a hydrophobic environment at the active sites, the large pore nature of this material may allow more access to water. Even TS-1 exhibits very low activity for cyclohexane conversion using H202 as the oxidant [10]. The tBI-IP oxidations are more efficient but produce a range of products. Scheme 1 illustrates the oxidation products generally observed in cyclohexane oxidation, where the preferred product is adipic acid. The major product in the case of Ti-UTD-1 catalyzed reactions is generally cyclohexanone with smaller amounts of the alcohol and adipic acid. It should be noted that decomposition of the cobalt complex used as the template generates nanoclusters of a cobalt oxide on the surface of the zeolite [4]. Cobalt oxide and cobalt complexes are effective as peroxide decomposition catalysts [11,12] and should be removed from the :~,eolite to better evaluate the role of the titanium sites.

1009 -~

I V

copt Adipic acid

\

o _ , . Q"-

CO2H Glutaric acid CO2H

~~~C02H 0

Succinic acid

Scheme 1 Table 1 shows some representative results for cyclohexane oxidation with t-BHP over calcined Ti-UTD-I(Co) and Ti-UTD-1 (acid treated to remove cobalt [4]). In acetone the cyclohexane is converted to cyclohexanone as the major product with trace amounts of adipic acid as well as the other carboxylic acids and diketones in Scheme 1. Changing the solvent from acetone to t-butanol improves the conversion and the yield of adipic acid, however, the peroxide efficiency decreases. Increasing the amount of peroxide in the system also improves conversion but the peroxide efficiency is decreased. The product distribution observed with the Ti-UTD-l(Co) samples suggets a homolytic process catalyzed by cobalt. Removal of the cobalt lowers the activity in t-butanol but improves the peroxide efficiency. Table 1. Results for cyclohexane oxidation using t-BHP at 60~ Catalyst

Hours

% Conversion

% Yield

% Efficiencyr

K e A e AA ~ 24 10.8 7.2 3.5 0 36 24 22.2 11.3 4.0 1.7 20 24 29 12.5 3.5 0.4 17 24 13.8 1.5 4.3 trace 34 50 20.5 7.2 3.7 1.2 32 a. acetone solvent, b. t-butanol solvent; c. cyclohexane:t-BHP = 2:1, d. cyclohexane:t-BHP : 1:2; e. K = cyclohexanone, A = cyclohexanol and AA = adipic acid; f. efficiency = (mmol cyclohexane converted/retool t-BHP converted) x 100

Ti-UTD- 1 (Co)a'c Ti-UTD- l(Co) b'c Ti-UTD- I(Co) a'd Ti-UTD- 1b,c

1010 3.2 Cyclohexene oxidation

Oletins are generally much easier to oxidize than alkanes but again the selectivity can be a challenge. Cyclohexene was employed as a model substrate with the preferred product being the corresponding epoxide. Table 2 shows the results for cyclohexene oxidation catalyzed by Ti-UTD-1 (cobalt free) using H202 in t-butanol solvent. At 50~ a reasonable conversion of cyclohexene was observed with reasonable selectivity for the epoxide. Even though the substrate/peroxide ratio was 4:1 there was a fairly high efficiency for hydrogen peroxide utilization. The diol product is the result of epoxide hydrolysis while the allylic products (K + A) reflect a competing homolytic process. In acetonitrile solvent the Table 2 Results for cyclohexene oxidation using H202 in t-butanol solvent at 50~ Hours

% Selectivitya'b

% Conversion

% Efficiency

CYO A K diol 1 1.5 33.3 78 11.1 22.2 33.3 3 3.4 15.0 78 10.0 25.0 50.0 27 11.0 1.5 79 9.4 23.4 65.6 a. CYO = cyclohexene oxide, A = 2-cyclohexen-1-ol, K = 2-cyclohexen-1-one, diol = 1,2cyclohexendiol.; b. Selectivity was calculated as ratio (mmol product/mmol total product) xl00. conversion improved over the same time period in t-butanol. Additionally, the selectivity for cyclohexene oxide improved with very little hydrolysis early in the reaction. The peroxide efficiency improved in acetonitrile as well. These results suggest that there is a strong solvent dependency for this process. Therefore, the oxidation of cyclohexene was run in several solvents and mixtures including acetone and methanol but the best results so far for conversion and selectivity to the epoxide were obtained in acetonitrile. Polar solvents such as water, acetone and methanol have been previously been noted to have an inhibitory effect on the epoxidation reaction in mesoporous titanosilicates [ 11] and Ti-beta [ 13]. In these cases it is proposed that the solvent coordination at the oxide surface blocks the active sites. Table 3 Results for cyclohexene oxidation using H202 ill acetonitrile solvent at 50~ Hours

1.5 3 23

% Conversion

9.5 12 26

% Efficiency

% Selectivity CYO 79.3 70.3 45.1

A 6.9 8.1 12.1

K 13.8 16.2 23.1

diol tr. 5.4 19.7

,

,,

79 84 99

1011 This trend in large pore materials appears to be the opposite from that observed in the medium pore TS-1 catalyst [ 14]. The effect of acetontrile maybe to poison acid sites and/or to allow a greater concentration of substrate near the active site. The later has been proposed to be the dominant factor in Ti-beta [ 13]. The aprotic acetonitrile does not form the extensive hydrogen bonded networks expected for water or the alcohols. However, the acetonitrile is polar enough to allow more substrate into the relatively hydrophobic zeolite. A hydrogen bonded titanium hydrogen peroxo complex has been suggested as an important intermediate where a silanol, titanol or protic solvent stabilizes the structure. In the case of Ti-beta it was proposed that acetonitrile solvent allowed water to function in this capacity where as in other solvents the water would be displaced [13]. As a donor the acetontrile may also bind weakly enough to the titanium to affect the peroxide decomposition. Scheme 2 shows titanol stabilized titanium hydrogen peroxo species in equilibrium interacting with acetonitrile. This would acount for in part why the efficiency is lower and level of autoxidation products is higher in protic solvents. It is interesting that the large pore titanosilicates seem to exhibit similar solvent behavior.

I

H a C - CN:.- -~Ti

/

"o

\ H

o

/n...

\

I H a C - CN:- 9.,..-Ti

O/

+

H+

~O

~0 Scheme 2

A disadvantage of using n202 as the oxidant is low stability with respect to radical decomposition especially in presence of the catalytically active Ti(IV) species. Solvent can certainly influence the homolytic process but we may also add radical traps to supress the formation of allylic products. If one adds hydroquinone to the H202/Ti-UTD-1 system in aeetonitrile then the yield of cyclohexene oxide improves to over 85% while the allylic ketone and alcohol are reduced to trace impurities. The amount of diol formed also decreases but the presence of water in the system is unavoidable and some hydrolysis of the epoxide is expected. Changing the oxidant to t-butylhydroperoxide is expected to improve selectivity. The large pores of Ti-UTD-1 allow the use of t-BHP in contrast to TS-1 where this oxidant is too big. Figure 2 shows the results of cyclohexene oxidation using t-BHP (80 wt.% in ditert-butylperoxide) and t-BHP (90 wt.% in water) after two days of reaction at 60~ where the solvent was acetonitrile. Clearly the presence of water lowers the cyclohexene conversion and the selectivity for epoxide formation.

1012

70

%

60-

50-

~d

40-

30"

20

1

2

3

4

5

6

7

CONVERSION, % Figure 2. Selectivity for cyclohexene oxide versus conversion of cyclohexene for 1 - t-BHP in di-t-butyl peroxide (80%) and 2 - t-BHP in water (90%).

3.3 Oxidation of 2,6-Di-tert-butyl phenol A characteristic feature of Ti-UTD-1 which we would like to exploit is the 14 membered ring system where substrates too large for other zeolites can effectively be oxidized. The large 2,6-di-tert-butyl phenol (2,6-DTBP) has been used as a susbtrate for titanium containing mesoporous molecular sieve catalyzed oxidations [15] to yield the corresponding quinone as shown in Scheme 3 below. Table 4 shows the results for 2,6-DTBP oxidation at 65~ using H202 as the oxidant in acetone solvent. The selectivity to the quinone

-'~OH

+ H20~

0 = = ~

\

N Scheme 3

0

1013 is quite high over both Ti-UTD-1 and the mesoporous titanosilicates with no evidence of the dimer product after 5 hours. The conversion wth the Ti-UTD-1 catalyst is comparable to TiMCM-41 over a similar time frame while the Ti-HMS is somewhat more active. Clearly these results indicate that the large pore Ti-UTD-1 is capable of addressing large substrates. Table 4 Results for the oxidation of 2,6-di-tert-butyl phenol Samples

Pores size (A)

Hours.

% Conversion

Ti-UTD-1

10 x 7.5

Ti-MCM-4 la Ti-HMS a a. Reference 15

- 30 - 30

4.5 40 2 2

20 90 20 83

% Selectivity > 99 95 >98 >95

4. CONCLUSIONS The novel zeolite UTD-1 with titanium in the framework (up to 3.5% by weight) is an effective catalyst for the oxidation of cyclohexane, cyclohexene and 2,6-di-tert-butyl phenol. The catalytic behavior is similar to that of other large pore zeolites and mesoporous molecular sieves modified with titanium which includes solvents effects. Additionally, TiUTD- 1 allows the use of oxidants and substrates too large for the commercial TS- 1 catalyst. We are currently evaluating further the role of solvent and oxidant in an effort to improve selectivity as well as expand the utility of this material in oxidation catalysis. ACKNOWLEDGMENT The financial support of this work by the National Science Foundation, the Robert A.Welch Foundation and Chevron Research and Technology Company is gratefully acknowledged. REFERENCES

.

K. J. Balkus, Jr. and A. G. Gabrielov, The Synthesis of Novel Molecular Sieves using Metal Complexes as Templates, U.S.Patent No. 5,489,424 (1996). K.J. Balkus, Jr., A.G. Gabrielov and N. Sandier, Mater. Res. Soc. Symp. Proc., 368 (1995) 369. K. J. Balkus, Jr., A. G. Gabrielov and S. I. Zones, Stud. Surf. Sci. Catat, 97 (1995) 519. K. J. Balkus, Jr., M. Biscotto and A. G. Gabrielov, Stud. SurE. Sci. Catal., 105 (1997) 415. C.C. Freyhardt, M. Tsapatsis, R.F. Lobo, K.J. Balkus, Jr. and M.E. Davis, Nature, 381 (1996) 295.

1014

,

9. 10. 11. 12. 13. 14. 15.

K. J. Balkus, Jr. and A. G. Gabrielov, The Synthesis of Novel Molecular Sieves using a Metal Complex as Template, U.S.Patent No. 5,603,914 (1997). K, J. Balkus, Jr., A. Khanmamedova, A. G. Gabrielov and S. I. Zones, Stud. Surf. Sci. Catat, 101 (1996) 1341. B. Notari, Adv. Catal. 41 (1996) 253 and references therein. G. Strukul in Catalytic Oxidations with Hydrogen Peroxide as Oxidant, Kluwer, Boston (1992), 6. U. Schuchardt, H.O. Pastore and E.V. Spinace, Stud. Surf. Sci. Catal., 84 (1994) 1877. Z. Liu, G. M. Crumbaugh and 1L J. Davis, J. Catal. 159 (1996) 83. C.B. Roy, J. Catal.,12, (1968), 129. A. Corma, P. Esteve and A. Martinez, J. Catal. 161 (1996) 11. M.G. Clerici, Appl. Catal., 68 (1991) 249. P.T. Tanev, M. Chibwe and T. J. Pinnavaia, Nature, 368 (1994) 321.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1015

Zeolite T i t a n i u m Beta" A selective catalyst in the M e e r w e i n - P o n n d o r f V e r l e y - O p p e n a u e r reactions. J.C. van der Waal a, P.J. Kunkeler a, K. Tan b and H. van Bekkum a a Laboratory of Organic Chemistry and Catalysis, Delft University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands b Tianjin Research Institute of Chemical Industry, Dingzigu, Hongqiao, 300131 Tianjin, PR China Zeolite titanium beta has been tested in the liquid- and gas-phase Meerwein-PonndorfVerley reduction of cyclohexanones and the Oppenauer oxidation of cyclohexanols. A high selectivity towards the thermodynamically unfavourable cis-alcohol was observed, which has been ascribed to transition-state selectivity in the pores of the zeolite. Under gas-phase conditions the dehydration of alcohols to cycloalkenes is observed as a side reaction. The catalyst was found to be active even in the presence of water and ammonia.

1. INTRODUCTION The Meerwein-Ponndorf-Verley reduction of carbonyl compounds and the Oppenauer oxidation of alcohols, together denoted as MPVO reactions, are considered to be highly selective reactions. For instance, C =C double bonds are not attacked. In MPV reductions a secondary alcohol is the reductant whereas in Oppenauer oxidations a ketone is the oxidant. It is generally accepted that MPVO reactions proceed via a complex in which both the carbonyl and the alcohol are coordinated to a Lewis acid metal ion after which a hydride transfer from the alcohol to the carbonyl group occurs (Fig. 1) [1]. Usually, metal sec-alkoxides are used as homogeneous catalysts in reductions and metal t-butoxides in oxidations [1]. RI-_

"')==o+ ;XR, Fig. 1

u

R3

-

.,

The Meerwein-Ponndorf-Verley-Oppenauer reaction [1].

Zeolites are crystalline metal oxides which have potential as regenerable heterogeneous catalysts in various organic reactions [2]. Because of their unique microporous structure, zeolites are especially promising in the field of shape-selective

1016 catalysis. As far as we know, only very few examples of the use of zeolites in MPVO reactions have been reported [3-7,9]. The reactions were carried out in the gas-phase over zeolites A, X and Y, exchanged or impregnated with alkali or alkaline-earth cations [3-5]. Shape-selectivity was only observed by Shabtai et aL[4] in the conversion of citronellal over zeolite X. It was shown that selectivity could be tuned by the size of the exchanged metal ion. In NaX there was enough space for the substrate to perform an intramolecular ring closure to isopulegol, whereas over CsX reduction to the linear citronellol was observed. Similar steric effects were also found for various other substrates. Recently Creyghton et aL [6,7] reported the use of zeolite beta in the MPVO reduction of 4-t-butylcyclohexanone. The high selectivity towards the thermodynamically less favoured cis-alcohol is explained by a restricted transition-state around a Lewis-acidic aluminium in the zeolite pores. When using an aluminium-free zeolite, titanium beta, in the epoxidation of olefins, we have shown that Ti-beta has acidic properties when alcoholic solvents were employed [8]. This was ascribed to the Lewis-acidic character of titanium in the zeolite framework. As we reported very recently [9], Ti-beta is found to be an excellent catalyst in MPVO reactions with a tolerance for water. Here, results are presented on the high selectivity, stability and low by-product formation of the catalyst, Ti-beta, in both the liquid-phase and gas-phase MPVO reactions.

2. EXPERIMENTAL

Zeolite titanium beta (Ti-beta) was synthesized according to van der Waal et al. [8,9] using di(cyclohexylmethyl)dimethylammonium hydroxide (DCDMA.OH) as the template. For a typical synthesis, 0.25 g titanium(IV) ethoxide (TEOT) was added to 30.0 g of a 19.5 % w/w DCDMA.OH solution and the mixture was stirred until all TEOT was dissolved. To facilitate the dissolution of TEOT, 1 ml of H20 2 (30 % w/w aqueous) was added. To the resulting clear solution 3.0 g Aerosil 200 (Degussa), 0.15 g seeds (all-silica beta [10] and 11.8 g water were added and the gel was aged for at least 24 h. After crystallisation (14 days at 140 ~ the zeolite (1.4 g) was filtrated, washed, dried and calcined at 540 ~ in air. Elemental analysis, performed using a LINK EDX system, showed a Si:Ti molar ratio of 76 and confirmed the absence of oligomeric titanium dioxide phases. Zeolite aluminium beta (Al-beta, Si:AI = 10) was prepared according to Wadlinger and Kerr [11] and all-silica beta (Si-beta, Si:A1 > 5000) was prepared according to van der Waal et aL [10]. Liquid-phase MPVO reactions were performed in 25 ml isopropanol (reductions) or 25 ml 2-butanone (oxidations) at 85 ~ using 2.5 mmol of the appropriate substrate: 4-t-butylcyclohexanone (4-Bu-ONE), 4-methylcyclohexanone (4-Me-ONE) or 4-t-butylcyclohexanol (4-Bu-OL, cis/trans mixture); 0.5 g zeolite or 0.25 mmol aluminium isopropoxide as the catalyst and 1,3,5-tri-t-butylbenzene as the internal standard. Samples were taken at regular intervals and analyzed by GC on a Carbowax CP-52 column and GC/MS. Gas-phase MPVO reactions were performed at 85 to 400 ~ in a fixed bed continuous down-flow reactor operated at atmospheric pressure under plug flow conditions. The catalyst, Ti-beta or N-beta (0.30 g), was diluted with 1.20 g a-quartz powder and processed to pellets then crushed to particles with a diameter of 0.7 - 1.0 mm. Reactant mixtures were pumped into a stream of preheated carrier gas (usually

1017 nitrogen) by means of a motor-driven syringe pump. The gas flow contained 10 vol.% isopropanol (reductions) or acetone (oxidations) and 1 vol.% of the appropriate substrate: 4-methylcyclohexanone (4-Me-ONE) or 4-methylcyclohexanol (4-Me-OL). The total gas flow was 50 ml/min and the molar gas flow of 4-Me-ONE or 4-Me-OL was 2.04 10-5 mol/min (WHSV = 2.9 gtotal goat-1 h'l) 9Samples of the reactor effluent were taken regularly and analysed on a CP-Sil-19 column. 3. RESULTS

3.1 Liquid-phase MPVO In the liquid-phase, the MPVO reactions catalysed by Ti-beta are highly selective and have low by-product formation. In the reduction of 4-t-butylcyclohexanone and 4-methylcyclohexanone a high selectivity (99 % and 98%, respectively) towards the thermodynamically unfavourable cis-alcohol is observed, similar to that observed with A1beta (Table 1). The high selectivity towards the cis-alcohol over beta-type zeolites was explained by Creyghton et aL [6] on the basis of transition-state selectivity. It can be seen in Fig. 2, that the transition-state leading to the cis- and trans-alcohol differ substantially in spatial requirements. The transition-states leading to the cis-isomer is aligned with the zeolite channel while that leading to the trans-alcohol occupies a much more axial position. Table 1 MPVO reduction of cyclohexanones over beta-type catalysts in refluxing isopropanol. Substrate

Catalyst

TOF a

Conversion b (%)

Selectivityb cis : trans

4-Me-ONE

Ti-beta

1.04

33.7

99:1

4-Me-ONE

M-beta

> 12

100.0

98:2

4-Me-ONE

AI(OPr)3

16.7

99.8

27:73

4-Bu-ONE

Ti-beta

2.26

64.9

98:2

4-Bu-ONE

Al-beta

> 12

100.0

95:5

4-Bu-ONE

Si-beta

4-Bu-ONE

AI(OPr)3

0.0 8.72

100.0

36:64

a Initial turn-over-frequency in mol ketone per mol of titanium or aluminium per hour. b Conversion and selectivity after 6 h reaction based on the initial amount of ketone added.

1018

H-~Me Me

....

/ Fig. 2

O~ /0

/

Proposed transition-states for the formation of cis-4-t-butylcyclohexanol (left) and trans-4-t-butylcyclohexanol (right).

In the Oppenauer oxidation of an equimolar mixture of cis- and trans-4-tbutylcyclohexanol, the cis-alcohol is converted selectively over both Ti-beta and Al-beta (Table 2), which is probably due to the same spatial restriction on the transition-states as for the reduction depicted in Fig. 2. In this way, the MPVO reduction and oxidation are in harmony as to the high transition-state selectivity. Table 2 Liquid-phase Oppenauer oxidation of 4-t-butylcyclohexanol (1:1 cis/trans mixture)with butanone as the oxidant. Catalyst

Conversion a of cis-alcohol (%)

Conversion a of trans-alcohol (%)

By-products a'b (mg)

Ti-beta

52.1

4.6

22.4

Al-beta

98.6 49.4 d

47.3 5.7c

317.3 63.8c

Ti(OiPr)4

6.1

12.9

22.7

Al(OiPr)3

5.3

7.4

56.5

a After 6 h reaction time. b Not MPVO related, predominantly from 2-butanone via aldol condensation, c After 15 min reaction time. d After 15 min. The most important side reaction in heterogeneously catalysed MPVO reactions is the acid-catalysed aldol condensation. Aldol products are usually observed during the Oppenauer oxidation of alcohols, when a surplus of ketone or aldehyde is used as the oxidizing agent and the solvent. The low amount of by-products formed when Tibeta was used as the catalyst, demonstrates the advantage of the titanium system over A1beta. This is probably caused by the much weaker BrCnsted acidity of the solvated titanium site [8] compared with the strong H+-acidity of the aluminium site in Al-beta. As we have shown earlier Ti-beta has a high tolerance towards water, which further shows the catalytic potential of Ti-beta in MPVO reactions [9].

1019

3.2 Gas-phase MPVO From an industrial point of view, gas-phase reactions are often preferred due to their ease of operations. The reduction of 4-methylcyclohexanone and the oxidation of cis- and trans-4-methylcyclohexanol over Ti-beta and Al-beta in the gas phase were studied at 100~ As can be seen from Fig. 3, both Ti-beta and Al-beta are active, but Tibeta has a considerably lower rate of deactivation. The deactivation of Al-beta is probably caused by the higher acidic strength of the protonic aluminium site compared to the non-protic titanium site. Another difference between the two catalysts is the pronounced dehydration of the alcohols formed over Al-beta. Two important differences between the gas-phase and liquid-phase reaction were observed. The most striking is the selectivity towards the cis-alcohol. Under liquid phase conditions (Table 1.) a selectivity of 99 % towards the cis-alcohol was observed, while in the gas-phase over Ti-beta only 53 - 62 % cis-alcohol was obtained. Furthermore, dehydration of the alcohols formed to 4-methylcyclohexene (4Me-ENE) was observed in the gas phase, while no trace of dehydrated products could be detected in liquid phase MPV reductions. The somewhat higher temperature of 100 ~ compared with 85 ~ for the liquid-phase experiments, could not explain this behaviour, since lowering the temperature to 85 ~ still resulted in a 4-Me-ENE selectivity of approximately 10 % over Ti-beta. The remarkable differences in selectivity between the gas- and liquid-phase MPVO reaction induced us to study the evolution of products as a function of the temperature, in the reduction of 4-methylcyclohexanone. From Fig. 4 it can be seen that the selectivity to the cis-alcohol decreases at higher temperatures; initially the selectivity to the trans-alcohol increases but at higher temperatures dehydration to 4-methylcyclohexene becomes the major reaction. At still higher temperatures (> 200 ~ the 4-methylcyclohexene is subsequently isomerised to the more stable 1-methylcyclohexene. I00

80

'N

60

20

0

2

4 Time on stream [h]

6

8

a) Ti-beta Fig. 3a

The gas-phase MPV reduction of 4-methylcyclohexanone with isopropanol over Ti-beta at 100~ 9 = conversion; 9 = cis-4-Me-OL; x = trans-4-Me-OL; n = 4-Me-ENE.

1020

100

80

~.43

------'~

!o A

o

-

i

A

4 ' Time on stream [h]

i

b) Al-beta Fig. 3b

The gas-phase MPV reduction of 4-methylcyclohexanone with isopropanol over Al-beta at 100~ 9 = conversion; 9 = cis-4-Me-OL; x = trans-4-Me-OL; [] = 4-Me-ENE. 100

80 ,___,

c~

.~ 40

0

Fig. 4

50

100

150 200 250 Temperature [ * C]

300

350

Influence of the temperature on the selectivity in the reduction of 4-methylcyclohexanone with isopropanol. Temperature increment was 0.2 ~ 9 = conversion; 9 = cis-4-Me-OL; x = trans-4-Me-OL; [] = 4-Me-ENE.

Since the 4-methylcyclohexene can be formed via dehydration of either of the two alcohols formed, both alcohols were tested in the Oppenauer oxidation, using acetone as the oxidant at 100 ~ It was observed that the cis-alcohol (Fig. 5a) is easily oxidized to the corresponding ketone, while the trans-alcohol showed a much lower activity. This is in accordance with the transition-states assumed for liquid phase MPVO reactions (Fig. 2). It can be observed further that the dehydration of the trans-alcohol only occurred to very limited scale (Fig. 5b) whereas dehydration of the cis-alcohol was an important side-

1021 reaction. This can be understood by assuming an E2-mechanism, in which the axial OH group in the cis-isomer is in the ideal position for water elimination. Another important side reaction for both alcohols was isomerization. It is therefore proposed that the alkene is formed mainly from the cis-alcohol and that at low temperatures the trans-alcohol can only be dehydrated if it is first isomerised to the cisalcohol via an MPVO transition-state. Comparing the deactivation of the Ti-beta catalyst in the oxidation (Fig. 5) and reduction reactions (Fig. 3) shows that the deactivation is more pronounced during oxidative conditions. This is probably caused by the high amount of ketones present, which easily form aldol condensates which may plug the zeolite channels, thus inhibiting access to the micropore system. 100

80

r~

..~ 40

i

r,.)

0

2

4 Time on stream [h]

6

a) Oxidation of cis-4-Me-OL 100

80

"~

60

.~ 40

I 2o

r.,)

0

2

4 Time on stream [h]

6

8

b) Oxidation of trans-4-Me-OL Fig. 5

The gas-phase Oppenauer oxidation of the 4-methylcyclohexanols with acetone over Ti-beta at 100~ a) cis-4-Me-OL and b) trans-4-Me-OL. 9 = conversion; 9 = cis-4-Me-OL; x = trans4-Me-OL; 9 = 4-Me-ONE. [] = 4-Me-ENE.

1022 The commonly used MPVO catalysts consist of metal alkoxides, which are easily hydrolysed to inactive oxides in the presence of water. Since the proposed catalytic species for the MPVO reaction also consists of an alkoxide intermediate [6,9] the influence of water and strong Lewis bases on the catalytic activity and selectivity was investigated. As already reported for the liquid-phase reaction [9], Ti-beta has a high tolerance for water due to its hydrophobic interior. As can be seen from Fig. 6a the presence of water is not detrimental to the activity of Ti-beta in the MPV reduction of 4-methylcyclohexanone. The temperature of 110 ~ at which a ketone conversion of 50% is measured, is identical to the temperature required for 50% conversion in the absence of water (Fig. 3), i.e. water has no effect whatsoever on the overall MPVO activity of the titanium site. 100

50

100

150 200 250 Temperature [ * Q

300

350

a) Reduction in the presence of water 100

9

--

w--r-~

80

60

.~

40

i 0 100

150

200 250 Temperature * C]

300

350

400

b) Reduction in the presence of ammonia Fig. 6

Temperature-programmed gas-phase MPV reduction of 4-methylcyclohexanone over Ti-beta in the presence of a) 2.66 vol.% water or b) 5 vol% ammonia. 9 = conversion; 9 = cis-4-Me-OL; x = trans-4-Me-OL; [] = 4-MeENE.

1023

The selectivity towards alcohols increased from 8 1 % to 97% at 85~ This enhanced selectivity can be ascribed to either a kinetic suppression of an irreversible dehydration or to a change in the alcohol/olefin equilibrium due to the higher amount of water present. In the case of a shift in equilibrium, it should be possible to oxidise olefins with ketones in the presence of water via in situ formed alcohols. In an attempt to oxidise cyclohexene with acetone in the presence of a water, no conversion to cyclohexanone was observed between 85 and 400 ~ It is therefore concluded that under the experimental conditions used, 4-methylcyclohexene is formed irreversibly from the cisalcohol (Fig. 7) and the increased selectivity towards alcohols should therefore be ascribed to a deactivation of the dehydration sites in the presence of water. OH

cis-4-Me-OL

4-Me-ENE

1-Me-ENE

Ti 4-Me-ONE

~

OH

trans-4-Me-OL

Fig. 7

Proposed reaction scheme for the MPV reduction of 4-methylcyclohexanone.

For liquid-phase reactions at 85 ~ we reported that small amounts of a strong base, e.g. pyridine, completely poisoned the catalyst [2]. It can be seen from Fig. 6b that, in the presence of ammonia, higher temperatures are required to reduce 4-methylcyclohexanone. The higher temperatures are required for the desorption of ammonia from the catalytically active site. The relatively low temperature of 305 ~ at which 50 % conversion is observed, suggests that the ammonia is not bonded to a strong acidic site. Since Br0nsted-acidic aluminium sites desorb ammonia at about 480 ~ this confirms that the MPVO reactions proceed via the titanium sites and not via any residual aluminium sites (Si:AI > 5000). The latter can also be concluded from Table 1; the all-silica analogue of zeolite beta [10] was found to be completely inactive even though it has a Si:A1 ratio similar to Ti-beta. It was also observed that in the presence of ammonia, no isomerisation of 4-Me-ENE to 1-Me-ENE occurred even at 400 ~ suggesting that the isomerisation requires strong Brcnsted acid-sites, most probably the residual aluminium sites. As was already shown in Fig. 3 and 5a, Ti-beta exhibits a low deactivation rate in the reduction of 4-methylcyclohexanone and a significantly higher deactivation rate in the oxidation of cis-4-methylcyclohexanol. In both cases, frequent regeneration of the catalyst will be necessary. Catalyst stability was tested by regeneration at 480~ in air after each run. No significant loss in activity, selectivity or product distribution was observed, even after 35 consecutive runs. Similar results were also obtained for the Tibeta used in liquid-phase reductions; after regenerating the Ti-beta catalyst 5 times, the same initial catalytic activity per gram of catalyst was observed.

1024 4. CONCLUSION Ti-beta is found to be an excellent catalyst in MPVO reactions under both liquid- and gas-phase conditions. Under liquid-phase conditions, a very high selectivity in the reduction of 4-substituted cyclohexanones towards the thermodynamically unfavourable cis-alcohols was observed. By-products were observed only during the oxidation of alcohols using ketone solvents and consisted primarily of aldol condensation products. Remarkable differences exist between the liquid-phase and gas-phase reactions under otherwise similar conditions. The selectivity towards the cis-alcohol is still above the thermodynamically expected value but significantly lower than under liquid-phase conditions. In contrast to the liquid-phase reactions, dehydration of the alcohols to the corresponding alkene is an important side-reaction. The oxidation of both the cis- and the trans-alcohol clearly showed that the olefin is exclusively formed from the cis-alcohol. Dehydration of the trans-alcohol is assumed to proceed by isomerisation via a MPVO mechanism to the corresponding cis-alcohol. The catalytic potential of the titanium-based catalyst is shown by the low amount of by-products formed over Ti-beta compared with Al-beta, the high resistance to water and the excellent stability of the catalyst with respect to regeneration.

ACKNOWLEDGEMENT Dr. Eddy Creyghton is thanked for valuable discussion and the Dutch Institute for Scientific Research (NWO/SON) for financial support. REFERENCES

1

For a review, see : C.F. de Graauw, J.A. Peters, H. van Bekkum and J. Huskens, Synthesis, 10 (1994) 1007. 2 P.B. Venuto, Microporous Mater., 2, (1994) 297; W. H61derich and H. van Bekkum, Stud. Surf Sci. CataL, 68, (1991) 631. 3 J. Shabtai, R. Lazar and E. Biron, J. Mol. Catal., 27, (1984) 35. 4 M. Huang, P.A. Zielinski, J. Moulod and S. Kaliaguine, Appl. CataI. A, 118, (1994) 33. 5 M. Berkani, J.L. Lemberton, M. Marczewski and G. Perot, Catal. Lett., 31 (1995) 405. 6 E.J. Creyghton, S.D. Ganeshie, R.S. Downing and H. van Bekkum, J. Chem. Soc. Chem. Commun., (1995) 1859. 7 E.J. Creyghton, S.D. Ganeshie, R.S. Downing and H. van Bekkum, J. Mol. Catal., in press (1997). 8 J.C. van der Waal, P. Lin, M.S. Rigutto and H. van Bekkum, Stud. Surf Sci. Catal., 105, 1093 (1997). 9 J.C. van der Waal, K. Tan and H. van Bekkum, Catal. Lett., 41 (1996) 63. 10 J.C. van der Waal, M.S. Rigutto and H. van Bekkum, J. Chem. Soc., Chem. Commun., (1994) 1241. 11 R.L. Wadlinger and G.T. Kerr, US patent, Appl. 3.308.069 (1967).

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1025

Selective Oxidation of Cyclohexane over Rare Earth Exchanged Zeolite Y

Emerson L. Pires, Martin Wallau, Ulf Schuchardt; Instituto de Quimica, UNICAMP, Caixa Postal 6154, 13083-970 Campinas-SP, Brasil; e-mail: [email protected]

Rare earth oxides are known to exhibit strong redox properties and to catalyze, in the presence of oxygen, the oxidation of hydrocarbons like n-butane, propylene and benzene to carbon dioxide and water. Recently the selective oxidation of cyclohexane to cylohexanol and cyclohexanone in the liquid phase catalyzed by SmC13 was reported. 1 Also Ce(IV) impregnated on a cation exchange resin was used as catalyst for the oxidation of alcohols to the corresponding carbonyl compounds. 2 Although these results demonstrate the performance of rare earth cations as catalysts for the oxidation of hydrocarbons, no attention has been given to the redox properties of rare earth cation exchanged zeolites, which are very well studied as acid catalysts in petrochemical processes. In this work we report first results on the redox properties of rare earth exchanged zeolites with FAU structure as catalyst for the oxidation of cyclohexane with tert- butyl hydroperoxide (TBHP) in the liquid phase. The catalysts were prepared from_ a sodium zeolite Y (NAY) with a Si/Al ratio of 3, by solid state ion exchange with CeC137H20, NdCI36H20, SmCI3"6H20 or YbCI36H20, respectively. The NaY was mixed with the rare earth chloride (REel3) in a molar ratio AI/REC13 of 3. The homogenized mixture was placed in a reaction tube, evacuated and heated to 450~ for 6 h. After cooling to room temperature, the solid was separated, carefully washed with water and dried at 120~ for 12 h. The catalysts were characterized by X-ray diffraction (XRD), FTIR spectroscopy and elemental analysis. The oxidation of cyclohexane was carried out in a suspension of the REY (200 mg) in 15.6 g of cyclohexane. 10 mmol of TBHP dissolved in 1.1 g of cyclohexane was added and the reaction mixture was stirred for 24 h under reflux. After filtering off the solid catalyst, the reaction products were quantified by gas chromatography, using internal standard and calibration curves. All peaks observed in the XRD patterns of the REY can be assigned to the FAU

1026 structure and the unit cell parameter ao, calculated by linear regression, does not vary compared to the parent NaY. The absence of a broad reflection between 20 and 25 o (20), which would indicate the presence of amorphous silica, demonstrates that the FAU structure remains intact after the solid state ion exchange. The decreased intensity of the reflections observed in the patterns of the REY may be attributed to an enhanced adsorption of the X-rays in the presence of the rare earth cations, rather than to a decreased crystallinity. It can be observed in the FTIR spectra that the ratio of intensities between the bands around 1000, 710 and 460 cm "1, attributed to internal tetrahedra vibrations (structure insensitive) 3 and the bands around 1130, 780 and 570 cm"1, attributed to the vibrations of the external linkages (structure sensitive), 3 do not differ between the parent NaY and the REY. This is a further confirmation that the FAU structure remains intact after the solid state ion exchange. The results of the catalytic oxidation of cyclohexane with TBHP are given in Table 1. The amounts of oxidized products obtained in a blank experiment are subtracted. Besides the main products cyclohexyl hydroperoxide (chhp), cyclohexanol (ol) and cyclohexanone (one), small amounts of other still unidentified products were also obtained. While NaY is inactive for the oxidation of cyclohexane, the activity of the REY increases in the following order: SmY NdY < YbY < CeY. An ol/one ratio close to unity, observed for NdY, SmY and YbY, indicates that both products are formed simultaneously from a common intermediate as already claimed for the cyclohexane oxidation catalyzed by SmC13.1 A probable reaction mechanism is the hydrogen abstraction catalyzed by the REY. The resulting cyclohexyl radicals react with molecular oxygen to cyclohexylperoxo radicals, which decompose in a bimolecular reaction to cyclohexanone, cyclohexanol and oxygen, or are reduced and protonated to cyclohexyl hydroperoxide. The enhanced cyclohexanol selectivity of CeY may be due to a selective decomposition of cyclohexyl hydroperoxide to cyclohexanol. The relative low activity of the REY is due to its high hydrophilicity, which leads to a rapid deactivation by adsorption of the polar reaction products. Studies to increase the activity of REY by using NaY with a higher Si/AI ratio are in progress. To the best of our knowledge, this is the first report on the redox activity of rare earth cations supported on zeolites in liquid phase oxidation.

1027

Table 1- Results of the oxidation of cyclohexane with TBHP catalyzed by REY catalyst NaY CeY NdY SmY YbY

ol (mmql) traces 1.14 0.22 0.39 0.39

.

one (mmol) traces 0.44 0.21 0.30 0.43

chhp (mmol) traces 0.74 0.56 0.22 0.46

ol/one 2.6 1.0 1.3 0.9

I. Yamanaka, K. Otsuka, J. Mol. Catal. A 1995, 95, 115 S. Kanemoto, H. Saimoto, K. Oshima, H. Nozaki, Tetrahedron Lett. 1984, 25, 3317 D.W. Breck, Zeolite Molecular Sieves, Wiley & Sons, New York, 1974

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1029

RATIONALLY DESIGNED OXIDATION CATALYSTS : FUNCTIONALIZED METALLOPORPHYRINS ENCAPSULATED IN TRANSITION METAL-DOPED MESOPOROUS SILICA (abstract) Lei Zhang, Tao Sun and Jackie Y. Ying Dept. of Chemical Engineering Massachusetts Institute of Technology, Cambridge, MA 02139, USA Catalytic oxidation by metalloporphyrins plays an important role in the conversion of both saturated and unsaturated hydrocarbons to valuable fine chemicals. The advantages in the development of metalloporphyrin systems that mimic cytochrome P-450 mono-oxygenases include substrate specificity, chemoselectivity and high catalytic activity under mild reaction conditions. Substantial efforts have been devoted to the development of an effective supported metalloporphyrin heterogeneous catalyst system that prevents self-oxidation of the active centers and allows for facile recovery of the catalyst. Mesoporous materials with a well-defined pore structure have recently attracted a great deal of research attention in catalytic applications. In this study, a rational design strategy has been established to prepare heterogeneous metalloporphyrin oxidation catalysts encapuslated in a hexagonally-packed mesoporous molecular sieve, providing (i) superior catalytic performance through a rational design of the matrix structure, and (ii) improved metalloporphyrin fixation through a tailored interaction between the catalyst and support. A series of mesoporous silicas known as MCM-41 has been synthesized with selective dopants for the metalloporphyrins. The microstructure of the materials was found to depend strongly on processing parameters, such as dopant concentration, aging time and temperature. In the case of Nb-doped silica (Nb/Si-TMS8,Nb/Si-TMS9), the formation of covalent bonding between the surface-exposed niobium sites and the functionalized groups of the iron porphyrin is crucial for immobilizing the iron complex within the mesoporous structure~ The fixation mechanism is different from the Coulombic forces and hydrogen bonding interaction involved in the conventional supported system. It effectively prevents leaching of the porphyrin and guarantees continued usage of the heterogeneous catalyst system. The well defined spacious mesoporous channels further allow for free diffusion of reactants and products. The functionalized metalloporphyrins, such as iron(III) meso-(tetra-aminophenyl)porphyrin bromide (FeTrva2PPr), encapsulated in transition metal-doped mesoporous silica showed high catalytic activity and improved selectivity for the epoxidation of olefins and hydroxylation of alkanes at ambient conditions. In the hydroxylation of cyclohexane, exclusive formation of cyclohexanol was observed with a yield of 53% after 6 hr. In epoxidation of cyclooctene, a single product of cyclooctene oxide was obtained with a conversion of 57% after 5 hr. No catalyst leaching was detected during the reaction over the supported catalystys, FeTNlI2PPBr/Nb/Si-TMS8 and FeTrcmPPBr/Nb/Si-TMS9. This novel rational design strategy provides a significant improvement on the fixation mechanism to achieve a better heterogneous catalyst. By manipulating the structural characteristics, such as the surface area, pore size, nature of the dopant and dopant concentration of the support material, as well as the functional groups of the iron porphyrin, the catalytic behavior can be controlled systematically via synthesis parameters.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1031

Catalytic Oxidations with Biomimetic Vanadium Systems I.W.C.E. Arends, M. Pellizon Birelli and R.A. Sheldon Laboratory for Organic Chemistry and Catalysis, Delft University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands.

ABSTRACT Schiffs base complexes of vanadium were encapsulated in the supercages of zeolite NaY and their catalytic activities were tested in the epoxidation of several alkenes and allylic alcohols with tert-butylhydroperoxide The complexes investigated were VO(HPS) (vanadylN-(2-hydroxyphenyl)salicylideneimine) and VO(salen) (vanadyl-N,N'(bis)salicylideneimine). Particular attention was devoted to the question of leaching of vanadium during reaction. 1. INTRODUCTION The widespread presence of vanadium-containing enzymes, such as bromoperoxidase, has only recently been appreciated 1. For example, vanadium-bromo-peroxidase (V-BrPO) catalyzes oxidative bromination of compounds in seawater. This action can be mimicked effectively to halogenate compounds for synthetic purposes 2. The mechanism involves oxidation of Br- to the brominating species in media with vanadium and hydrogen peroxide. The identity of the intermediate is still poorly understood 3. On the other hand peroxocomplexes of vanadium have long been known for their oxidative capabilities in organic media 4. Mo, W, Ti, and V are effective for the epoxidation of alkenes with ROzH. This process forms the basis for the commercial production of propylene oxide using either tert-butyl hydroperoxide or ethylbenzene hydroperoxide and Mo or Ti as catalysts 5. There is general agreement that the mechanism involves an alkylperoxometal complex which transfers an oxygen atom to the olefinic double bond. For a particular alkylperoxovanadium(V) complex the crystal structure has been determined, which showed coordination by both oxygen atoms 6. As to the intricate details of the oxygen transfer step, analogous to molybdenum, both direct nucleophilic attack of the olefin on the electrophilic oxygen of the alkylperoxometal species, and coordination of the olefin to the vanadium, followed by rate limiting insertion into the metal oxygen bond seem possible (Figure 1, pathways (a) and (b) resp.) 7. As epoxidation of unfunctionalized alkenes catalyzed by vanadium is rather slow compared to e.g. molybdenum, the ligand environment is critical in allowing the above described heterolytic mechanism to be dominant 5,7,8 Epoxidation of allylic alcohols by peroxometal complexes has been demonstrated to be even faster and more selective due to coordination of the allylic alcohol through its alcohol group.

1032 Vanadium shows exceptional reactivity in this respect 9, which is consistent with the strong affinity of vanadium(V) for alcohol ligands.

/,,,,,O

But

vV_o 0

"'"'~

Figure 1

But

\But

Our aim was to design a highly selective heterogeneous stable vanadium catalyst, optimal for epoxidation, which can be fine-tuned via its ligands. A biomimetic approach consists of encapsulating a complex in the cages of a zeolite. The inorganic backbone in this sense mimics the environment in the enzyme 1~ Recent examples of these so called "zeozymes" comprise the encapsulation of iron-phthalocyanines 11 and manganese. b~pyndlnes . The principle is that owing to space restriction the complex stays in the zeolite and cannot diffuse out, provided that the complex is stable. This is essentially different from isomorphously substituted vanadium molecular sieves, like vanadium-silicalites and V-APO's. Especially, the latter ones have demonstrated problems in leaching 13'14. Few examples of vanadium encapsulated complexes are known. Recently VO(bipyridyl)215 and VO(salen) ( 1 ) 16 - a Schiffs base complex - were incorporated in zeolite NaY and studied at room temperature. VO(salen) however is known to be not very selective for catalytic epoxidation of alkenes. Mimoun showed that complex (2) vvO(HPS)(OOtBu) is a stable and much more efficient epoxidation catalyst 3 and HPS (N-(2-hydroxyphenyl)salicylideneimine) is, therefore, an interesting choice as a ligand. In this study VO(HPS) was synthesized within the cages of zeolite NaY and tested as a catalyst for oxidation with TBHP. In order to reach significant reaction rates, our studies were performed at 70~ Different substrates were studied, including allylic alcohols, and an attempt was made to distinguish between catalysis in homogeneous solution, versus catalysis in such a catalytically constrained environment within the cage of a zeolite. VO(salen) was studied as a comparison. 9

9

12

O

.o

o. II/o \

!

(1)

(2)

1033 2. E X P E R I M E N T A L

2.1 Materials NaY was acquired from AKZO, with a Si/A1 ratio of 2.44, washed with aq. NaOAc and water before use. For a typical synthesis vanadium was introduced into the zeolites via ion exchange in 2 mM VOSO4 solutions. The zeolite was dried under nitrogen at 200~ and the dry material was mixed under nitrogen with a 2-5 fold excess of ligand and incubated for 6 h at 125~ (salen) or 190~ (HPS). Incubation with HPS gave a brown material, whereas with salen a yellow-orange material was obtained. The zeolite was soxhletted with acetonitrile or acetone and consequently with dichloromethane. Differences in treatment are given in Table 1. HPS (N-(2-hydroxyphenyl)salicylideneimine) was synthesized via condensation of 2aminophenol and salicylaldehyde in boiling ethanol and recrystallized from methanol. Homogeneous vvO(HPS)(OEt)[EtOH] and vVo(salen)(OEt) complexes were synthesized by complexation of VO(i-OPr)3 and the ligand in ethanol 17 Complexes were checked with 51V and 1H-NMR. Materials; anhydrous TBHP was made by extracting a 70% TBHP solution (in water) in chlorobenzene and drying over 3~, molecular sieves, to result in a 5 M solution. Cyclohexene and cyclooctene were distilled and were passed through a column of basic alumina before use. Vanadyl(IV) sulphate hydrate, Salen (N,N'-bis(salicylidene)ethylenediimine), 1,2,-dichloro-ethane (DCE), dichloromethane (DCM), acetone, and acetonitrile were reagent grade and used as received. (-) Carveol was purchased as a 1.48:1 mixture of trans- and cis-carveol and used as such. Cumylhydroperoxide was purchased as a 80% solution in cumene and used as received.

2.2 Catalyst characterization The crystal structure was checked by X-Ray diffraction using a Philips PW 1877 automated powder diffractometer with CuK~ radiation. AAS and ICP-AES measurements were carried out with Perkin Elmer instruments type 1100, 5000 Z and Plasma-40. UV/VIS measurements were carried out on a Varian Cary spectrometer. Diffuse reflectance spectra were recorded against a barium sulphate reference on the Varian Cary-3 and reflectance spectra were converted according to the Kubelka-Munk equation.

2.3 Catalytic oxidation experiments The catalytic oxidation experiments were carried out in a round bottom flask equipped with condenser and stirrer. Typically, 6 or 12 mmol of substrate, 2 mmol bromobenzene (internal standard), 9 ml solvent (1,2-dichloroethane) and 100 mg zeolite (which contains typically around 2.9 lamol V (0.15wt%), TBHP/V ratio = 2070) were heated to 70~ after which 6 mmol of TBHP in a 5.3 mmol chlorobenzene solution (which at the same time can function as an internal standard) were added to start the reaction. A sample was taken immediately afterwards. Before and during the reaction the mixture was purged with nitrogen for oxidations with cyclohexene and cyclooctene. Samples were filtrated over cotton wool and/or alumina, and triphenylphosphine was added to remove TBHP. In case of acetone as the solvent at 70~ reactions were performed in a 50 ml autoclave and the reaction mixture was only purged with nitrogen before heating. After reaction TBHP was determined by iodometric titration.

1034 Quantitative analyses were carried out by GC with a semi-capillary column, CP-Sil-5 CB (50 m x 0.53 mm) for cyclooctene or on a CB wax 52 CB (50 m x 0.53 mm) for cyclohexene and carveol. Qualitative identification of peaks was made by reference samples and GC/MS analyses. Also cyclohexenyl-tert-butylperoxide could be analyzed in this way. Quantitative analysis was performed by using molar responses with respect to the internal standard. 3. RESULTS 3.1 Characterization Schiffs base complexes exhibit very intense UV absorptions. In Figure 2, UV spectra in Nujol and/or DRS spectra are shown for the synthesized materials, VO(HPS)-Y and VO(salen)-Y and their homogeneous counterparts. The spectra for the homogeneous complexes were recorded in DMF, to which two drops of a 0.01 M NaOH solution was added. Apparently in this matrix ethoxy ligands are replaced by OH as a ligand and the characteristic absorption maxima at 424 nm for VO(HPS) and 360 nm for VO(salen) can be seen. This environment closely resembles the ligand environment in the zeolite, where vanadium is surrounded by oxygens from the framework. In dichloromethane for VO(HPS)(OEt) only shoulders at 309 and 370 nm, and for VO(salen) at 260 and 290 nm were observed. Spectra of in-situ generated complexes of vIVoso4 and HPS or salen look identical to those of the preformed V v complexes. Figure 2 clearly indicates that the complexes are present within the zeolites, and for VO(salen) a 20 nm shift occurs upon complexation. Separate Shiffs bases display maxima at 350 and 317 nm for HPS and salen resp.

~(iii)

VO(HPS)

~,i)

VO(salen)

(i) i

200

400 600 wavelength (nm)

800

200

400 600 wavelength (nm)

800

Figure 2a + b: UV/VIS spectra for VO(HPS) and VO(salen). (i) homogeneous complex, 0.1 mM dissolved in DMF + OH added (ii) DRA spectra for VO(HPS)-Y [B] and VO(salen)-Y [HI (c) Solution spectra in nujol for VO(HPS)-Y [D].

ICP analysis indicates that in a typical procedure the V content decreases from approximately 0.23-0.30 wt% before ligand incubation, to 0.15-0.29 wt% V after incubation with HPS and soxhlet extraction, apparently depending on the efficiency of the incubation. Washing with NaOAc, followed by soxhlet extraction with acetone removes another 0.03 0.13 wt% of vanadium. Data are given in Table 1. VO(salen) [J] was prepared with a 50 fold excess of salen and 24 h incubation, approaching as much as possible the conditions in ref. 16.

1035

One unit cell in the faujasite structure contains 8 supercages and in our case 58 A1 atoms. A ratio of 58/1 for A1/V therefore corresponds to 1 V atom per 8 supercages. Most materials used contain 1 vanadium per 8 to 16 supercages. Table 1. Vanadium content as determined by ICP, after several treatments.

VO-Y [A] VO(HPS)-Y [B] VO(HPS)-Y [C] VO(HPS)-Y [D] VO(HPS)-Y [E] VO-Y [F] VO(HPS)-Y [G] VO(HPS)-Y VO(salen)-Y [H] VO(salen)-Y [I] VO(salen)-Y [J]

details treatment

V(wt%)

AI/V (mol/mol) a

2mM exchange at RT [A] + acetonitrile soxhlet [B] + NaOAc wash + soxhlet [A] + acetone soxhlet [D] + NaOAc wash + soxhlet 4mM exchange at 80~ [F] + acetone soxhlet, + NaAc wash + soxhlet [G], washed TBHP/70~ low wt% VO-Y b + acetonitrile soxhlet [A] + acetonitrile soxhlet [A] + 24h incubation + acetone soxhlet

0.23 0.15 0.12 0.29 0.14 1.2 0.61 0.36 0.017 0.076 0.30

68 n.d. 142 62 132 n.d. 32 47 957 251 62

(a) In some cases no reliable data were obtained for the AI/V ratio, and could therefore not be determined (n.d.). (b) VO-Y synthesized by exchange of NaY with 2mM solution of VO(i-OPr)3 in demineralized water.

3.2 Oxidation experiments with cyclohexene In Table 2, the first experiments with cyclohexene as a substrate are given. The necessary blanks and homogeneous experiments are given as a comparison. As can be seen, it is very difficult with vanadium to get 100% selectivity towards the epoxide. Especially at the reaction temperatures of 70~ allylic oxidation interferes. In an attempt to reduce its influence, reactions were performed under nitrogen. As a consequence the main product becomes the cyclohexenyl-tert-butylperoxide, according to reactions 1-7. The In (itiator) could either be tBuOo, tBuOO~ or VOO~ /---X In.

+

+ InH

~

(1)

tBuOOH

+ In.

~

~~

+ 02

~

~~'-OO 9

(3)

~ .

+tBuOO"

,~

~-OOtBu

(4)

~

~--O

~---~-OO-

~-OO.

2 tBuOO.

+tBuOO"

tBuOO 9 + InH

+ ~~-OH

~ ~ = O

~

tBuO. + 02

(2)

+ O2

(5)

+tBuOH + O2

(6)

(7)

1036

Radicals are formed upon one-electron oxidation of tBuOOH 5, but also a v~Vooo diradical has been suggested as an intermediate 6. As can be seen from the formation of ketones and alcohols, it is impossible to completely eliminate the influence of oxygen, also because oxygen is formed in situ according to reaction 7. Table 2. Oxidation of cyclohexene catalyzed by vanadium a catalyst...... time yield epoxyb selectivity product selectivities %c (%) epoxide (on TBHP epoxide, CyOOtBu, Cy=O+CyOH consumed) no VO(acac)2 VO(HPS) VO(HPS) VO-Y [A] VO(HPS)-Y (B) VO(HPS)-Y (D) VO(HPS)-Y (D) VO(salen) VO(salen) VO(salen)-Y [I]

lit e g h i j RT lit e

23h 5h 3h 5h 24h 48h 5h 24h 3h 5h 21h

0 25 32 55 18 (148 TON) 23 5 0.3 (6 TON) 0.1 4 0.6 (12 TON)

23 43 n.d. n.d. n.d. n.d. n.d. 1 2 n.d.

46

44

100d 0

43 f

1f

1f

57 34 33 72 100 1f 5 19

17 58 28 19 77 50

25 7 39 9 10f 18 31

-::iai-i~eaction condiiions; as in experimental, 70~ i00 mg cat (zeolite) 0r 0.06 mmol complex, solvent~gml DCE, 0.55 M TBHP, TBHP:cyclohexene = 1:1. n.d. = not determined (b) Yield on intake cyclohexene. (c) Defined as product/total products. CyOOtBu, refers to cyclohexenyl-tert-butylperoxide, Cy=O + CyOH comprises cyclohexenone (major), cyclohexenole, epoxycyclohexenole (major) and epoxycyclohexenone as products. Mass balance was usually between 60 and 80% due to vaporization of cyclohexene. (d) 3% cyclohexenole (on intake cyclohexene) formed as sole product. (e) Ref.7, conditions; 60~ solvent DCE, 1.40 M cyclohexene, 1.43 M tBuOOH, cat. 0.07 M. (f) Products in % on TBHP consumed. (g) After 0.75h, 14% yield of epoxy, and 67% product selectivity to epoxide, 17% CyOOtBu and 16% Cy=O+CyOH.(h) After 5h, 14 TON and 26% selectivity to epoxide.(i) only 50 mg cat. (j) ratio cyclohexene:oxidant = 2:1. The oxidations catalyzed by VO(HPS)-Y, gave slow reactions, with rather low selectivities. The latter phenomenon is likely a direct result of the lower rate, giving allylic oxidation more chance to interfere. At room temperature almost no reaction was observed. The VO(salen)-Y material tested here gave a limited activity, although the selectivity towards the epoxide was higher than that obtained with the homogeneous VO(salen) complex.

3.3 Oxidation experiments with cyclooctene Conversion of cyclooctene to epoxycyclooctene was studied with different catalysts. As can be seen products of homolytic oxidation are much less prominent compared with cyclohexene. Heterolytic epoxidation of cyclooctene compared to cyclohexene is a faster reaction, and in Table 3 it is shown that whereas VO(salen) did not give significant epoxide yields with cyclohexene, with cyclooctene 78% epoxide selectivity was found. The cyclooctene oxidation is therefore less sensitive to ligand effects, and the reaction was mainly used to study catalyst stability. Experiments to test for leaching were performed in two ways. In the first method the reaction mixture was filtered after 1 h, and the filtrate allowed to react further (Table 3, note g). It is important to filter the solution at the reaction temperature because readsorption can take place on cooling. In the second method the zeolite is incubated

1037

with T B H P for 1 h at 70~ filtered and the reaction started by adding c y c l o o c t e n e to the filtrate (note h) 18. As s h o w n in Table 3 the results were very similar, although in general a slightly l o w e r selectivity to the epoxide was observed after filtering the catalyst as well as a lower selectivity on T B H P consumed. Recycling or extensive incubation with t B u O O H , gave d i m i n i s h e d activity. A p p a r e n t l y T B H P is capable o f extracting the v a n a d i u m from the m o l e c u l a r sieve. Table 3. Oxidation o f cyclooctene, catalyzed by v a n a d i u m a catalyst time yield epoxy sel. epoxlde (on (%) TBHP consumed) no d 24h 7e 23 VO(HPS) 5h 62 70 VO(HPS) d 5h 53 63 VO(HPS)-Y [B] d 24h 28 n.d. VO(HPS)-Y [C] 24h 26 49 VO(HPS)-Y [D] 24h 50 75 VO(HPS)-Y [D] recycle f 24h 26 59 VO(HPS)-Y [D] leach g 24h 43 62 VO(HPS)-Y [D] leach h 24h 33 n.d. VO(HPS)-Y [D] RT ~ 24h VO(HPS)-Y [a] 24h 48 65 VO(HPS)-Y [G] leach h 24h 50 52 VO(HPS)-Y [G] leachj 24h 29 38 VO(HPS)-Y [E] CHP k 24h 5 n.d. VO(HPS)-Y [E] CHP k 24h 10 n.d. VO(salen) d 5h 26 25 VO(salen)-Y [I] l d 24h 1.5 18 VO(salen)-Y [J] 24h 0 "

-

.

.

.

.

.

.

.

.

.

.

~ . . . . . . . . . . .

" .......

13

. . . . . . . . . . .

' .

.

.

.

.

.

.

.

9

'~

....

'

........

product selectwmes (%) epoxide Cy=O CyOOtBu 52 5.0 33 100 97 1.4 0.7 95 2.3 0.7 96 0.7 1.3 97 <0.5 1.1 85 2.6 4.7 82 5.8 4.3 79 6.2 3.7 91 97 81 100 100 78 100

1.8 < 1.0 3.6 6.7

3.9 2.0 6.2 14

(ai:~::Gaci:(0-n:::con'ditions; 70o~C;~]~00mg zeolite or 0.06 mmo[VO(HPS)I 9ml DCE, 0.55 MTBHP, cyciooctene: TBHP = 2 : 1 . (b) Yield on input TBHP. (c) Defined as product/total products. Mass balances are usually around 95%. Cy=O is cyclooctenone (2 isomers), and CyOOtBu refers to the cyclooctenyl-tert-butylperoxide. The rest of the products were cyclooctanediol (usually <2%) and epoxycyclooctenole (up to 10%). (d) Ratio cyclooctene:TBHP = 1:1. (e) After 5h, 4% epoxide is formed. (f) Zeolite from reaction before. (g) One hour after start of reaction, the reaction mixture is filtered and the reaction continued. (h) The zeolite was incubated l h with TBHP at 70~ filtered and cyclooctene was added to start the reaction. (i) Reaction performed at RT. (j) The Zeolite was first washed with TBHP at 70~ for three hours, rinsed with solvent and then used in a leaching experiment as described under (h). (k) Cumylhydroperoxide is used as the oxidant, ratio CHP:cyclooctene = 1:1. (1) Washed with NaOAc and soxhletted with DCM. A possibility is that tBuO-, f o r m e d through reaction o f t B u O O H with vanadium, or t B u O O itself replaces H P S as a ligand, resulting in leaching o f the n e w l y f o r m e d V - c o m p l e x . At r o o m t e m p e r a t u r e no oxidation was observed. As to VO(salen), the material [J], with highest v a n a d i u m incorporation and best complexation, appeared to be inactive, w h i c h could be due to pore blockage. The use o f c u m y l h y d r o p e r o x i d e as an oxidant did not lead to less leaching. One might have e x p e c t e d that if RO- or ROO- replaces H P S as a ligand, the use o f P h C ( C H 3 ) 2 O O H instead o f t B u O O H w o u l d diminish the leaching due to steric reasons. H o w e v e r , an even higher activity w a s o b s e r v e d after filtering the catalyst.

1038

3.4 Oxidations with allylic alcohols As a test reaction a mixture of cis- and trans-carveol was oxidized with TBHP. A n a l o g o u s to what has previously been observed for compounds like geraniol and linalool, epoxidation was fast and selective 9 (reaction 8). Trans-carveol was converted only to the corresponding epoxide, whereas cis-carveol gave a mixture of cis-epoxide and carvone. With V O ( H P S ) - Y [E] as the catalyst, reaction in the presence of the zeolite gave a m u c h better selectivity towards cis-epoxide, than without zeolite present. This suggests that carveol is able to c o m p l e x with the vanadium present within the molecular sieve. Moreover, the selectivity on TBHP consumed decreased from 95% to 45%, suggesting that the ligand environment has changed.

.......OH

(8a)

......OH ..

"V" TBHP trans-carveol

"v"

jo

o

(8b)

TBHP

cis-carveol

Table 4. Oxidation of cis- and trans-carveol, catalyzed by vanadium a catalyst solvent time yield yield trans-epoxide b cis-epoxide b no VO(acac)2c VO(HPS) d VO(HPS) VO-Y VO(HPS)-Y [B] VO(HPS)-Y [E] e VO(HPS)-Y [E] f VO(salen) VO(salen)-Y [H] VO(salen)-Y [J] VO(salen)-Y [J]g

acetone DCE DCE acetone DCE acetone DCE leach

DCE DCE DCE leach

5h 1h 5h 5h 18h 43h 24h 24h 5h 22h 24h 20h

0 100 100 100 100 61 77 66 100 6 23 31

0 34 62 31 25 5 8 2 52 1 3

ratio cis-epoxide /carvone 6% carvone 0.89 2.2 1.5 0.79 0.14 0.41 0.042 1.4 0.042 20% carvone 0.071

....(a) Reaction C0nditi0ns, as in the experimental, 70~ 100 mg zeolite or 0.06 mmol complex; 0.55 M TBHP, TBHP:cis+trans carveol = 1:1. Trans:cis carveol = 1.48:1. (b) Yield on input carveol; mass balances always 100+5%. Selectivity for trans-carveol to the epoxide always 100%.(c) Trans-carveol already fully converted in 5 min. (d) Trans-carveol fully converted in three hours. (e) After 4h, 5% trans-epoxide was formed, selectivity for the formation of epoxide on TBHP consumed was 95% after 24h. (f) As previous experiment, leaching method, note h, Table 2. After 4h, 18% trans-epoxide formed, selectivity of total epoxide on TBHP consumed = 45% after 24h. (g) As previous experiment; leaching method, note h, Table 3.

1039 Furthermore the influence of acetone as a solvent was examined. For the homogeneously catalyzed epoxidation reaction in general, lower yields and selectivities are obtained with acetone as solvent. However for the transport of substrates and oxidant in and out of the zeolite channels, acetone could have a beneficial effect. This could not be confirmed by our experiments. 4. Discussion and conclusions. We were able to synthesize an active zeolite-encapsulated VO(HPS,) complex. However epoxidation catalysis seems to take place mostly outside the zeolite. Apparently the TBHP is capable of extracting vanadium from the sieve, probably by replacing HPS as a ligand. It has previously been reported 16 that VO(salen)-Y affords epoxycycohexene with 100 turnovers in 24h at room temperature, with TBHP as the oxidant and chloroform as the solvent. However, vanadium was detected in the solution afterwards. When trying to reproduce these results, we observed some activity with materials that were incubated with small excesses of salen, but not at room temperature. It appears that if one is able to efficiently complex all the vanadium and to maintain encapsulation, almost no reaction takes place. In literature VO(bpy)2-Y 15 was able to activate TBHP as well as H202 and PhIO for the epoxidation of cyclohexene. Turnover frequencies of 200/day were observed, which is comparable to our results at 70~ with VO(HPS)-Y. We emphasize, however, that oxygen transfer with the latter two oxidants presumably involves different active oxidants. Vanadiumperoxo intermediates, for example, formed with hydrogen peroxide might be better accommodated. The zeolite supercages of 13 ~, in diameter, with intersecting channels of 7.4/;,, might have problems accommodating the alkylperoxo mechanism as depicted in Figure 1. On the other hand it seems that if the reaction is fast enough (faster than leaching) as with carveol, reaction does take place inside as well. It must be emphasized that alcohol and traces of water retard epoxidations catalyzed by vanadium, and in this respect the NaY zeolite which is quite hydrophilic, might contribute to this effect.

Acknowledgement The research of Dr. I.A. has been made possible by a fellowship of the Royal Netherlands Academy of Sciences.

References

[1] [2] [3] [4]

(a) R. Wever, B.E. Krenn, in N.D. Chasteen (ed.), Vanadium in Biological Systems, Kluwer, Boston, 1990, pp. 81; (b) A. Butler, J.V. Walker, Chem. Rev. 93 (1993) 1937; (c) A. Messerschmidt, R. Wever, Proc. Natl. Acad. Sci, 93 (1996) 392. (a) M. Bianchi, M. Bonchio, V. Conte, F. Coppa, F. Di Furia, G. Modena, S. Moro and S. Standen, J. Mol. Catal, 83 (1993) 107. (b) M.J. Clague, A. Butler, J. Am. Chem. Soc., 117 (1995) 3475. V. Conte, F. Di Furia, S. Moro and S. Rabbolini, J. Mol. Catal. A, 113 (1996) 175. R.A. Sheldon, J. Mol. Catal. 7 (1980) 107; (b) K.A. Jorgensen, Chem. Rev., 89 (1989) 431; (c) A. Butler, M.J. Clague and G. E. Meister, Chem. Rev. 94 (1994) 625.

1040

[5]

[6] [7]

[8]

[9] [10] [11] [12] [13] [14]

[15] [16] [17] [18]

(a) R.A. Sheldon, in W. Herrman and B. Cornils, (eds.), Applied Homogeneous Catalysis with Organometallic Compounds, Vol 1, VCH, Weinheim, 1996, pp. 411; (b) R.A. Sheldon and J.K. Kochi, Metal Catalyzed Oxidations of Organic Compounds, Acad. Press, New York, 1981. H. Mimoun, P. Chaumette, M. Mignard, L. Saussine, J. Fischer and R. Weiss, Nouv. J. Chim. 7 (1983)467. H. Mimoun, M. Mignard, P. Brechot and L. Saussine, J. Am. Chem. Soc. 108 (1986) 3711. H. Mimoun, L. Saussine, E. Daire, M. Postel, J. Fischer and R. Weiss, J. Am. Chem. Soc. 105 (1983) 3101. K.B. Sharpless, T.R. Verhoeven, Aldrichimica Acta, 12 (1979) 63. R.F. Parton, D.E. de Vos and P.A. Jacobs, in E.G. Derouane et al. (eds.), Zeolite Micoporous Solids: Synthesis, Structure and Reactivity, Kluwer, Dordrecht, 1992, pp. 555. N. Herron, G.D. Stucky and C.A. Tolman, J. Chem. Soc., Chem. Comm. (1986) 1521. P.P. Knops-Gerrits, D. de Vos, F. Thibault-Starzik and P.A. Jacobs, Nature, 369 (1994) 543. M.J. Haanepen, A.M. Elemans-Mehring and J.H.C. van Hooff, Appl. Catal. A., in press. J.S. Reddy, P. Liu and A. Sayari, Appl. Catal. A, 148 (1996) 7. P.P. Knops-Gerrits, C.A. Trujillo, B.Z. Zhan, X.Y. Li, P. Rouxhet and P.A. Jacobs, Topics in Cat. (1996) 437. K.J. Balkus Jr., A.K. Khanmamedova, K. M. Dixon and F. Bedioui, Appl. Catal. A. 143 (1996) 159. M.J. Clague, N.L. Keder and A. Butler, Inorg. Chem. 31 (1993) 4754. H.E.B. Lempers, R.A. Sheldon, Stud. Surf. Sci. Catal. 105 (1997) 1061.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1041

H i g h l y s e l e c t i v e p h o t o c h e m i c a l a n d d a r k o x i d a t i o n o f h y d r o c a r b o n s b y 0 2 in zeolites H. Frei Laboratory of Chemical Biodynamics, MS Calvin Laboratory, Lawrence Berkeley National Laboratory, University of California, Berkeley, CA 94720, USA

1.

INTRODUCTION

In partial oxidation of low alkanes, olefins, and small aromatics, molecular oxygen is the oxidant of choice because of economic factors. Yet, autoxidation of small hydrocarbons is inherently unselective, whether conducted in the gas or liquid phase, or whether catalyzed by transition metals or not [1-4]. One reason is that the desired products such as alcohols or carbonyls are more easily oxidized by 0 2 than the parent hydrocarbon. Overoxidation can only be minimized by keeping conversions low. Another factor is diversion of the radical chain reaction leading to the primary product (allyl, alkyl, benzyl hydroperoxide) by termination steps that result in the formation of oxy radicals. The highly reactive oxy radicals undergo several competing reactions that lead to a multitude of products. Hence, oxidations by 0 2 exhibit most often little chemo- or regioselectivity. A major challenge in the field of hydrocarbon + 0 2 chemistry is, therefore, to find reaction paths that afford the primary product with high selectivity at high conversion. We have developed a method that affords partial oxidation of small alkanes, alkenes, and alkyl substituted aromatics by 0 2 to commercially important products at very high selectivity. The approach is based on access to hydrocarbon.O 2 charge-transfer states by loading the gas phase reactants into cation-exchanged zeolites. These ionic states of hydrocarbon.oxygen collisional pairs are strongly stabilized by the high electrostatic fields inside the pores of alkali or alkaline-earth zeolites, such as zeolite Y, in which collisional pairs of light hydrocarbons and 0 2 can be formed at high concentration. Energies of charge-transfer states are lowered sufficiently to allow direct access by visible photons without the need of a photosensitizer, or even in the absence of light by thermal excitation in some cases. Selectivity is a direct result of the low energy path of these reactions and the positional constraints imposed by the zeolite pores. In this paper, we present a brief overview of physical aspects of zeolites crucial to our chemistry, and then describe the most interesting reactions found so far.

2.

S O M E I M P O R T A N T P H Y S I C A L ASPECTS OF Z E O L I T E S

Zeolites are crystalline aluminosilicates that exist in a variety of structures [5]. Ideal for our purpose are zeolites with faujasite structure, especially those of type Y (Si/A1 ratio of 2.4). The structure consists of a three-dimensional array of 13 ,~ spherical cages (so-called supercages) which are connected by windows of 8 ,~ diameter (Fig. 1). The size of the cages is just right for the formation of high concentrations of collisional pairs of oxygen and small hydrocarbon molecules. At room temperature, small molecules can roam freely inside the

1042 supercage network of the dehydrated zeolite (nanosecond hopping time from cage to cage) [6]. Hence, what one has in the zeolite is a high steady state concentration of short-lived reactant collisional pairs.

visible light I

>

Figure 1. HydrocarbonoO 2 collisional pair in a supercage of zeolite Y. An important ingredient for our reactions are the charge-compensating cations inside the zeolite cages. In the case of the materials used in most of our studies, namely Na + and Ba 2+exchanged zeolite Y, there are 3 to 4 cations located in each supercage. The wall of the cage carries a formal negative charge of 7, which resides on the framework oxygens. The electric shielding of the cations by these framework oxygens is poor, resulting in high electrostatic fields inside the supercage, especially in the vicinity of the cations. They are manifested, for example, in the induced infrared absorption of the fundamental vibration of 0 2 or N 2 gas loaded into the zeolite. The vibrations are infrared forbidden but become active in the presence of an electrostatic field. Fig. 2 shows the induced infrared absorption of 0 2 at 1555 cm-1 in NaY and BaY at a loading of 1-2 molecules per supercage. Intensities are proportional to the square of the field, and values of 0.3 V/k-1 (NAY) and 0.9 V ,~-1 (BAY) are derived from such measurements [7]. This method has previously been used by Cohen de Lara for the study of electrostatic properties of zeolite A [8]. Electrostatic fields in various zeolites have also been determined by other methods such as ESR of guest radicals [9] and analysis of electron densities from X-ray data [ 10]. They agree well with predictions from model calculations [ 11 ]. Aside from these electrostatic fields, the zeolite matrices we use are inert. Specifically, zeolites are on purpose kept free of acid sites (Br6nsted, Lewis acid sites). These are catalytic sites which play a crucial role in hydrocarbon cracking in the petroleum industry and synthetic applications in the chemical industry. Na + and Ba2+-exchanged zeolite Y used in our work can readily be prepared free of acid sites [12].

1043 0.03-

0.02o o

O ,.o

0.01-

<

0.00-

-0.01, ~ 1500 1520 1540 1560 1580 1600cm-1 Wavenumbers

Figure 2. Infrared fundamental absorption of 0 2 induced by the electric field of the NaY and BaY cage.

3. 3.1.

VISIBLE L I G H T DRIVEN OXIDATIONS

Selective oxidation of toluene to benzaldehyde When we loaded a dehydrated BaY matrix with toluene and 0 2 from the gas phase and exposed it to visible light at room temperature, we found that benzaldehyde is produced without side reaction [ 13]. The chemistry was monitored in situ by FT-infrared spectroscopy. Zeolites are transparent to infrared light except for some regions below 1200 cm -1 where framework vibrations absorb. Bands of small organic guest molecules are sharp, which makes identification of products by comparison with spectra of authentic samples and by isotopic labeling straightforward. We used the visible output of a tungsten lamp or the emission of a CW dye laser for photolysis, but did not note any difference in products or yields. When we ran the photochemistry at -70~ benzyl hydroperoxide was trapped. Warm-up of the matrix to ambient temperature resulted in dehydration to benzaldehyde, thus establishing the hydroperoxide as a reaction intermediate (Scheme 1). Benzaldehyde was the sole final oxidation product even upon conversion of as much as half of the toluene loaded into the zeolite matrix. Such toluene to benzaldehyde oxidation by 0 2 without overoxidation or side reaction has not been achieved by any other method. Cobalt (III) catalyzed autoxidation of toluene in solution practiced currently in industry lacks this selectivity, primarily because of further

1044 oxidation of the aldehyde to benzoic acid [ 1,2]. Benzaldehyde is an industrial intermediate for the synthesis of agrochemicals, flavors and fragrances.

//o CH 3 ~ ~

CH2OOH + O2

~~,<600nm BaY CaY ,

~

C ~

~

+H20

Scheme 1. Selective oxidation of toluene to benzaldehyde While alkali or alkaline-earth zeolite Y has no optical absorption in the visible, zeolite BaY or CaY loaded with toluene and 0 2 showed a continuous absorption tail extending into the green spectral region. This band appears only when toluene and 0 2 are simultaneously present in the zeolite, which means that it originates from an electronic transition of a collisional complex of the two molecules. Pressed zeolite pellets strongly scatter visible light, hence most spectroscopic studies in the visible/UV were done by diffuse reflectance measurements [13,14]. Recently we have been able to monitor the same hydrocarbon.O 2 absorption by transmission spectroscopy of monolayers of large (40 micron) zeolite Y crystals [15], thus confirming the results from diffuse reflectance measurements. Of the two assignments for this band we can conceive of, namely, an O2-enhanced triplet absorption of the hydrocarbon, or a toluene~ 2 charge-transfer transition, only the latter is consistent with our data. UV absorptions of hydrocarbon.O 2 charge-transfer complexes were discovered in the 1950s in high pressure 0 2 gas phase and O2-saturated liquids by the group of Evans and Mulliken [ 16,17]. These transitions appear as continuous absorption tails whose onset lies in the UV and depends on the ionization potential of the hydrocarbon. The interaction of the high electrostatic field inside the zeolite supercage with the large dipole generated upon excitation of the tolueneoO2 pair to the charge-transfer state (about 15 Debye) results in a stabilization of the excited state by 1.5 eV (-~t.E) (generally between 1 and 3 eV depending on hydrocarbon and exchanged cation). The result is a very large red shift of the absorption from the UV into the visible region. This electrostatic field effect of the zeolite cage on the charge-transfer absorption is what allows us to access the mild hydrocarbon oxidation path. 3.2.

Propylene to acrolein and propylene oxide Selectivity is an especially tough challenge in the case of autoxidation of small olefins in fluid phase because of diversion of the radical chain reaction by termination steps that result in the formation of oxy radicals. These are very reactive species that attack indiscriminately the starting hydrocarbon and primary products. The photochemical reaction in zeolites opens up much better controlled oxidation of low olefins, which are illustrated here briefly for the commercially important case of propylene oxidation. Irradiation of propylene and O2-1oaded zeolite BaY at room temperature with green or blue light induced partial oxidation of the olefin [18]. Readily identified products were acrolein, allyl hydroperoxide, and propylene oxide. The hydroperoxide was found to be stable when the zeolite was kept at -100~ Hence, photolysis experiments at this temperature allowed us to find out about the origin of the aldehyde and epoxide. Allyl hydroperoxide was the main product at -100~ the remaining 13% were propylene oxide. Warm-up of the zeolite after photo-accumulation of the hydroperoxide produced propylene oxide if excess propylene was kept in the matrix, but only acrolein if the olefin was removed prior to warm-up. Hence, allyl

1045 hydroperoxide is the primary photoproduct and acrolein originates from thermal dehydration of the hydroperoxide. Heterolytic thermal rearrangement to the corresponding carbonyl compound under elimination of H20 occurs in the case of all hydroperoxides with an t~ C-H group. Propylene oxide is produced by O transfer from allyl hydroperoxide to excess olefin (Scheme 2). The thermal rearrangement of the hydroperoxide to acrolein exhibits a steep temperature dependence while the epoxidation reaction does not. Therefore, the aldehyde is the preferred final oxidation product of the visible light-induced propylene oxidation at elevated zeolite temperature. For example, when conducting the propylene + 0 2 photochemistry at 55~ the acrolein to propylene oxide ratio is 2 to 1. Variation of the propylene loading level gives an additional handle on the acrolein/propylene oxide branching. Since the competition between the two thermal paths is influenced by the mobility of the propylene, which in turn depends on the structure of the zeolite, we are currently exploring the effect of zeolite structure on the aldehyde to epoxide branching. The most important result, however, is the unprecedented selectivity in terms of the allyl hydroperoxide intermediate (>98% at ambient temperature). The selectivity is undiminished even upon consumption of 20% of the propylene loaded into the zeolite. On the basis of spectroscopy on the visible propylene~ 2 chargetransfer absorption and the measured infrared product growth, a rather high reaction quantum yield of 20% was estimated. Photochemical Reaction CH3

C H - ' - CH2

+

02

h,v visible

CH2 ~

"~

Subsequent Thermal Reaction CH2

C H ~ CH2OOH

,//o

C H ~ CH2OOH

~-

CH2

CH~ C\

+

H20

\

H CH2

CH m C H 3 ~ ~ , , , , ~

/% CH2

CH

CH3

+

CH2

C H ~ CH2OH

Scheme 2. Origin of acrolein and propylene oxide upon photooxidation of propylene. Photooxidation experiments with other small olefins add strong support to the assignment of the visible absorption of hydrocarbon~ 2 complexes in zeolites as charge-transfer absorptions [19,20]. For example, the onset of the absorption tail (and the photolysis threshold) shifts linearly to higher energies with increasing ionization potential of the olefin. We have also found that, as expected, the optical absorption onset is sensitive to the magnitude of the electrostatic field. For example, the increase of the electrostatic field upon substitution of Na + by Ba 2+ is accompanied by a red shift of the 2-butene.O 2 absorption onset of 100 nm (4000 cm- 1) [ 14].

1046 3.3.

Selective oxidation of alkanes

Particularly interesting is the finding that small alkanes can be partially oxidized by oxygen using the same photochemical method in alkali and alkaline-earth zeolite Y [21-25]. Reactions established thus far are summarized in Scheme 3. Alkyl hydroperoxides and carbonyl compounds are produced with very high selectivity in each case even upon high (>50%) conversion of the hydrocarbon. These partially oxidized alkanes are of commercial importance, and the search for more selective and direct processes is intense [3,4]. H OOH

~

+O2

~<520nm~ ~ N a

0

Y

~

~H3 ~,<520 nm ~ BaY

CH3-- C--- OOH [

CH 3

CH 3

CH3\

CH3CH3 + 02

~,<500nm BaY

~<500 nm ~ CaY

+H20

~H3

CH3-- C ~ H + 0 2 ]

CH3CH2CH 3 + 02

~

C~ 3

/CHOOH

~

CH 3

CH3CH2OO H

/ C ~ O + H20 CH3

~

CH 3 ~ C \

+ H20 H

Scheme 3. Visible light-induced oxidation of small alkanes with complete selectivity. The first reaction we explored was oxidation of a tertiary alkane, isobutane. Since the ionization potential is only a few tenths of an eV higher than that of proplclene, we speculated that light in the range 400-500 nm would promote reaction with 0 2 in Baz + -exchanged zeolite Y. Blue light-irradiation of a BaY matrix loaded with isobutane and 0 2 gave t-butyl hydroperoxide with 98% selectivity [21], with the remainder due to trace amounts of acetone and methanol, the known thermal products of the hydroperoxide. Fig. 3 shows the infrared difference spectrum taken after photochemical conversion of 60% of the isobutane. All product bands originate from t-butyl hydroperoxide except the 1686 cml absorption which belongs to acetone. The negative bands show the depletion of isobutane. The reaction quantum efficiency was determined to be around 15%.

1047

0.3

0

cxa3

,o,

9 0.2

II II,-,

CH3--C--CH 3

0.1-

0 (/J

<

J

0.0-

-0.1 -

0.4

I

I

I

I

I

1700

1600

1500

1400

1300

|

9

900

800

-

tD u 0.2-

m

CH3 I CH3---C---OOH

0 ~ O.O,.Q

I

<

CH3 CH3__C__ H

-0.2I

I

I

I

I

I

3600

3400

3200

3000

2800

2600

v tern-1 ) Figure 3. Infrared difference spectrum upon 60% conversion of isobutane and 0 2 with blue light in zeolite BaY. An interesting aspect of the synthesis of t-butyl hydroperoxide in the zeolite is the opportunity of in situ use for olefin epoxidation [21 ].

1048 CH3 CH3

CH3 +

I

C----H

I

+

hv

02

~

CH3

CH3

O2-

I

CH3

CH3

CH3

I CH3

I

CmH

CmOOH

I

CH3

I ~

CH3

Ce

I

+

O2H

CH3

Scheme 4. Proposed mechanism for hydrocarbon photooxiation. The initial step of the proposed mechanism is proton transfer from the isobutane radical cation, formed by photoexcitation, to 02- (Scheme 4). Isobutane radical cation (and other alkane or alkene radical cations) are spectroscopically established transients [26]. Proton transfer from the radical cation to 02 - is expected to be very efficient since hydrocarbon radical cations are known to be highly acidic [27]. Efficient proton transfer quenching of the chargetransfer pair is probably the main reason for the rather high quantum yields to reaction of these hydrocarbon photooxidations (typically in the tens of percent) because it furnishes a path that is competitive with back electron transfer. The latter process prevents chemical reaction of charge-transfer states in many other systems. The alkyl and hydroperoxy radical so produced are expected to undergo cage recombination to yield the observed alkyl hydroperoxide. This mechanism, summarized in Scheme 4, is proposed to be the same for all visible light-driven alkane, alkene, and arene oxidations encountered thus far. The still higher ionization potential of propane (11.1 eV) and ethane (11.5 eV) compared to isobutane (10.6 eV) or cyclohexane (9.8 eV) prevented us from direct detection of the C3H8.O 2 and C2H6.O 2 charge-transfer absorption tail by diffuse reflectance spectroscopy. Nevertheless, judging from the observed trend of photooxidation yields versus ionization potential of the hydrocarbon, we could expect to observe propane and even ethane oxidation products in zeolites with doubly-charged cations such as BaY and CaY. Indeed, propane was found to be oxidized in zeolite BaY under irradiation with blue light when several hundred Torr of the alkane and one atmosphere of 0 2 were loaded into the pellet [23]. The sole final products were acetone and H20, while isopropyl hydroperoxide was observed as an intermediate. It dehydrates spontaneously to acetone in the ambient zeolite. Ethane photooxidation was very slow in BaY under irradiation with blue light, but was much accelerated in zeolite CaY. Acetaldehyde was the exclusive final oxidation product. No CO 2 was formed, which is unprecedented for ethane oxidation by 0 2. Build-up of a small steady state concentration of ethyl hydroperoxide intermediate was indicated by several small infrared bands [23]. These completely selective ethane and propane oxidations open up the possibility of using low alkanes, the constituents of natural gas, as new feedstocks for acetaldehyde and acetone in place of petroleum-derived ethylene and propylene currently in use. Moreover, these

1049 visible light-induced oxidations by 0 2 constitute a mild and very selective method for alkane CH bond activation.

4.

THERMAL ALKANE OXIDATION

Oxidation by 0 2 in cation-exchanged zeolite Y in the absence of light was observed for two alkanes, namely cyclohexane at T>50~ in NaY [22] and propane in BaY at T>50~ and in CaY at room temperature [23]. Final oxidation products were cyclohexanone and acetone, respectively, the same products that were generated in the visible light-driven reactions. Remarkably, the selectivity of these thermal oxidations was 100% even upon high (>30%) conversion of the alkane. A systematic study of cyclohexane autoxidation in alkali and alkaline-earth zeolite Y by diffuse reflectance infrared spectroscopy was recently reported by Jacobs and coworkers [28]. While still speculative, our observations point to a charge-transfer mechanism involving cation sites with especially strong electrostatic fields (there exists a heterogeneity in terms of cation shielding in zeolites) [ 11]. It is the high mobility of small alkane and oxygen molecules in the zeolite that allows easy access to such minority sites [6]. At these poorly shielded cation sites, spontaneous charge-transfer could result in alkane radical cation and 0 2- formation. We propose the same mechanism for the subsequent reaction leading to alkyl hydroperoxide and ketone as proposed for the light-induced oxidation. At elevated temperature, the familiar chain propagation of autoxidation may play a role because sufficient energy is available to overcome the activation energy for H abstraction from the alkane [ 1,2]. We attribute the lack of similar thermal reactivity of unsaturated hydrocarbons such as toluene or propylene to the fact that these molecules are fairly strongly attracted by the cations, especially the most poorly shielded ones. The diminished electrostatic fields may not be large enough to promote thermal hydrocarbon-oxygen charge-transfer in these cases.

5.

CONCLUSIONS

Partial oxidation of small alkenes, alkanes, and alkyl substituted aromatics by 0 2 to organic building blocks and industrial intermediates has been achieved in alkali or alkaline-earth zeolite Y and L at unprecedented selectivity. All reactions are induced by visible light, some even under dark thermal conditions. One key factor responsible for the tight control of these reactions is the very strong stabilization of the hydrocarbon.O 2 charge-transfer state by the electrostatic field of the zeolite cage. This makes them accessible by low-energy visible photons or even by thermal excitation at modest temperatures. Hence, the primary products, hydrocarbon radical and HOO radical, are generated with minimal excess energy, which reduces their mobility and likelihood of random coupling reactions that would destroy the selectivity. For the same reasons, homolytic fragmentation of the alkyl hydroperoxide intermediates, either thermally or by secondary photolysis does not occur. No overoxidation has thus far been observed, even at high (>50%) conversion of the hydrocarbon. Studies by time-resolved FT-infrared spectroscopy are in progress to establish more firmly our proposed mechanism for these hydrocarbon oxidations in zeolites. In order to scale up these oxidations from the micromolar scale employed so far to practical quantities, the most important issue to be addressed is the optimization of the process conditions with regard to desorption of the products from the zeolite.

1050 ACKNOWLEDGMENTS

This work was supported by the Director, Office of Energy Research, Office of Basic Energy Sciences, Chemical Sciences Division of the U. S. Department of Energy under Contract No. DE-AC03-76SF00098. I am indebted to my postdoctoral collaborators Fritz Blatter, Hai Sun, and Sergey Vasenkov for their hard work and enthusiasm. REFERENCES

1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25.

26. 27. 28.

R.A. Sheldon and J. K. Kochi, Metal-catalyzed Oxidation of Organic Compounds, Academic Press, New York, 1981. G.W. Parshall and S. D. Ittel, Homogeneous Catalysis, 2nd ed., Wiley, New York, 1992. J.E. Lyons and G. W. Parshall, Catal. Today, 22 (1994) 313. C.B. Dartt and M. E. Davis, Ind. Eng. Chem. Res., 33 (1994) 2887. D.W. Breck, Zeolite Molecular Sieves; Structure, Chemistry and Use, Wiley, New York, 1974. J. K~irger and D. M. Ruthven, Diffusion in Zeolites, Wiley, New York, 1992, Chapt. 13. F. Blatter and H. Frei, submitted. B. Barrachin and E. Cohen de Lara, J. Chem. Soc., Farad. Trans. 2, 82 (1986) 1953. J . A . R . Coope, C. L. Gardner, C. A. McDowell and A. I. Pelman, Mol. Phys., 21 (1971) 1043. M.A. Spackman and H. P. Weber, J. Phys. Chem., 92 (1988) 794. E. Preuss, G. Linden and M. Peuckert, J. Phys. Chem., 89 (1985) 2955. J.A. Rabo, Catal. Rev., 23 (1981) 293. H. Sun, F. Blatter and H. Frei, J. Am. Chem. Soc., 116 (1994) 7951. F. Blatter, F. Moreau and H. Frei, J. Phys. Chem., 98 (1994) 13403. S. Vasenkov and H. Frei, J. Phys. Chem., submitted. D.F. Evans, J. Chem. Soc., (1953) 345. H. Tsubomura and R. S. Mulliken, J. Am. Chem. Soc., 82 (1960) 5966. F. B latter, H. Sun and H. Frei, Catal. Lett., 35 (1995) 1. F. B latter and H. Frei, J. Am. Chem. Soc., 115 (1993) 7501. F. Blatter and H. Frei, J. Am. Chem. Soc., 116 (1994) 1812. F. Blatter, H. Sun and H. Frei, Chem. Eur. J., 2 (1996) 385; Angew. Chem., Int. Ed. Engl., 35 (1996). H. Sun, F. Blatter and H. Frei, J. Am. Chem. Soc., 118 (1996) 6873. H. Sun, F. Blatter and H. Frei, Catal. Lett., in press. H. Frei, CHEMTECH, 26 (1996) 24. H. Frei, In: Heterogeneous Hydrocarbon Oxidation, T. Oyama and B. K. Warren (eds.), ACS Symposium Series No. 638, American Chemical Society, Washington, D. C., 1996, p. 409. M. Iwasaki, K. Toriyama and K. Nunome, J. Am. Chem. Soc., 103 (1981) 3591. O. Hammerich and V. D. Parker, Adv. Phys. Org. Chem., 20 (1984) 55. D.L. Vanoppen, D. E. De Vos and P. A. Jacobs, Proceedings of the 1 lth International Zeolite Conference, Seoul, Korea, 1996.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997Elsevier Science B.V. All rights reserved.

1051

Catalytic oxidations by in situ generated peroxotungsten complexes immobilized on layered double hydroxides (LDH): Relation between catalytic properties and peroxotungstate micro-environment B.F. Sels, D.E. De Vos and P.A. Jacobs* Centre for Surface Chemistry and Catalysis, Katholieke Universiteit Leuven, Kardinaal Mercierlaan 92, 3001 Heverlee (Belgium)

ABSTRACT Tungstate is exchanged on a C1--LDH and converted into peroxo complexes by the action of H202. This hydrophilic material is an efficient catalyst for the truly heterogeneous epoxidation of allylic alcohols. The alkaline nature of the LDH suppresses hydrolysis of the epoxides. For simple olefins such as cyclohexene, the use of a lipophilic LDH, exchanged with WO42- and an organic anion, is more appropriate. In these reactions, the apolar surface provides a suitable micro-environment for epoxide formation with over 90 % selectivity. Finally it is demonstrated that a co-anion such as SO4 ~- favorably influences the selective oxidation by peroxoW, probably through formation of heteronuclear S-W peroxoclusters. 1. I N T R O D U C T I O N The interest in hydrogen peroxide as primary oxidant has grown steadily over the last decades, because of its low cost and for environmental reasons. Unfortunately, hydrogen peroxide itself has a restricted oxidizing potential and is therefore unadequate to oxidize e.g. olefins and amines in appreciable amounts. Epoxidations and amine oxidations can only be achieved after activation of the hydrogen peroxide by means of a suitable catalyst [1]. Tungsten(VI) compounds are superior amongst other transition metals of groups 5B and 6B in the activation of H202 [2-8]. In these reactions, reactive peroxotungsten derivatives are formed, which act as highly efficient monooxygen donor species. Numerous investigations are directed towards the synthesis of novel peroxo tungsten complexes [9-12]. In general, these are assembled starting from WOj- and S042, RSO3, PO43-, RPO32- or AsO43. Single cristal X-ray analyses, 31p. and 183W-NMR, Raman and IR spectroscopy are helpful tools in characterizing the peroxo species. The oxidation ability of the This work is supported by the Belgian Government as an I.U.A.P. project. We thank I.W.T. (BS) and F.W.O. (DDV) for research positions. E-mail: [email protected]

1052 complexes has been demonstrated in both stoichiometric and catalytic epoxidation of olefins, oxidation of amines and oxidation of alcohols, mainly in the Ishii-Venturello biphasic conditions. Notwithstanding the superiority of the tungsten catalysts, only few publications deal with the immobilization of these complexes on support materials [13,14]. Therefore, our research objective is the development of a truly heterogeneous W catalyst, which can successfully oxidize both watersoluble and insoluble nucleophilic substrates. Recently, we explored the potential of layered double hydroxides (LDHs) as supports for active tungsten catalysts in the epoxidation of olefins, such as allylic, homoallylic alcohols and non-functionalized olefins [15]. Although LDHs are mostly used as precursors to basic oxides, profit can also be taken from their excellent anion exchange properties [16]. The anionic character of tungsten complexes in aqueous peroxidic conditions makes LDHs obvious candidates for heterogenization attempts. Moreover, the highly variable chemical nature of both the structure-building octahedral layers and the anion population allow to adapt the catalytic center's environment to widely changing reaction requirements. In the present paper, the main objectives are (i) to prepare reactive peroxocomplexes in situ at the material's surface starting from a precursor material, and (ii) to control the catalytic properties (activity, selectivity, oxidant efficiency) via modification of the micro-environment of the catalytic center through variation of the anion population. The catalyst precursors and the in situ formed peroxocomplexes are characterized by means of XRD, IR, TGA/DTA and UV-Vis reflectance spectroscopy. 2. EXPERIMENTAL

2.1. Apparatus The powder pattern of the samples was obtained with a Siemens D5000 matic X-ray diffractometer using filtered Cu-I~ radiation, operating at 40 kV and 50 mA. Samples were scanned from 3 ~ to 65 ~ 20. The FTIR absorption spectra of the samples as KBr pellets were recorded between 400-1300 cm 1 with a Nicolet 730 spectrometer. UV-Vis Reflectance spectra were recorded on a Cary 5 spectrophotometer, using anhydrous BaSO4 as the standard over the whole spectral range. The Mg, A1, C1 and W contents were determined by emission spectroscopy with an inductively coupled plasma. The DTA and TGA diagrams were recorded in OJHe (1/3) for 25-35 mg of air-dried samples, with a heating rate of 5~ from 20 to 800~

2.2. Preparation of Catalysts The host compound IVIg/A1-LDH-[CI] was obtained by the hydrolysis of MgC12 and A1CL in weakly alkaline solutions (pH=9-10, NaOH) under nitrogen atmosphere at room temperature [17]. All water was deionized before use and was decarbonated by boiling for at least two hours. The LDH was

1053 isolated by centrifugation and dried by lyophilization. Elemental analysis gave the molar ratios expected for the Mg233A1derivative, [Mgo 7Alo.3(OH)2][C1-]. The catalyst precursors (e.g., LDH-[WO42, CI] (A), LDH-[ WO42-, SO42] (B) and LDH-[WO42-, p-toluenesulphonate] (C)) were prepared by ion-exchange reaction of [Mgo 7Alo~(OH)2][C1-] in aqueous suspension (pH=9) with tungstate and an appropriate co-anion, e.g. sulfates and sulphonates. Syntheses were performed so that tungstate takes 10-15% of the anion exchange capacity; the rest is t ~ e d with the co-anion. The precursors A, B and C were isolated routinely by centrifugation and lyophflization.

2.3. Epoxidation procedure 0.06 g of component A (or B) or 0.075 g of component C, corresponding to 0.01 mmol W, is suspended in 3 ml MeOH, together with 3 mmol substrate. Cold aqueous 35% H202 (7.5 mmol, diluted with MeOH in a 1:1 volume ratio) is added stepwise to the suspension in five equal portions. Reactions are performed at 298 K under efficient stirring. Reaction mixtures were analyzed on a GC equipped with a 50 m CP-Sil 5 column. Product identity was checked by GCMS. 3. CATALYST CHARACTERIZATION

~

The purity of the Mg/AI-[CI] LDH phase and the basal spacings were checked by X-ray A diffraction. The thickness of the intermediate layer (3.01 A) was calculated by substracting B the thickness of the brucite layer (4.77/~) from the thickness of the unit layer (doo3 - 7.78 A) and agrees well with previously published C values [ 18]. The X-ray patterns for the catalyst 0 20 40 60 80 precursors A, B and C are given in Figure 1A, B and C. The basal spacings (003) appear at 20 F i g u r e 1. X-ray patterns = 11.387 ( d A - 7.764 A), 20 - 10.606. (rib = for the precursor 8.335 A) and 20 - 5.021 ( d c - 17.586 /~). compounds A, B, and C. Because of the predominance of the co-anion (85-90% of the AEC), it is clear that the co-anion (CI, SO42 or p-tosylate) determines the height of the interlayer gallery. Figure 2 shows the DT/TG profiles for the host compound and for the organo-LDH C, measured between room temperature and 800 ~ Endothermic peaks due to loss of physisorbed and interlayer water are observed around 105 ~ for the host compound, and at 105 and 210 ~ for compound C. The total amount of desorbed water is clearly less for the organo-LDH, in agreement with its hydrophobic nature. The dehydroxylation of the brucite-like layer in the host compound results in a double endothermic peak near 360 ~ For compound C, this structural feature is largely masked by the highly

1054

exothermic combustion of the p-tosylate at 540 ~ intercalation of the organic anlon.

This confirms the successful

EXO

5 wt.%

o

45o

T(oC)

E)

s5oo

T(~

soo

F i g u r e 2. Weight loss (%) and heat flow for the host compound (left) and component C (right). Figure 3 shows the IR spectra (400-1300 cm 1) for the starting LDH and for the catalyst precursors A, B and C. The characteristic frequencies of the exchanged anions are listed in Table 1. The main W-O stretching frequency is indicative of the coordination number and the nuclearity of W. Polynuclear, octahedrally coordinated W typically has bands in the 900-1000 cm 1 range, while the v3 stretching of tetrahedral WO42- absorbs at 838 cm -1 [19]. In materials A and B, a broad shoulder between 900 and 800 cm -1 is superimposed on the LDH spectrum. It seems therefore t h a t the tungstate ions do not form % oligomers on the LDH surface. For T the sulfate in compound B, the two IR active vibrations v3 and v4 at A 1110 and 602 cm 1 are quite similar to those observed for the free ion [20]. The v3 stretching vibration is broadened; this splitting results from a distortion of the t e t r a h e d r a l lObO 500 geometry because of weak Wavenumbers (cm-1) bidentate bonding of the anion with F i g u r e 3. IR spectra for the host the interlayer surface. The compound (upper curve) and the sulphonates in compound C give catalyst precursors A, B and C. The characteristic very strong Vs-o dotted line is at 830 cm 1. absorptions in the 1300-1000 cm -1 range. Spectrum C displays another strong and sharp band at 820 cm "1. Even if a W-O absorption is expected in this region, an assignment of this band to the C-H out-of-plane vibrations of the aromatic ring is more plausible.

1055 T a b l e 1. Characteristic frequencies (cm 1) of anions on LDH materials.

vs-o

vw-o

other assignments

A

-

833

B

1110

830

602 (O-S-O bending)

C

1195, 1130,1050, 1025

*

820 (C-H out-of-plane)

9hidden by the C-H out of plane vibrations The catalyst precursors can be transformed into active 'O' transferring materials by addition of HzO2, as proved by UV-Vis reflectance. For instance, a suspension of LDH-[WO42-, CI] in methanolic H202 displays two clearly distinguishable absorption bands at ~=238 nm and )~=325 nm, and visually the sample turns slightly yellow [21]. These frequencies are close to the literature data for the charge transfer bands (O22-~W) in di- and tetraperoxotungstate anions respectively.

4. CATALYTIC E P O X I D A T I O N S The in situ generated peroxocomplexes were tested for the catalytic epoxidation of various olefins, such as allylic alcohols, homoallylic alcohols and non-functionalized olefins. The results of these H202 oxidations in an alcoholwater system are summarized in Table 2 for the hydrophilic catalyst A, and in Table 3 for the lipophilic material C. Especially for the more reactive alkenes, the turnover number comes close to the maximum of 300. The epoxide selectivity generally exceeds 90%, with minimal solvolysis. With catalyst A, some substrates gave a lower selectivity. For instance, the product distribution for cyclohexene is 65% epoxide, 27% of allylic oxidation products and only 4% of the diol. The epoxycyclohexane selectivity increases to 91% with the hydrophobic material C. Within the class of hydrophilic LDHs, tungstate can be surrounded by different co-anions and these might affect its catalytic properties. As the comparison of A and B in Figure 4 shows, substitution of CI with SO42 increases the epoxide selectivity e.g. from 65 to 85% for cyclohexene.

1056 T a b l e 2. E p o x i d a t i o n of simple olefins, allylic a n d homoallylic alcohols w i t h a q u e o u s H202 in t h e p r e s e n c e of t h e hydrophilic LDH-[WO42-, el-] a

Substrate

Sepoxides (%)

Xsubstrate (%) EHOOH (%)~

TON d

allyl alcohol

91

15

10

38

methallyl alcohol

92

40

22

104

crotyl alcohol

100

73

39

218

trans)

61

32

179

cis)

54

31

162

trans cis

2-hexen-l-ol

2-hexen-l-ol

98 (all

100 (all

1,4-but-2-endiol

85

52

33

156

geraniol

92

90

74

242

nerol

85

89

68

222

perillyl alcohol

78

63

40

154

3-buten-l-ol

100

13

8

39

3-methyl-3-buten-l-ol

100

33

17

99

3-MeOH-cyclohexene

86

33

17

84

cyclopentene

63

72

42

151

cyclohexene

65

44

26

93

3-methylcyclohexene

28

49

40

52

2, 3-dimethyl-2-butene

39

79

58

105

cyclooctene e

88

47

22

130

18

11

35

54

26

151

trans

2-hexene ~

norbornene

60 (all

trans)

93

a S = selectivity, X = (retool substrate converted)/(mmol substrate initially present) and E = efficiency of oxidant use, measured as (meq "O" in oxidation products / mmol H202 consumed), b Control experiments prove the truly heterogeneous character of the catalysis. Maximally 0.3% of the total W was found in solution, which is by far insufficient to account for the observed activity, c H202 was analyzed via cerimetry, a TON = mmol (epoxide + solvolysis products) /mmol W. Maximum TON value is 300. e in 4.5 ml MeOH.

1057 T a b l e 3. E p o x i d a t i o n of allylic alcohols a n d non-functionalized olefins with a q u e o u s HeO2 in the presence of LDH-[p-toluenesulphonate-, WO42]. a

Sepoxides (%)

Substrate

Xsubstrate

EHOOH(%)

TON b

crotyl alcohol

99

23

25

69

fferaniol

99

98

84

294

nerol

99

94

74

282

perillyl alcohol

91

80

56

243

3-Me O H - c y e l o h e x e n e

96

38

39

113

cyclohexene

91

46

46

130

2, 3 - d i m e th y l - 2 - b u t e n e

79

84

80

201

cyclooctene c

1O0

64

51

192

95 (all trans)

51

40

148

trans 2-hexene c

a The catalytic activity of a tungstate-free p-toluenesulphonate exchanged hydrotalcite was too little to be detected within reasonable time. u as in Table 1. ~in 4.5 ml MeOH.

~A compound A D compound B 85

Sel.(%) 10c

2 F i g u r e 4. Epoxide selectivities for [WO42",C1] (n) VS. [WO42, SO42"] (B) catalysts.

5. D I S C U S S I O N The physicochemical c h a r a c t e r i z a t i o n gives a r e h a b l e idea on the s t a t u s of the c a t a l y s t p r e c u r s o r s before peroxide addition. XRD a n d TGA prove t h a t the n o r m a l l y hydrophilic m a t e r i a l s are t r a n s f o r m e d into a hydrophobic p r e c u r s o r by the successful i n t e r c a l a t i o n of the p-tosylate. IR suggests t h a t at the surface of a Mg/A1-LDH, W is largely p r e s e n t as monomeric, t e t r a h e d r a l t u n g s t a t e .

1058 Reaction of this WO42 with H~Oz produces peroxocomplexes, which in an aqueous methanolic medium epoxidize allylic alcohols. The reactivity of our system agrees well with that of tungstate salts, dissolved in a single polar liquid phase. The alkaline nature of the LDH support seems however to prevent solvolysis reactions. In the epoxidation of (homo)allylic alcohols, selectivities are therefore better with the WO42--LDH A than with the homogeneous W salts [2,3]. However, for some of the simple olefins, allylic oxidation is not negligible. This allylic oxidation can be suppressed by working with the lipophilic catalyst C. Hydrophobic olefins such as cyclooctene are expected to interact favorably with the aromatic p-tosylate on the peroxoW-LDH surface. The transfer of the electrophilic oxygen to the olefin then takes place in an essentially hydrophobic compartment of the reaction system, as in the Venturello-Ishii reactions. In such an environment the olefin is enriched with respect to hydrogen peroxide and this improves the efficiency of the oxidant use. Another potential advantage of our system C is that there is no need for a chlorinated solvent to constitute a second, apolar liquid phase. The relation between catalytic performance and polarity of substrate and catalyst becomes obvious when the yields with catalysts A and C are compared. In Figure 5, AYleld is defined as Yield with C minus Yield with A. Epoxide yields are largest when the polarities of substrate and catalyst are matched.

A Yield

HOCH HOCH~

+

28 [~:."i':] ~ J

16

I

23 :r:_.:.:

~.......~.

~2 OH

35

50 .

.

.

.

.

.

.

.

.

.

.

.

.

.

.

Hydrophilic

.

.

.

.

.

.

.

9

. .

9

9

.

.

.

.

.

.

.

9

.

.

.

.

.

.

.

~ , .

Hydrophobic

F i g u r e 5. Effect of catalyst polarity on epoxide yield. While the p-tosylate undoubtedly increases the hydrophobicity of the W micro-environment, formation of heteronuclear S-W peroxoclusters might also contribute to the improved epoxide selectivities in Table 3. In order to check this hypothesis, S042- and CI exchanged LDHs were compared. Even if these materials have a similar polarity, selectivities are systematically better for the

1059

SO42- catalyst (see Figure 4). Thus the LDH seems to organize formation of oligomers by concentrating the anions at the surface. In conclusion, this report introduces a truly heterogeneous catalyst based on the in situ generation of active peroxotungstate complexes. By taking profit of the variability of the anion population, a new class of tailor-made catalysts is available. The hydrophobicity of the surface can be tuned, and heteronuclear anion oligomers can be synthesized at the sohd-liquid interphase.

REFERENCES

1. R.A. Sheldon and J.K. Kochi, Metal-Catalyzed Oxidation of Organic Compounds, Academic Press, New York, (1981). 2. D. Prat and R. Lett, Tetrahedron Lett. 27 (1986) 707. 3. D. Prat, B. Delpech and R. Lett, Tetrahedron Lett. 27 (1986) 711. 4. J. Prandi, H.B. Kagan and H. Mimoun, H. Tetrahedron Lett. 27 (1986) 2617. 5. C. Venturello, R. D'Aloisio, J.C.J. Bart and M.J. Ricci, J. Mol. Catal. 32 (1985) 107. 6. M. Quenard, V. Bonmarin and G. Gelbard, New J. Chem. 13 (1989) 183. 7. D.C. Duncan, R.C. Chambers, E. Hecht and C.J. Hill, J. Am. Chem. Soc. 117 (1995) 681. 8. Y. Ishii, K. Yamawaki, T. Ura, H. Yamada, T. Yoshida and M. Ogawa, J. Org. Chem. 53 (1988) 3587. 9. L. Salles, C. Aubry, R. Thouvenot, F. Robert, C. Dor6mieux-Morin, G. Chottard, H. Ledon, Y. Jeannin and J.-M. Br6geault, Inorg. Chem. 33 (1994) 871. 10. L. Salles, F. Robert, VI Semmer, Y Jeannin and J.-M. Br6geault, Bull. Soc. Chim. Fr. 133 (1996) 319. 11. W.P. Griffith, B.C. Parkin, A.J.P. White and D.J. Williams, J. Chem. Soc., Chem. Commun. (1995) 2183. 12. W.P. Griffith, B.C. Parkin, A.J.P. White and D.J. Williams, J. Chem. Soc. Dalton Trans. (1995) 3131. 13. R. Neumann and H.J. Miller, J. Chem. Soc., Chem. Commun. (1995) 2277. 14. E. Duprey, J. Maqut, P.P. Man, J.-M. Manoh, M. Delamar and J.-M. Br6geault, Applied Catalysis A: General 128 (1995) 89. 15. B.F. Sels, D.E. De Vos and P.A. Jacobs, Tetrahedron Lett. 37 (1996) 8557 16. C. Cativiela, F. Figueras, J.M. Fraile, J.I. Garcia and J.A. Mayoral, Tetrahedron Lett. 36 (1995) 4125. 17. S. Miyata, Clays Clay Miner. 23 (1975) 369. 18. F. Cavani, F. Trifiro andA. Vaccari, A. Catal. Today 11 (1991) 173. 19. K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination compounds, Wiley-Interscience, New York (1986). 20. S. Miyata and A. Okada, Clays and clay minerals 25 (1977) 14. 21. J.A. Conner and E.A.V. Ebsworth, Adv. Chem. Radiochem. 6 (1964) 279.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1061

Homobimetallic Heptadentate coordinated Iron Complexes in Montmorillonites as Methane MonoOxygenase Mimics P.P. Knops-Gerrits

, S.Dick w A.Weiss w M.Genet § P.Rouxhet § X.Y.Li ~ P.A.Jacobs §

+Centrum voor Oppervlaktechemie en Katalyse, KU Leuven, Kardinaal Mercierlaan 92, B-3001 Heverlee, Belgium; w Institut mr Anorganische Chemie der Universit/it Mianchen, Meiserstrasse 1, 80333 Miinchen, Germany, +Unit6 de Chimie des Interfaces, UC Louvain, Louvain-la-Neuve, Belgium; ~ Department, Hong Kong University for Science and Technology, Clear Water Bay, Kowloon. * To whom correspondence should be addressed.

1. INTRODUCTION Many metallo-proteins that catalyse the incorporation of oxygen atoms derived from dioxygen into organic substrates usually contain iron in mono- or polynuclear form. In case of iron-containing non-heme-oxygenases such as methane monooxygenase (MMO) [1-3] our understanding is more limited as only a few suitable models are known, compared to the ironcontaining heme-oxygenases, where knowledge has been substantially advanced. A good example of the latter is the extensive characterisation ofhigh-valent iron-oxo porphyrins [4]. The binuclear iron(III) complex [Fe2(HPTP)(p-OH)(NO3)2](C104) 2 (1) with HPTP = N,N,N',N',-tetrakis(2-pyridyl methyl)-2-hydroxy-l,3-diamino-propane [5], reacts with H2Oz to yield a blue species [6]. The crystal structure of (1) and its interaction with montmorillonites (MMT) has been reported [7]. The occlusion of [Fe2(HPTP)(p-OH)] 4+ complexes in new families of hexagonal mesoporous materials such as MCM-41 and HMS will be described elsewhere [8]. Related HPTB complexes, with benzimidazole instead of pyridine, have their tripodal N atoms trans to strong basic peroxo-ligands, the benzimidazole ligands perpendicular to the Fe203 plane and trans to each other on every iron [9]. 3+ Oxygen binding in [Fe2(p-I,2-O2)(N-Et-HPTB)(L)] with L =(p-I,2-O,O'-O2C-C6Hs) or (Ph3PO)2 is stable for months at -40~ [11-12]. Knowledge about the way oxygen coordinates is important to understand how MMO converts methane to methanol [14]. A shuttling between the diferrous and diferric states of natural MMO allows variation between weak ferromagnetic and antiferromagnetic coupling [3, 9]. For partial oxidation of alkanes with iron complexes, examples can be found in radical Fenton chemistry as well as in iron oxo chemistry [4, 12]. Other MMO mimicks include [Fe2(~t-O)2(5-Me-TPA)](C104)2, (TPA = tris(2-pyridyl methyl)amine) [15] and [Fe2(H2Hbab)2 (N-Me-Im)2 ] (H2Hbab = 1,2-bis(2-hydroxobenzamido)benzene [16]. Interesting mimics of hemerythrin that bind 02 reversibly have related structures [ 19]. In the present work the spectroscopic and catalytic fate of [Fe2(HPTP)(~t-OH)(NO3)2] (C104) 2 complexes are compared with their analogs heterogenised in MMT. Mrssbauer, XRD, ESCA, EPR, DRS, FT-IR & Raman, TGA and SEM characterisation are used to elucidate the nature of the active site prior to and under conditions operative in oxidation catalysis.

1062

2. EXPERIMENTAL 2.1. Materials 1,3-diamino-2-hydroxy-propane (95%, ALDRICH), para-chloropicoline (99%, ALDRICH), dichloromethane (99%, ACROS), cyclohexane (99%, ACROS), cyclohexanol (99%, ACROS), adamantane (99%, ACROS), hydrogen peroxide (30% solution in water, ACROS) and tBHP (tbutyl-hydroperoxide, 70% solution in di-tert-butyl-peroxide, ACROS) are used.

2.2. Synthesis

The N,N,N"N',-tetrakis(2-pyridyl-methyl)-2-hydroxy-1,3-diamino-propane was prepared by crushing para-chloropicoline (0.097 mol) and mixing it with 2-hydroxy-l,3diamino-propane (0.016 mmol) [5-6]. The mixture is heated to 170-180~ for 1 hour, till gas formation stops. After cooling the red glassy structure is diluted with HC1 (150 ml, 4M), thus a blue precipitate arises. After filtration the precipitate is washed several times with acetone. The precipitate is then dissolved in water and neutralised with a diluted ammonia solution. The white precipitate is recrystallised in acetone and crushed to a fine powder and then dried under vacuum. To an ethanolic solution with Fe(NO3)3.6H20 (0.31 g) the HPTP (0.3 g) is added. The precipitate ofbinuclear iron complex is collected. Montmorillonite, a bentonite sample from Linden/Bavaria, was used for the introduction of the binuclear complexes by cation exchange as previously reported [7]. The complexes can be removed from the interlayers only by highly concentrated solutions of NaC104 or KC104. As inferred from TGA when the clay sample is heated in an Ar/02 atmosphere, the Fe coordinated nitrate is removed between 160 and 260~ and the ligand is cracked between 350 and 450~

2.3. Spectroscopy The ESCA surface-concentration is performed on a SSX 100/206 spectrometer (Xprobe) from FISONS, with A1Kot monochromatised X-ray source (hv = 1486.6 eV). The spot size is 600 ~tm corresponding to an area of 0.493 rnlTl 2. The Fe203 standard shows the 2pl/2 and 2p3/2 lines at 724.3 and 710.7 eV, respectively. The values obtained for the organic complexes within the zeolites are in close proximity of this of the K3Fe(CN)6 and K4Fe(CN)6 complexes with values for the 2pl/2 and 2p3/2 lines at 710.2 eV and 708.3 eV. Use of ESCA for the characterisation of clays is difficult as for the non-conducting MMT positive charge accumulation occurs and binding energy referencing becomes problematic, either gold decoration or flood gun static charge neutralisation was used [ 18]. To correct for these effects of surface charge, the Si2p line of 103.5 eV for SiO2 was used as a reference, rather than the Cls at 285 eV of adventituous carbon, the Ols at 530.0 eV for metallic oxides and at 534.6 eV for metalloid oxides. For the determination of atom concentrations in MMT, the Auger line of Mg at 306 eV is used since the 51 eV band for Mgzp in oxides coincides with the 55 eV band of Fe3p, the 90 eV band for Mg2~ in oxides coincides with the 93 eV band of Fe3~. FT-IR spectra were recorded on a Nicolet F-730 spectrometer, FT-Raman spectra on a Bruker IFS 100 spectrometer. The sample was irradiated by a CW Nd:Yag laser (fundamental at 1064 nm) with a 50-500 mW power in a 180 ~ scattering geometry, and the signal was detected with a Ge detector cooled with liquid N2, accumulated 500 times to obtain a resolution of 4 cm -~. An EPR Bruker ESP 300E spectrometer is operated in the X band mode at 9.58 GHz. The experiments were performed at 300 K, using a modulation frequency and amplitude of 100 kHz and 10 G, respectively, a time constant 1.28 ms, a resolution of the field axis of 2048, a sweep time of 83.89 s and microwave power of 6.33 mW. The work at 95 GHz A W-band EPR spectrometer with two superconducting magnets was constructed at the IERC (Illinois

1063 EPR Research Center). Zero-field splittings in metal based systems can be overcome at these high microwave frequencies ( sensitivity for 108 spins). X-Ray Diffraction (XRD) patterns were recorded on a Philips HTK-KC diffractometer with a CuKc~ X-ray source, linked to Philips 386 computer. Thermo-Gravimetry (TG), was recorded on a Setaram TG-DTA 92 thermobalance. Diffuse Reflectance Spectra (DRS) were recorded on a Cary-5 spectrofotometer with a BaSO4 integration-sphere in the UV-VIS-NIR region. M6ssbauer spectra are recorded on a vertical constant acceleration drive in transmission geometry with a 28mCi 57Co(Rh) source. Isomer shift data are expressed relative to metallic Fe at 293 K which has an isomer shift of ~5- -0.0888 mm/s relative to natural c~-Fe.

3. RESULTS AND DISCUSSION The characterisation of binuclear heptapodate coordinated iron(III)-complexes of HPTP, i.e. [Fez(HPTP)(~t-OH)(NO3)2](C104) 2 and their heterogeneous [Fe2(HPTP)(OH)(NO3)z]Z+-MMT analogs (Fig.1.) will be performed with a wide number of techniques. Apart from XRD data and exchange isotherms and XRD characterisation of the interaction with the MMT such as the are known to date [7].

H Figure 1.a Structure of [Fez(HPTP)(g-OH)(NO3)2] 2+ ( the nitrates are omitted for clarity ).

1.b. Stick and ball representation of the montmorillonite unit cell.

3.1. Structure and Distribution of Dinuclear Complexes in Clays

From ~ it can be seen that the basal spacing of montmonllonite increases from 0.96 nm to 1.84 nm upon adsorption of the [Fea(HPTP)(~-OH)(NO3)2 ]2+ complexes. Basal spacings between clay sheets vary as a iiiiii i /84 ii , iiiii i i,.... . . ./'iii . Jii result of swelling. In ethanol, acetone and acetonitrile swelling increases the d-spacing of the clay to 1.69, 1.85 and 1.96 nm, respectively. In Fig 2. the scanning electron micrographs of montmorillonite are shown, containing [Fez(HPTP)(g-OH) (N03)2]2+. The montmorillonite sheet structure is retained upon inclusion of the metal complexes. Furthermore the image of the external surface Figure 2. Scanning Electron Micrograph of remains identical prior to and atter occlusion [Fe2(HPTP)(OH)(NO3)2 ]2+-MMT. of the complexes. In ESCA the pure [Fe2(HPTP)(~-OH)(NO3)2](C104) 2 shows a 2p3/2 line with an asymmetric form which can be decomposed as 710.6 eV (1.42 %), 714.16 eV (0.29 %) and

1064 717.73 eV (0.39 %). The asymmetry of the high energy peak is generally observed for polynuclear iron complexes. The concentrations of the [Fe2(HPTP)(~t-OH)(NO3)2] 2+ complex occluded in montmorillonite, expressed as the Fe / (Si +A1 +Mg) ratio amounts to 6 %. Table 1. ESCA concentrations and binding energies (BE) of the [Fe2(HPTP)(~t-OH)(NO3)2] (C104) 2 salts and the montmorillonite occluded [Fe2(HPTP)(I.t-OH)(NO3)2] 2+ complexes. Sample Fe2p3 Cls I Cls z 01 s ~ Ols z N1s ~ NlsZ C12p3/2~ C12p3/2"~ HPTP/Fe 2.10 57.69 1.64 0.50 22.32 9.15 1 . 4 3 0.73 4.45 BE(eV) 710.6 285.0 292.2 529.1 531.7 399.5 406.5 198.1 207.1 MMT/HPTP/Fe 1.25 22.57 1.39 44.36 3.67 BE(eV) 710.2,712.1 285.6 292.6 531.6 400.1 Upon heterogenisation of the iron complex on MMT the charge compensating chlorates and nitrates are below the XPS detection sensitivity. In this respect the binding energy of the nitrate N atoms (406.5 eV) varies significantly from that of the N atoms of the pyridine type ligands (399.51 eV). The removal of the chlorates and (part of the) nitrates has implications on the complex structure. Apart from the ~t-OH interactions of iron, the lattice anions may now compensate for the loss in these charge compensating interactions. Table 2. ESCA ratios, concentrations and binding energies [Fe2(HPTP)(~-OH)(NO3)2]2+-MMT. Fe/(Si +A1 +Mg) Fe2p 3 Ratio or Cone. (%) 0.058 1.25 BE (eV) 710.2, 712.1

(BE) ofFe, Si, A1 and Mg for

,,

Si (2p) 16.32 102.50

A1 (2p) 5.39 74.46

Mg (Auger) 2.68 305.59

3.2. Vibrational Spectroscopy For pure MMT in FT-IR several intemal structure insensitive vibrations of T O 4 tetraedra are seen. Asymmetric and symmetric stretch vibrations at 1044cm -1 and 720-650 -1 -1 cm and vibrations of the T-O bond at 524 and 466 cm are seen. External structure sensitive vibrations between individual TO 4 tetraedra occurs as asymmetric stretch vibrations at 918 -1 -1 cm and symmetric stretch vibrations at 798 cm are seen. FT-IR and FT-Raman vibration modes for linear or bent Fe-O-Fe (180 ~ or 110 ~ can be distinguished and are seen at 900 and 700 cm -1 for Vas (Fe-O-Fe) and at 350 and 550 cm 1 for Vs (Fe-O-Fe), respectively [2-3]. . FT-Raman spectroscopy is also ideally suited to detect supported complexes as little interference from the support occurs. [Fe2(HPTP)(p-OH ) (NO3)2](C104)2 complexes show two equally strong modes at 655 and 624 cm -1 for Fe202 and FeO(OH) respectively. Upon hetero1500 1000 500 I O0 genisation in MMT of RAMAN SHIFT (cm -1) [Fe2(HPTP)(Ia-OH)(NO3)2] 2+ complex the 655 and 624 Figure 3.a. FT-Raman of [Fe2(HPTP)(OH)(NO3)2](C104) 2 11129.3 ' ~ 3 o q

325

i

i

,

w

|

|

I

i

7

1065 -1

cm Fe202 modes have a decreased intensity. However, all ligand modes of the [Fe2(HPTP)(~t-OH) (NO3)2] 2+ complex on montmorillonite show reduced intensities. The other modes relate to (in)organic ions such as perchlorate that disappear and nitrate modes that 1500 10~00 ~ ' | - r 500 100 decrease in intensity upon heterogenisation (Fig.3.b). RAMAN SHIFT (cm -1) The theoretical perchlorate Figure 3.b. FT-Raman of [Fe2(HPTP)(OH)(NO3)2]2+-MMT modes Vl, v2, v3 and v4 occur at 935, 462, 1102 and 628 cm ~, respectively. The 931 cm -~ is clearly shifted due to interaction with the complex (Fig.3.a). The four characteristic modes of nitrate i.e. v~, v 2, v 3 and v4, are found around 1050 cm -1, 831 cm 1, 720 cm -1 and 1330 cm -1 respectively. In our spectrum of the pure complex Vl and v2, bands at 1058 cm 1 and the shoulder on the 851 cm -1 band around 830 cm 1 are seen. (Fig.3.a). Two theoretical modes of pyridine i.e. vl and vl2 around 1000 cm -1 and 1035 cm 1 respectively. The 1029 cm -1 band assigned to Fem bound pyridine in the complex, is shifted to 1028 cm -1 upon heterogenisation in MMT. In FT-IR the anti-symmetrical Fe202 modes for [Fe2(HPTP)(g-OH)(NO3)2] 2+ complexes are seen at 764 cm -1 with a shoulder at 782 cm 1. The absorptions are broadened for the MMT supported [Fe2(HPTP)(~t-OH)] 4+ complexes at 752 cm 1 with a shoulder at 783 cm1 as seen from difference spectra after substraction of the MMT bands. Part of the nitrate remains associated with the complexes in MMT as the 1384 cm 1 (v 3 of nitrate) shifts to 1352 -1 cm , after introduction of the ~SN nitrate labelled compound. 21

to2s 1

57 ~ I Lrl)t}~ " 1448 ]

"~

i

l

9

I ~~4 I -"

1

T

i~

i

,~1,,) g

7(17 A / , , / "

'

I-

~

u

i

3.3. EPR/DRS-spectroscopy The X band EPR spectra show a very complicated signal from 500 G to 6000 G. The spectra are fairly comparable to these observed for the [(Febpy)20](SOa)2.5H20 and [(Fephen)zO](SOn)2.5H20 complexes [3]. Interpretation of such coupled Fe 3§ d 5 system remains complex. Therefor a W-band EPR study of the complexes was performed. X- and Wband measurements were recorded at 293K. Below 80K, only the S'=0 state is significantly populated, the complexes are EPR inactive and create a diamagnetic matrix. An antiferromagnetic interaction would imply the thermal population of the S=I or S=2 states from the S=0 state. The S=I states should first of all show central transitions both from Ms = 0 to Ms=l and from Ms= -1 to Ms=0, which both cooccur and remain splitted. The [Fe2(HPTP)(NO3)5].4CH3OH complex has magnetic moments of 3.98 ~teff at 297K and 2.10 ~eff at 81K with minor antiferromagnetic coupling (J = -23.9 cm-~). The effect of antiferromagnetic interaction of the complexes requires larger D parameter to fit these spectra in comparison to the mononuclear high spin Fe nI modes. The effect of a small E parameter on these spectra is of very great importance, they affect the splitting of the evenly spaced 5/2---~ 3/2, 3/2---~ 1/2, 1/2---~ -1/2, -1/2--+-3/2 and -3/2---~-5/2 transitions, thus complicating the pattern, as additional splitting

1066 caused by E is not constant, but function of the transition. From the characteristic of the spectral ,q. parameters of multi-frequency EPR, obtained by simulation an idea about the type of interaction can be obtained. At 293K they g~ show properties of an S = 5/2 state with W band spectra with parameters g = 2.000, D = 0.180 v-q -1 cm , % --- (E/D) 0.070 (rhombicity / axiality), with line width AB (p-p)= 10.000 0 5000 60mT and v = 94 GHz. In the X Gauss band, spectra are observed with parameters S = 5/2, g = 2.000, D = Figure 4. X band EPR spectra at 293 K of 0.003 cm -1, % = (E/D) 0.070, with [Fe2(HPTP)(OH)(NO3)2] (C104)2 line width AB (p-p) = 1.5mT and v = 9.250 GHz. Only after cooling to around 100 K the appearance of S = 1 antiferromagnetic interactions are expected. The interpretation is difficult since small variations in zero-field parameters have large effects on the phenomenology of the EPR spectra. The DRS spectra of the pure complex 1.8 and that after heterogenisation in 1.6 MMT show a very broad absorption. 1.4 :3 The absorption is hyperbolic and the 1.2 "band-gap" (defined as an half1 .u0 maximal wavelength) shifts from 459 '~ 0.8 0 0.6 to 435 nm upon heterogenisation w (figure 5). Electronic absorption shows <9Q 0.4 0.2 a broad n--~n* -absorption region that 0 I I I I t is relatively independant of 2~ 3~ 4~ 500 600 7~ 8~ complexation and tailing of the oxoWavelength(nm) to-Fe ]I] charge transfer bands from the Figure 5. Diffuse Reflectance Spectra at 293K of UV in the visible region. With H202 [Fe2(HPTP)(OH) (NO3)2]2+- MMT and tBHP the maxima around 580 nm are assigned to peroxo-to-re charge transfer. The pure complex shows a shoulder around 630 nm probably due to a (p-OH) or (p-OR)-to- Fe III charge transfer. The latter shoulder typically occurs close to the site of oxygen coordination. The binuclear iron(III) HPTP complexes are known to form blue,violet adducts with H202 in a 1 : 1 ratio, the complex being surprisingly more unstable than the HPTB complex [10, 18].

=l-

9

~

III

3.4. M6SSBAUER SPECTROSCOPY For diiron complexes Mrssbauer spectroscopy allows to asses (1) oxidation and spin states of the iron atoms, (2) diamagnetism and ferromagnetism of the ground state for diferric and mixed-valent oxidation levels and (3) valence (de)localisation in the solid state for mixedvalence complexes [2,3]. Isomer shifts (IS) in the range 0.35-0.60 mm/s are characteristic of 5- or 6-coordinate high-spin diferric p-hydroxo complexes [2,3]. Tetrahedral high-spin ferric iron has lower isomeric shifts in the range of 0.22 mm/s [2,3]. For isolated ferric iron with S =

1067 5/2, relatively similar IS values of 0.34 mm/s are seen. In diferric complexes the most important difference between ~t-oxo complexes and g-hydroxo complexes is that the quadrupole shift for the former is > 1 mm/s, while for the latter it is < 1 mrn/s. The reduction in electric field gradient indicated by the smaller quadrupole splittings in the hydroxo-bridged complexes may be due to the lengthening of the Fe-O bond upon protonation of the oxo bridge. This seems a general trend for both dibridged and tribridged complexes [2,3]. The quadrupole splitting is also sensitive to other changes in the coordination sphere. Table 3. Mrssbauer data of related complexes Species IS* QS [(N-b)PcFe]20 0.17-0.20 1.58-

Rel.Int.(%) Reference Murray, 1974.

1.76

[Fe202(N-Et-HPTB)(OBz)] 0.52 [Fez(N-Et-HPTB)(OBz)] 1.07 [Fe2(HPTP)(OH)(NO3)2] (C104)2 0.3088* 0.2988* 0.2188" [Fe2(HPTP)(OH)(NO3)2] z+0.3088* MMT 0.3288*

0.72 3.13 0.42 1.38 0.67 0.38 1.54

100 0.32 0.34 0.84 0.78 0.25 0.54

Que et al., 1990. Que et aL, 1990. **

45 55 49 18 33

* Relative to metallic Fe (for transformation to a-Fe values add +0.0888 mm/s); ** Present work at 4 K. 381"

x

x

x

x

c509 ~. . . . . . . .

379

507

377

5031

'

,

~

505 "~

375 373 42K

dPK -3

-2

-I

0

I

2

3

-3

-2

-I

0

I

2

3

VELOCITY (mm/s) VELOCITY (mm/s) Figure 6. Mrssbauer spectra of [Fe2(HPTP)(OH)(NO3)2](C104)2, and [Fe2(HPTP)(OH)]- MMT. For pure [Fe2(HPTP)(kt-OH)(NO3)2](C104)2 (table 3), the iron speciation at 4.2 K shows two almost identical isomer shifts of 0.29 mm/s and 0.30 mm/s. Their quadrupole splitting however indicates that 55% of the complexes show bent (~t-O) (~-OR) bridging, whereas 45% has bent (~-OH) (~-OR) bridging, corresponding to two QS doublets at 1.38 mm/s and 0.42 mm/s, respectively ~t-O bridging may by related to strongly hydrogen bound p-OH, interacting strongly with charge compensating anions. For the [Fe2(HPTP)(Ia-OH)(NO3)2]2+-MMT complexes the situation is distinctly different at 4.2K. Here two species with either full 0t-O) or full (~t-OH) bridging are encountered, with an IS of 0.31 mm/s and QS of 0.38 mm/s for the ~t-hydroxo complexation

1068 and an IS of 0.32 mm/s and QS of 1.54 mm/s for the ~t-oxo complexation. For the first type some association with nitrates may remain possible. A significant amount (49 %) shows a reduced IS of 0.22 mm/s and a QS of 0.67 mm/s, maybe due to interaction with the clay.

3.5. Alkane Oxidation

Iron salts or iron exchanged MMT deactivate H202 with the generation of 02, and show low alkane oxidation activity (< 0.5%). Complexation of iron by organic ligands has very specific effects on the spin state (II, III) and its electrophilicity and allows non-Fenton-type chemistry [13,14]. Mononuclear ferrous N,N'-bis(2-pyridinecarboxamide)-l,2/3-R complexes in Na-Y zeolite (with R = alkane or benzene) are oxo-type catalysts [13]. Here a different reactivity pattern is expected for the dinuclear ferric complexes [ 14].

Fe/ ',, '

~ F e +H202._ Fe ,/ -H3O+ , ~

\o / H

'

u

he ' OOH

-02 -C6H~1" o

OH

C6Hll

Fe/

,,'\

~Fe

/',,

Fe.

Fe

~ =~

Figure 7. Interaction of [Fe2(HPTP)(~-OH)] 2+ complexes with peroxides (H202 or tBHP) proceeding via a ~t-peroxo intermediate. This intermediate either generates oxygen (catalase activity) or reacts with cyclohexane yielding both cyclohexyl and hydroxyl radicals. The latter start a radical chain reaction yielding cyclohexyl-hydroperoxide from which cyclohexanol/one are obtained. Cyclohexane oxidations by [Fe2(HPTP)(OH)(NO3)2](C104)2 complexes are perfomed at low substrate conversions around 2-5 %. The use of MeCN as a solvent results in an increase of the reaction rate and the selectivity for cyclohexylhydroperoxide (CHHP). In the tBHP oxidation the formation of CHHP and cyclohexyl-tbutylperoxide (CHBP) points to a radical proces. In the presence of organic peroxides triphenylfosfine oxidation occurs as well. Decomposition products such as triphenylphosphine-oxide and cyclohexanol from CHHP confirm the radical nature of the reaction. In the H202 oxidation CHHP is again formed. As the [Fe2(HPTP)(OH)(NO3)2](C104) 2 complexes catalyse fast decomposition of CHHP a rather complete deperoxidation occurs. The peroxo-adducts are more stable in clays than in solution. In solution with [Fez(HPTP)(OH)(NO3)2](C104)2 as catalyst, a cyclohexanone / cyclohexanol ratio of 1.72 is observed. For [Fez(HPTP)(OH)]-MMT a much lower ratio of 1.10 is obtained (table 4).

1069 Table 4. Selectivity in the cyclohexane oxidation for 24h at 293K Conversion chexanol chexanone chexylhydrochexyltbutyl(%) (%) (%) peroxide (%) peroxide (%) HPTP / Fe 6.43 33.4 57.5 3.65 5.36 HPTP / Fe / MMT 12.0 40.8 45.1 8.66 5.43 0.1mmol HPTP/Fe, 0.2 g cyclohexane (2.4 mmol) and 1.8 g tBHP (16 mmol) in 2 g MeCN. In the presence of the radical scavenger ascorbic acid, that supplies the necessary electrons for transfer of free radicals into anions, the complete lack of oxidation of cyclohexane is apparent. Thus involvement of radicals in the oxidation of cyclohexane is proven (table 5). Table 5. Cyclohexane oxidation in absence and presence* of ascorbic acid for 24h at 293K Conversion chexanol chexanone chexylhydro(%) (%) (%) peroxide (%) HPTP / Fe 7.4 31.2 60.3 3.4 HPTP / Fe* +/- 0 +/- 0 +/- 0 +/- 0 0.1mmol HPTP/Fe, 0.2 g cyclohexane (2.4 mmol), (*) 0.15 ascorbic acid and 2 g hydrogen peroxide (30% soln. in H20, 16 mmol) in 2 g MeCN. In the direct oxidation of cyclohexanol a significantly higher conversion is obtained with HPTP / Fe complexes, contrasting with that of monouclear iron complexes [12] or the blanc reaction. Thus cyclohexanone is not only formed through CHHP decomposition, but also by direct cyclohexanol oxidation (table 6). Table 6. Cyclohexanol oxidation in absence and presence of catalyst for 2 - 24 h at 293K. chexanol (%) chexanone (%) no catalyst - 2 h 98 2 no catalyst -24 h 82 18 HPTP /Fe* -24 h 30 70 *0.1mmol HPTP/Fe, 0.2 g cyclohexanol (98% purity), 0.1 g chlorobenzene and 1.8 g tBHP (16 mmol) in 2 g MeCN. The reaction mechanism is studied in batch-reactors for adamantane oxidation (table 7) in dichloromethane, as the solubility of adamantane in MeCN is very low. An 80% tBHP solution in ditertbutylhydroperoxide is used as oxidant to circumvent possible phase separation. For the HPTP / Fe complex adamantan-l-ol is a major oxidation product. For HPTP / Fe / MMT adamantanyl-tbutyl-peroxide is observed and seems stabilised on the clay. Both the homogeous and MMT supported reaction show relative reactivities of the secondary and the tertiary C-atoms in adamantane which are close to these of radical reactions. Table 7. Adamantane oxidation for 24 h at 293K. HPTP /Fe Conv. C Lo/, adam- adam- adamadam adam- adam- adam- adam(%) Cs o 001-ol OH, l-C1 2-one 2-ol 2 - C 1 1-oltBu OOtBu 2-one / (C104-)2 7 0.19 4.5 67.3 1.7 9.8 9.8 13.2 0 2.0 / MMT 11.4 0.14 27.1 54.7 0.5 4.4 4.3 6.6 0.7 1.6 0.05 mmol HPTP/Fe, 0.2 g adamantane (1.47 mmol), 1.8 g tBHP (16 mmol) in 4 g CH2C12

1070 All catalytic results are consistent with non-heme iron monooxygenase activity. The activation of methane by this system remains a challenging goal as iron in inorganic environments has already resulted in promising results [20]. Several clay effects are observed and relate to the support-complex interaction.

Acknowledgments The authors acknowledge IUAP-PAI sponsering from the Belgian Federal Government. PPKG thanks the Fund for Scientific Research-Flanders (FWO) for a research grant. The help of A.Smimov at IERC (Illinois EPR Research Center) for W-band EPR and A.-M. Van Bavel & G. Langouche at KULeuven for M~issbauer spectorscopy is acknowledged.

REFERENCES

1. K.S. Murray, Coord. Chem. Rev., 12 (1974) 1. 2. D. M. Kurtz, jr., D.F. Shriver, I.M. Klotz, Coord. Chem. Rev., 24 (1977) 145. 3. D. M. Kurtz, Chem. Rev., 90 (1990) 585. 4. B. Meunier, Bull. Soc. Chim. France, 4 (1986) 578. 5. H. Toftlund, Acta Chem. Scand., 35A (1981) 575. 6. Y. Nishida, M. Nasu, T. Akamatsu, J. Chem. Soc., Chem. Commun., (1992) 93. 7. A. Weiss, S. Dick, Z.Naturforsch. 49b (1994) 1051. 8. P.P. Knops-Gerrits, S. Dick, A. Weiss, A.M. Van Bavel, G. Langouche, P.A. Jacobs, Microporous Materials (1997) submitted. 9. M.P. Hendrich, E. Munck, B.G. Fox, J.D. Lipscomb, J. Am. Chem. Soc., 112 (1990) 5861. 10. Y. Nishida, M. Takekuchi, H. Shimo, S. Kida, Inorg. Chim. Acta, 96 (1985) 115. 11. Y. Dong, S.Yan, V.G.Young Jr., L.Que Jr., Angew. Chem. 108 (1996) 674. 12. B.A. Brennan, Q. Chen, C.J. Garcia, A.E. True, C.J. O'Connor, L. Que Jr., Inorg. Chem., 30 (1991) 1937. 13. P.P. Knops-Gerrits, M. L'abb6, W.H. Leung, A.-M. Van Bavel, G. Langouche, I. Bruynseraede, P.A. Jacobs Stud. Surf. Sci. Catal., 101 (1996) 811. 14. R.H. Fish, Konings, M.S., Oberhausen, K.J., Fong, R.H., Yu, W.M., Christou, G., Vincent, J.B., Coggin, D.K., Buchanan, R.M., Inorg. Chem., 30 (1991) 3002. 15. Y. Dong, H. Fujii, M.P. Hendrich, R.A. Leising, G. Pan, C.R. Randall, E.C. Wilkinson, Y. Zang, L. Que, Jr., B.G. Fox, K. Kauffman, E. Mtinck, J. Am. Chem. Soc., 117 (1995) 2778. 16. A. Stassinopoulos, G. Schulte, G.C. Papaefthymiou, J.P. Cardonna, J. Am. Chem. Soc., 113(1991)8686. 17. A. Earnshaw, J. Lewis, J.Chem.Soc. (1961) 376. 18. F. Rueda, J.Mendialdua, A.Rodriguez, R.Casanova, Y.Barbaux, L.Gengembre, L. Jalowiecki, D.Bouqueniaux, Surf.Interf.Anal., 21 (1994) 659. 19. W.H. Armstrong, S.J. Lippard, Inorg.Chem., 24 (1985) 981, J. Am. Chem. Soc., 105 (1983) 4837; 106 (1984) 4632; W.H. Armstrong, A. Spool, G.C. Papaefthymiou, R.B. Frankel, S.J. Lippard, J. Am. Chem. Soc., 106 (1984) 3653. 20. J.E. Lyons, P.E. Ellis, V.A. Durante, in "Struct.-Act. and Sel. Relat. in Het. Catal." R.K. Grasselli, A.W. Sleight, Eds., Elsevier, Amsterdam, 99 (1991) 116.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1071

M o d e l i n g the t r a n s i e n t C O o x i d a t i o n o v e r p l a t i n u m T.A. Nijhuis, M. Makkee, A.D. van Langeveld, and J.A. Moulijn Delft University of Technology, Department of Chemical Engineering, Section Industrial Catalysis, Julianalaan 136, 2628 BL Delft, The Netherlands

ABSTRACT

The platinum catalyzed oxidation of carbon monoxide was studied using an advanced transient reactor system, referred to as Multitrack. The experiments indicate, that the reaction is taking place according to the Langmuir-Hinshelwood reaction model. The desorption of CO 2 from the catalyst surface was shown to be a kinetically relevant step in the reaction. From experiments performed using 1802it was shown that isotopic mixing occurs on the catalyst surface due to the decomposition of CO 2present on the catalyst surface. Some experiments showed CO 2 pulse responses having two maximums in the pulse response. This strange occurrence is explained by a moving reaction front model.

1. INTRODUCTION The platinum-catalyzed CO oxidation was studied as a model reaction to evaluate the performance of the Multitrack system. Multitrack, developed in our laboratory, is an advanced version of a TAP (Temporal Analysis of Products) apparatus [1]. The advantages of this model reaction are that the reaction is fast, the only possible reported by-product is carbon, and that the reaction has already been studied in conventional TAP systems [2,3]. A review on the CO oxidation over platinum has been made by Engel and Ertl [4]. In this review they list the following relevant reaction steps: k1 CO( k2 )GOad

(1)

0 2

(2)

k3 ) 20ad k4 ) CO 2

Oad q- GOad Oad q-CO

k~9 )CO2

0 2 "~-2 GOad ~

2 CO 2

(3) (4) (5)

1072 These reactions are based on the assumptions: only irreversible dissociative adsorption of oxygen (2), and instantaneous desorption of the carbon dioxide produced (3)(4)(5). In the temperature range of this study (300-650 K) these assumptions were found to be valid. Reaction (3) represents a reaction according to the Langmuir-Hinshelwood model, whereas reactions (4) and (5) represent reactions according to an Eley-Rideal model.

2. EXPERIMENTAL

The CO-oxidation experiments were performed in the Multitrack system. The basic system of Multitrack consists of three ultra high vacuum chambers. The first contains a small reactor, which can be operated from room temperature up to 1273 K. The catalyst containing reactor had an internal diameter of 7 mm with a total bed height of 10 mm. Gas pulses can be given to the catalyst using high speed gas pulsing valves (pulse width of about 100 ~s). The size of the gas pulses was 2.3"1016 molecules for the experiments with 1602! 2"1016 molecules/pulse for the experiments with 1802! and 1.0"1017 molecules for the CO pulses. Using a flow valve it is possible to perform continuous flow experiments. The flow valve was only used to pretreat the catalyst. The second vacuum chamber is an intermediate (differential pumping) chamber. This chamber is used to prevent molecules to travel indirectly to the mass spectrometers in the third chamber (i.e., not straight out of the molecular beam leaving the reactor). This third section is the analysis section, containing three quadrupole mass spectrometers in-line. These mass spectrometers are used to measure the pulse response of up to three components simultaneously. The maximum possible sampling frequency for the mass spectrometers is 1 MHz. The catalyst used for the CO-oxidation experiments was platinum sponge of 99.99 % purity. The advantage of using platinum sponge as a catalyst is the macroporous structure, resulting in the absence of pore diffusion limitations, while it still has a relatively high surface area for the reaction. The platinum sponge consisted of 2 ~m spheres, sintered together to form 40-100 ~m particles. The amount of catalyst used for the experiments was 291 mg. The platinum was diluted in the reactor using 628 mg of 212-250 lum SiC (carborundum) particles resulting in a total bed height of 10 mm. Prior to use the platinum was pretreated in-situ by oxidation at 773 K, followed by reduction at 673 K. The platinum was analyzed using SEM and krypton BET. The surface area calculated from the SEM micrographs was 0.070 m2/g, the krypton BET gave a 0.081 m2/g surface area. Two basic types of experiments were performed. The first type of experiments are of the multiple-pulse type. In this experiment the catalyst was either first pre-covered with oxygen or carbon monoxide. After this the catalyst was subjected to carbon monoxide pulses or oxygen pulses, respectively. By determining the amount of carbon dioxide produced and the amount of CO or 02 consumed the active surface area of the catalyst can be calculated.

1073 The second type of experiments is of the single-pulse type. In this experiment type single gas pulses are given to the reactor. This type of experiment can also be done by alternately pulsing two different gases. By measuring the individual pulse responses at a high sampling frequency and then modeling these pulse responses, information on the reaction mechanism can be obtained.

3. RESULTS 3.1. Multiple-pulse experiments Multiple pulse experiments were performed in the temperature range of 328 to 423 K. Figure I shows the result of one of these experiments as measured directly. In these figures the amount of CO 2produced can be obtained by integrating the pulse responses, along with the amount of used oxygen or carbon monoxide via a similar integration and subtracting the value from the total amount of the gas pulsed. The results of these calculations are given in Table 1.

3.2. Single-pulse experiments A variety of different single-pulse experiments was performed, only two situations will be discussed here. The first is the steady-state of pulsing alternately CO and 02 over the catalyst. To obtain additional information on the oxygen source of the CO 2produced, this experiment was done using 1802and C160. Figure 2 shows the results of one of these experiments. In this figure it should be noted that a mass spectrometer measures fragmentation patterns of ionized molecules. The small peak at t=l s in both the m / e signal of 28 and 30 is in this case not caused by C1'O and C180 leaving the reactor exit, but mainly by fragments of ionized CO 2molecules.

I CO2

ii ii ii ii ii~

0

5

10

15

20

25

Time (s)

Figure 1. Results from multiple-pulse experiment. Pulsing 1.0"1017molecules of CO over oxygen-precovered platinum (333 K).

1074 Table 1 Results of series of multiple-pulse experiments over Pt sponge. The values given are the total amounts for the multiple-pulse experiment (accuracy +/- 10 %). 02 multiple-pulse (CO covered Pt) 0 2 used C O 2 produced 1 (1017 molecules) (1017 molecules) 328 1.6"10 .2 333 2.6 2.9 338 2.3 2.5 343 3.1 2.6 373 2.1 2.3 393 2.2 2.9 423 2.3 2.2 1. equals CO adsorbed at start of 0 2 pulsing. 2 equals adsorbed 0 2 a t start of CO pulsing.

CO multiple-pulse (O~ covered Pt) CO used CO 2produced 2 (1017 molecules) (1017molecules) 6.1 3.1 6.1 3.2 5.8 2.9 6.4 3.2 5.9 2.9 5.7 3.0 5.0 3.2

T (K)

48

C1802

46 C 1 6 O a 8 0

28

0

C160

I

I

I

I

I

0.5

1

1.5

2

2.5

Time

(s)

Figure 2. Steady-state results of CO oxidation with 1802over Pt sponge at 573 K. Experiment done by pulsing 2"1016 molecules of lSO2 at t=0 and 1.0"1017 molecules of CO at t=l s. At t=3.5 s the next 1~O~pulse is given (equals next t=0).

1075 The second series of single pulse experiments described here is the situation where the catalyst is first pre-oxidized and then slowly reduced by pulsing alternately 1017 molecules of CO and 2.3"10 ~ molecules of oxygen. The aim was to study the reaction at different oxygen surface coverages. From the second until fifth CO pulse-response measured a strange phenomenon was observed: two maximums in the measured CO 2pulse-responses, seemingly indicating two different types of reaction sites. An example of this is shown in figure 3. To gain more insight in this occurrence this experiment was repeated after pre-covering the platinum with '602 and then alternately pulsing C160 (1017 molecules) and '~O~ (2"10~ molecules). The result of one of these experiments is shown in Figure 4.

02

co 0

0.2

0.4

0.6

0.8

1

T i m e (s)

Figure 3. Results of single-pulse experiment. Pulsing CO over partially oxygen covered platinum (573 K). CO pulse given after 3 cycles of CO pulse followed by O 2 pulse over initial fully O-covered platinum. 4. DISCUSSION 4.1. Reaction m o d e l

In Table I it can be seen that at a temperature of 328 K the pulsing of oxygen over platinum precovered with carbon monoxide does not produce any carbon dioxide. This can be explained by a very strong adsorption of CO on the platinum, inhibiting the dissociative adsorption of oxygen. This is in agreement with the experiments done in the temperature range from 333 K to 393 K (not shown), where when pulsing oxygen over CO pre-covered platinum, there is initially no CO 2production and the reaction only starts after some CO has desorbed. The time necessary for the desorption of a small amount of CO is shorter at a higher temperature. Apparently it is necessary to have free sites on the surface for the reaction to start, indicating that the reaction takes place according to the Langmuir-Hinshelwood reaction model.

1076

C1601aO ,..,,

"~

C1602

9

0

~:,~.,,,.,~.~

I

I

I

I

I

I

I

I

I

I

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

Time (s)

Figure 4. Results of single-pulse experiment. Pulsing C160 over partially oxygen covered platinum (573 K). C~60 pulse given after 3 cycles of C~60 pulse followed by 1802 pulse over initial fully ~60-covered platinum. Pulsing CO over O-covered platinum does not show an inhibition. This can either be because the reaction of CO with adsorbed oxygen can occur according to the EleyRideal model, or because some free sites are always available for CO adsorption. Whichever is the case cannot be concluded from the experiments performed. 4.2. Number of active sites From the surface areas determined by SEM and Krypton BET the number of sites on the platinum surface can be estimated. This results in 2.7"1017 sites for SEM micrographs and 3.1"1017 sites for the BET measurement. If both the amounts of COmproduced in Table I are added and compared to the amounts of oxygen and carbon monoxide used, it can be seen, that the amount of oxygen and CO used, corresponds to the amount needed for the CO 2produced. The amount of CO 2produced during one multiple pulse experiment should be equal to the number of active sites on the catalyst. If the amount of produced carbon dioxide is compared to the calculated number of platinum surface atoms from the SEM micrographs or B.E.T. measurement, it can be seen that these are in good agreement. 4.3. Isotopic mixing at the catalyst surface Figure 2 shows, that in the steady-state situation when pulsing C160 and 1802 sequentially, three types of CO 2are produced: C1602, C160180, and C1sO2, although only one type, C1601~O, might be expected. The production of the two other CO 2 isotopes indicates that splitting of the C-O bond takes place. This is explained by: 1. CO decomposes on the surface and produces Cads and Oads; 2. CO2adsproduced decomposes to COad, and O ad` The decomposition of CO or CO 2in the gas phase is very unlikely at the low reaction temperatures of the experiments and is, therefore, not taken into account.

1077 Possibility I is not likely to be the main cause because of: 9 when only CO is pulsed over the platinum sponge, no CO 2 production is seen. 9 during the ~O2 gas pulse a significant amount of C~602 is produced, and it is not likely that the decomposition of CO is induced by the presence of oxygen. The second possible explanation for the occurrence of the three types of CO 2 is based on the mechanism proposed in figure 5. 0 II

0 II

II

II

c

O~

CO

c

\

CO

dII dII

/ 0

0

c

c

II,,

II,,

0

o" II

o*

--

II

.~

~

[o=c

II ..................

7d]

o

*

-o,c

o

II

c

II

coo* o" II

C02

o

o" II

c

II

cG

o

d .

II

c

tl

o

Ii

Figure 5. Model for the mechanism of the CO oxidation. The shaded areas represent a platinum surface. The oxygen atoms marked with a * denote oxygen-18. Adsorbed CO reacts fast and reversible with adsorbed oxygen to form adsorbed CO2~ As in the decomposition of this CO 2both oxygen atoms are identical, both have an equal chance to be split off. The Oads at the surface can then be either from the CO or the originally adsorbed O atom. By this pathway it is possible to form C180, that can then produce C1~O2without CO dissociation. As this mechanism produces adsorbed 160 from adsorbed C160 during the reaction between C160 and ~8Oads,it should also be possible to produce C1~ The C1~O2is produced simultaneously with both other types of CO2. As ClSO2 can only be produced by a Langmuir-Hinshelwood model (i.e., reaction only at the surface), it is most likely, that all CO 2 produced from the reaction between CO and adsorbed O atoms is produced in a similar manner. From the ratio between the produced amounts of C~~ C~60~80, and C1~O2the relative rate of scrambling compared to desorption can be determined. If the isotopic scrambling of the CO 2 is very fast compared to the desorption of CO 2, the ratio C1602 : C~O~aO : C~802 should be I : 2 : 1, if the desorption and isotopic mixing of CO 2 occur

1078 at the same rate the ratio should be I : 6 : 1. From the observed ratios (about 2:4:1 (573K) and 2:6:1 (373 K)) we conclude that the rates are of the same order of magnitude. The same isotopic mixing of the CO 2 produced was also found by Huinink [3], who performed comparable experiments on a conventional TAP system.

4.4. Double pulse-responses Figure 3 shows two maximums in the pulse response for CO 2 of a single CO pulse given. From the isotope experiment shown in Figure 4, it can be seen that the two maximums in the CO 2 pulse response originate from different oxygen sources, as the first m a x i m u m contains 180 and the second maximum does not. The first CO 2 production peak, therefore, comes from the freshly added 1802 and the second peak from the 1602of the precovering. The first CO 2 peak does not, however, only consist of C160180, but also of C1602, which can be explained by the isotopic mixing, as described previously. In Figure 4 a CO pulse response is detectable, whereas this is not the case in Figure 3. Figure 4 shows the last pulse-response measured at the time all 'old' oxygen is consumed, which is not yet the case in Figure 3, due to the slightly smaller size of the 1802 pulses compared to the 1602pulses. Three possible explanations for the difference in reactivity between 'fresh' and 'old' oxygen are: 9 Platinum has two different kinds of reactive zones for the CO oxidation, possibly from different crystal planes. It is, however, unlikely that these zones will cause such a sharp distinction between freshly added oxygen and oxygen with a longer residence time on the surface. 9 A second possibility might be that subsurface oxygen is taking part in the reaction. Subsurface oxygen could cause an extra amount of oxygen to be available for the reaction. The time necessary for the oxygen to move from subsurface to surface could cause the delay between the two peaks. The amount of oxygen atoms (sites), however, calculated to take part in the reaction, corresponds only to the amount of surface sites (see paragraph 4.2). If subsurface oxygen was to take part in the reaction, this amount should have been higher. 9 A third possibility, is that the reaction is taking place according to a 'moving reaction front model', i.e., the reaction rate is so fast, that the adsorption and reaction take place immediately when the gas pulse arrives at the catalyst bed. This model is described in detail in the next paragraph. 4.5. Moving reaction front model In experiments where the adsorption a n d / o r reaction of a component take place much faster than the transport of that component in the axial direction, this can cause the reaction to occur in a zone moving through the catalyst bed. A very fast reaction compared to the gas transport is consistent with Figure 3, where all CO is used, as enough oxygen is available on the catalyst. Figure 6A schematically shows the surface coverage as a result of this type of moving-reaction-front behavior in case CO is pulsed repeatedly over an oxygen-covered platinum-catalyst bed. Figure 6B shows what would happen to the surface coverage if alternately (excess) CO and 02 are pulsed over the same catalyst bed.

1079 A

ii!i!iiiiiiiiiiiiiiiiii~iiiiii!iiiiii~iii~ii~i!ii!iiii~ii~il

ii[!ii~i~iiiii~ii~i~ii]i ii iiliiliiiiiiii i i i!i#~i!~ii i!~ii i !i~!ii!~i~i i !iiiiiii)~iiiiiiii!i!!i !i !i~i ~!~!~i ii~iiiiii!i~i i~i~;!~i!~i~ii~@~i~.~i ..... ~!

..... ~.

3a

4

ji~~',i~i,'!i)i,'i!i!i!i~!]

li!ili!i .iiil~176 liilil

ll)iil

Figure 6. Schematical representation of the moving-reaction-front model. A: pulsing CO over O-precovered platinum. B: Pulsing CO (excess) and 02 alternately over initial O-covered platinum. Gray: O-covered; White: CO covered or clean platinum. The numbers denote the pulse cycle number. Figure 6B demonstrated that after the first oxygen pulse three zones with different surface coverages have formed (from left to right): 1. an oxygen covered zone, with oxygen from the last added pulse 2. a (partial) CO covered zone, from the CO pulse(s) 3. an oxygen covered zone, with oxygen on the surface from the initial coverage When a new CO pulse is given, zone I will immediately produce CO 2 and become (partially) CO covered (=part of zone 2), the first part of zone 3 will also produce CO 2 and also become part of zone 2. As CO continuously adsorbs on / desorbs from the platinum, CO travels slower through the catalyst bed than CO 2. This causes the CO 2 produced in zone I to leave the reactor earlier than the CO 2produced in zone 3, explaining the double pulse response. This model is consistent with the isotope experiment in Figure 4, where zone 1 would be 180 covered and zone 3 1~Ocovered, resulting in the production of an isotopic mixed CO 2 first and not labeled CO 2 later. At the time all 'old' 160 has been used, the CO will start 'breaking through' the catalyst bed, as visible in figure 4. The same model can also explain why in steady state, when alternately pulsing excess CO and 02, CO 2 is formed after both the CO and 02 pulses. If this moving-reaction-front behavior would not occur, the oxygen pulse would completely react with all CO on the platinum, leaving no oxygen left for the next CO pulse to react with. To confirm the validity of this model, computer simulations were carried out [5]. These computer simulations showed the same double pulse responses. This moving reaction front behavior implies that not the reaction rate as a function of a varying oxygen surface coverage is measured, as was planned, but the reaction rate as a function of the location of the reaction front in the reaction.

1080

5. CONCLUSIONS The CO oxidation takes place via the Langmuir-Hinshelwood model, with the following set of equations: k1 C O ( k~ ) GOad

(6)

02

"(7)

k~" ) 2 O a d k4

GOad -}- Oad < ks ) C O a a d

(8)

CO2a d

(9)

k~ > C O 2

9 The CO 2 is produced from CO and O on the catalyst by an equilibrium reaction that is faster than the CO 2 desorption. This implies that the rate of desorption of CO 2 from the platinum surface should be taken into account when modeling this reaction. 9 Multitrack offers the possibility to accurately determine the number of active sites on a catalyst. 9 The surface coverage on a catalyst can vary strongly as a function of the reactor position, even in a small reactor as used in Multitrack. 9 At higher temperature the CO oxidation in Multitrack takes place according to a moving reaction front pattern. 9 The moving-reaction-front model is able to describe the occurrence of the double pulse responses obtained in measurements.

REFERENCES

1. J.T. Gleaves, J.R. Ebner, and T.C. Kuechler, Catal. Rev. - Sci. Eng., 30(1), p. 49 (1988). 2. F.H.M. Dekker, J.G. Nazloomian, A. Bliek, F. Kapteijn, J.A. Moulijn, D.R. Coulson, P.L. Mills, and J.J. Lerou, Carbon Monoxide Oxidation over Platinum powder; A comparison of TAP and Step-Response Experiments, Appl. Catal. A, acceptedfor publication (1996). 3. J.P. Huinink, A Quantitative Analysis of Transient Kinetic Experiments: The Oxidation of CO by O2/NO on Pt, Ph.D. Thesis, Eindhoven University of Technology, 1995. 4. T. Engel and G. Ertl, Elementary steps in the catalytic oxidation of carbon monoxide on platinum metals, in D.D. Eley, H. Pines and P.B. Weisz (editors), Advances in Catalysis, vol. 28, Academic Press, New York,1979, p. 1. 5. T.A. Nijhuis, A.D. van Langeveld, M. Makkee, and J.A. Moulijn, paper in

preparation.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1081

Dioxygen Activation with Sterically Hindered Tris(pyrazolyl)borate Cobalt Complexes. K. H. Theopold a, O. M. Reinaud b, D. Doren a, and R. Konecny a aDepartment of Chemistry and Biochemistry, Center for Catalytic Science and Technology, University of Delaware, Newark, DE 19716, USA (e-mail: [email protected])* bEcole Nationale Superieure de Chimie de Paris (ENSCP), Laboratoire de Bioorganique et Biotechnologies, 11 rue Pierre et Made Curie, 75231 Pads Cedex 05, France Dioxygen is the most readily available strong oxidant. It may serve as an electron acceptorproducing the environmentally benign molecule water, or as a source of oxygen atoms to be incorporated into organic products - leaving no byproducts whatsoever. These obvious advantages are countermanded by the kinetic limitations of its reactions, which are marked by large initial kinetic barriers, low selectivities, and radical chain mechanisms producing undesirable "combustion" products (e.g. CO, CO2). The "activation of dioxygen", i.e. its transformation into a highly reactive, and yet selective reagent for the oxidation of a wide range of organic substrates, is thus a major challenge for research in catalysis. Herein we describe our efforts to design metal complexes for this ultimate purpose. 1. B A C K G R O U N D - C H O I C E OF THE TpCo-SYSTEM We began with the observation that dioxygen coordinated to a transition metal is not inherently a reactive species - i.e. mononuclear dioxygen complexes (M-O2) are not generally strong oxidants. To unleash the oxidizing power of 02 requires breaking the O-O bond, thus generating metal oxo moieties. Even the most cursory consideration of biological oxidation systems (e.g. cytochrome P450, methane monooxygenase, etc.) would suggest that the reactive species generated by these enzymes are metal oxo complexes featuring late transition metals (e.g. Fe and Cu) in relatively high formal oxidation states.[ 1] The instability of such species has been noted and supports the expectation of high reactivity.[2] Given that transition metals metals to the right of (and including) iron are unlikely to support a direct oxidative addition of 02 to a single metal center (a 4-electron process), the most straighforward scheme for the activation of dioxygen involves its coordination and cleavage by at least two metal atoms, e.g. according to eq. 1. Reaction of the M=O moiety with substrate regenerates the metal in its reduced form (M); catalytic turnover would thus be possible. M 02 ~

M M'-O 2 ~

0,, M--O2-M ~

M

M ~

2 M=O

(1)

',0 / Due to the extensive O2-binding chemistry of cobalt,[3] we chose the latter as the metal for our initial studies. However, while the available precedent consists largely of reactions of *This research was supported by the U.S. Department of Energy (ER14273 to KHT), the National Science Foundation (CHE-9401312to DD) and by CRNS and NATO (fellowshipsto OMR).

1082 octahedral Co(H) coordination compounds with 02, yielding Co(III) superoxo and kt-peroxo complexes, we felt that another coordination geometry might be more favorable for our purposes. Specifically, some earlier work of ours with unusually high-valent cobalt organometallics (Co(IV), Co(V))[4] led us to aim for tetrahedral cobalt, in the hopes of facilitating the formation of oxo species in high formal oxidation states. Keeping in mind that one of the ligands must be the 02 or 02- moiety, we were thus looking for a ligand to occupy the remaining three coordination sites of a tetrahedron, while at the same time preventing the formation of complexes with higher coordination numbers. Tris(pyrazolyl)borate ligands with bulky substituents in the 3-position of the pyrazole rings (so called "tetrahedral enforcers")[5] seemed an attractive choice.[6] This class of ligands offers great flexibilty in tuning the steric requirements of the remaining fourth coordination site;[7] their monoanionic nature compensates some positive charge of high-valent metal ions, and combined with the lipophilic character of the ligand solubility of the complexes in nonpolar organic solvents might be expected. The particular ligands used in this work were Tp/-Bu,Me and Tp i-Pr,Me (see below). H

H

%

Tpt-Bu,Me

Tpi-Pr,Me

Both the t-Bu and i-Pr substituents are bulky enough to prevent the formation of octahedral Tp2Co-derivatives; however, whereas the Tpt-Bu,MeCo-complexes were all found to be mononuclear, the Tpi-Pr,MeCo-series featured several dinuclear compounds, due to the lesser steric hindrance of this ligand. Indeed, while this difference may seem trivial, it turned out to be of crucial importance for the activation of 02. 2. STRUCTURE AND REACTIVITY OF COMPLEXES We have prepared and characterized a series of cobalt complexes representing substantially all of the species shown in eq. 1. This series of compounds provides a unique view of the stepwise transformation of 02 from an unreactive molecule into a highly reactive species capable of abstracting hydrogen from saturated alkyl groups.

2.1. Coordinatively unsaturated cobalt complexes Our scheme calls for the preparation of coordinatively unsaturated, or at least substitution labile, cobalt complexes that will bind 02. Reduction of halide precursor TpCo-I with magnesium metal under N2 yielded the corresponding dinitrogen complexes, see eq. 2.[8,9] TpR,MeCo-I

Mg, N 2 ~- Tpt'Bu,MeCo-N2 or Tpi'Pr,MeCo-N2-CoTp i'Pr,Me

(2)

1083 The result of a crystal structure determination of dinuclear [(Tpi-Pr,MeCo)2q.t-N2)]is shown in Figure 1. A notable feature of the structure is the position of the nitrogen ligand, which binds to the cobalt not in the expected symmetric fashion (pseudo C3v symmetry), but rather is bent away from the threefold axis established by the B-Co vector by a substantial margin (a = 260). This particular distortion is a consistent structural feature of 16-electron Co(I) (d 8) complexes of the type TpCo-L. We have analyzed the driving force for this distortion with the aid of density functional (DFT) calculations; at the same time we have also demonstrated that such calculations can now be carried out on large enough molecules to enable the modeling of the actual tris(pyrazolyl)borate complexes with high structural fidelity.[ 10]

7c,,0,

C(8'~ C(9LAC111)~ o"

"~

~P

At,-" .....~-

~

? ~

c(,(~p

N(3o}

~C(8al Figure 1. The molecular structure of [(Tpi-Pr,MeCo)2(~-N2)]; Selected interatomic distances and angles: Co-N(7), 1.790 A; N(7)-N(7a), 1.151 A; Co-N(7)-N(7a), 171.6 o. The dinitrogen complexes are very reactive, undergoing facile ligand substitution reactions with a variety of added Lewis bases. For example, exposure to carbon monoxide gas produced the carbonyl complexes TpR,MeCo-CO in quantitative yield. 2.2.

Mononuclear

dioxygen complexes

Exposure of a solution of Tpt-Bu,MeCo-CO to dioxygen gas yielded the dioxygen complex Tp tBuM " , eCo-02 1"n a rapid and quantitative reaction, see eq. 3.

C

~

%

o-o@_ Co /\

(3)

As its vibrational spectrum (vo-o = 961 cm -1) did not allow an unambiguous assignment of the nature of the coordinated dioxygen (superoxo vs. peroxo), the crystal structure of the complex was determined. It revealed a dioxygen ligand coordinated in the side-on fashion

1084 thought to be characteristic of peroxo complexes; however, the O-O bond distance measured only 1.262(8) A, consistent with a superoxide assignment. Tpt-Bu,MeCo-O2 was indeed the first structurally characterized example of a side-on bound superoxo ligand; a few other examples, all featuring tris(pyrazolyl)borates as ancillary ligands, have since been found.[11,12] Tpt'Bu,MeCo-O2 is a surprisingly stable compound; it can be crystallized from hot toluene without any decomposition. Significantly, it does not react with olefins or even triphenylphosphine under ordinary conditions. These observations underscore our notion that simple complexation of 02 to a transition metal does not "activate" the former. The analogous treatment of a solution of Tpi-Pr,MeCo-CO with 02 produced a dinuclear carbonate complex instead of the dioxygen complex, see eq. 4. This compound must result from a bimolecular reaction of the desired Tpi'Pr,MeCo-O2 with its carbonyl precursor; apparently the difference in steric protection offered by the two Tp-ligands can have dramatic chemical consequences. However, once recognized, a simple solution to the preparative problem suggested itself. Exposure of solid Tpi-Pr,MeCo-CO (finely ground) to 02 gas for several hours yielded analytically pure Tpi-Pr,MeCo-O2. Spectroscopic characterization of this material strongly suggested that it is a structural analog of the Tpt-Bu,Me-complex. Note, however, that its thermal stability is significantly different (vide infra).

I0 ~

0\

TpCd~~CoTp

~.,

02

s-'olution

Tp/-Pr,MeCo-CO

02

~ Tpi-Pr,MeCo-O2 solid state

(4)

2.3. Coordination of the second metal atom

Ultimate cleavage of the O-O bond of dioxygen requires the reducing equivalents of a second metal atom. There are several ways in which the formation of a dinuclear dioxygen complex might be achieved. Most obviously, the addition of a second equivalent of metal, e.g. in the form of the nitrogen complex, to the mononuclear dioxygen complex would serve this purpose. When one equivalent of Tpt'Bu,MeCo-N2 was added to a solution of Tpt'Bu,MeCo-O2, a rapid reaction occurred, which produced slightly less than two equivalents of Tpt-Bu,MeCo-OH (90 % yield by NMR), see eq. 5.

Tpt-Bu,MeCo-O2 + Tpt-Bu,MeCo-N2

"-,.--

2 Tpt-Bu,MeCo-OH

(5)

Essentially no incorporation of deuterium label was observed when the same reaction was carded out in toluene-d8. Since the formation of more than one equivalent of Co(H) hydroxide requires the transfer of hydrogen atoms, we hypothesized that the mechanism of this transformation involved the formation of a dinuclear g-peroxo complex, followed by rapid cleavage of the O-O bond and hydrogen atom abstraction from the Tp-ligand. The inferred activation of the primary C-H bonds of the tert-butyl groups of the ligand requires a highly reactive species; thus the "activation" of 02 seems to have been accomplished. In a bid to prove that the ligand is the source of the hydrogen atoms, the reaction was repeated with the analogous Tpi-Pr,Me-complexes. We reasoned that hydrogen atom abstraction from the isopropyl groups would generated isopropyl radicals, which should disproportionate into isopropyl- and isopropenyl-substituted ligands. Reaction of Tpi-Pr,MeCo-O2 with half an equivalent of [(Tpi'Pr,MeCo)2(g-N2)]_produced a mixture of products, one of which was the dinuclear hydroxide complex [(Tp~-Pr,MeCo)2(kt-OH)2]. Hydrolysis of the whole product mixture, and examination of the recovered pyrazoles indeed showed that 10 % of the pyrazole fraction was 3-isopropenyl-5-rnethylpyrazole. While the reaction of Tpt'tiu,meCo-O2 with Tpt'Bu,MeCo-N2 produced no observable intermediate under any. conditions we have tried, monitoring of the parallel reaction of Tp iPr,MeCo-O2 with [(Tp~'Pr,MeCo)2(I.t-N2)] by 1H NMR at low temperature (-30~ revealed the

1085 formation of a new species, which decomposed to the hydroxide upon warming. The same species was also generated quantitatively, when solid [(Tpi'Pr~eCo)2(I.t-N2)] was exposed to 02 gas at-78oc. Given the dinuclear structure of the dinitrogen complex (see Fig. 1), the product of the solid state reaction is likely to be a dinuclear complex also. We originally assigned to this intermediate the structure [(Tpi-Pr,MeCo)2(kt-O2)], i.e. a Ix-peroxo complex preceding the O-O cleavage step. Presumably the lesser steric hindrance of the Tpi-Pr,Me-ligand stabilizes this dinuclear intermediate sufficiently for it to be accessible. 2.4. Structure of the dinuclear complex - [(TpCo)2(l.t-O2)] vs. [(Tp.Co)7,(Ix-O)2]

Based on the close analogy to Kitajima's hemocyanin model complex [(TpZ-rr,~-erCu)2(lxO2)],[13-15] and in the absence of an X-ray structure determination, we had proposed a IXperoxo structure - i.e. [(Tp/-Pr,MeCo)2(Ix-O2)] - for the dinuclear intermediate.J9] The same structural motif has subsequently been suggested for a close analog, i.e. [(Tp/-Pr,i'PrCo)2(~t02)], based on spectroscopic data.[ 16] However, recent results reported by Tolman et al. cast some doubt upon this assignment.[17,18] Specifically, the possibility of a bis(Ix-OXO) structure, resulting from facile (reversible?) cleavage of the O-O bond must be considered. Since these two structures have been shown to be of essentially equal stability for copper, the greater tendency of cobalt to favor higher formal oxidation states (i.e. Co(II) for the (I.t-O2) complex and Co(III) for the bis(kt-O) species) suggests that [(Tpi-Pr,MeCo)2(kt-O)2] may be the more stable isomer in the cobalt system. Repeated attempts to obtain suitable crystals for an X-ray structure determination have not met with success. Therefore we have used DFT calculations to determine the structure of the dinuclear complex (omitting the alkyl substituents). Full geometry optimizations beginning with two different initial structures, namely a kt-peroxo geometry and a bis(kt-oxo) structure, ultimately converged on the same result. The structure of lowest energy is depicted in Figure 2; it is a bis(Ix-oxo) complex!

Figure 2. The structure of [(TpCo)2(kt-O)2], as predicted by DFT calculations. Selected interatomic distances and angles: Co-Co, 2.756 A ; Co-O, 1.886./k, 1.765/~ ; O-O, 2.399 A. The stepwise transformation of the Ix-peroxo structure into the bis(I.t-oxo) complex followed a smooth trajectory, along which the total energy decreased with every iteration. There is no apparent barrier to the cleavage of the O-O bond, once the dinuclear peroxo-complex is formed.

1086 While the g-rl2:rl2-peroxo complex is a hypothetical structure, the difference in energy favoring the bis(g-oxo) isomer was estimated to be 30 kcal/mol. Finally, the dissociation of the bis(goxo) dimer into two terminal oxo complexes (i.e. TpCo=O, see eq.1) is prohibited by a large energy difference (ca. 75 kcal/mol). If these calculations prove correct, the high reactivity of the "activated" 02 (i.e. H-abstraction from unactivated C-H bonds) may be associated with bridging rather than terminal oxo-ligands. We note that the most recent proposal for the reactive intermediate of methane monooxygenase features a "Fe2(g-O)2 diamond core structure" much like the one shown in Figure 2.[ 19]

2.5. Direct conversion of TpCo-O2 into [TpCo2(g-O)2] The high thermal stability of Tpt-Bu~eCo-O2 has been commented on above. It presumably reflects the prohibitive enthalpic cost of dissociation of 02 from the complex, which in turn would generate coordinatively unsaturated Tpt-Bu,MeCoI. The latter could then react with the dioxygen complex, leading to O2-cleavage and ligand oxidation (see section 2.3. above). Indeed, the complete absence of this dissociation equilibrium is a great obstacle for any attempts to run the O2-activation in a catalytic mode (i.e. in the presence of excess 02), because all available cobalt would be tied up in the form of unreactive rnononuclear dioxygen complex. However, somewhat to our surprise, mononuclear Tpi-Pr,MeCo-O2 proved unstable in solution at ambient temperature. Its decomposition produced [(Tp/-Pr, MeCo)2(g-O)2] as a transient intermediate and ultimately proceeded to the hydroxide. Significantly, no additional metal complex was needed to initiate the decomposition, and the reaction took place even in the presence of dioxygen. It is highly unlikely that the substitution of an isopropyl group for a tertbutyl substituent would change the electronic nature of the mononuclear dioxygen complex sufficiently to shift the 02 dissociation equilibrium. Thus there must be another .pathway, which is open to Tpi-Pr,MeCo-O2, but closed to the sterically more encumbered Tpt-t~u,MeCo02. A significant difference between the two mononuclear dioxygen complexes emerged from low temperature 1H NMR studies. Whereas the isotropically shifted resonances of paramagnetic Tpt-Bu,MeCo-O2 exhibited simple Curie-type temperature dependence, cooling of a solution of Tpi-Pr,MeCo-O2 revealed the formation of a diamagnetic species in a temperature dependent equilibrium. The result of a structure determination of this new compound is shown in Figure 3. It is a dinuclear complex featuring two superoxide ligands linking the cobalt atom in a wans g-rl 1:111 fashion; the Co204 core of the molecule adopts a chair conformation.[20] C(37) CI341

CI19) ~

~381

'~"

Figure 3. The molecular structure of [.(Tpi-Pr,MeCo)2(g-O2)2.]; selected interatomic distances: Co-O1, 1.836(3)/~, Co-O2, 1.839(4) A, O1-O2A, 1.354(5) A; Co-CoA, 3.50/~,.

1087 The thermochemical parameters governing the dimerization equilibrium (see eq. 6) have been determined; they were AH = -14.5(5) kcal/mol and AS = -60(3) cal/mol.K. The energetic accessibility of this structure suggest a mechanism for the formation of [(Tpi'Pr,MeCo)2(kt-O)2] via dimerization of Tpi-Pr,MeCo-O2, followed by loss of one molecule of 02, and cleavage of the O-O bond of the remaining dioxygen (eq.6).

/O(•o.• r

2 Tpi'Pr-MeCo-O2 ~

TpCo

CoTp

O

2.6. H y d r o g e n

atom

-02

~

,~ .-. /Ox lpt.:o\ ,,CoTp O

(6)

abstraction

The complex to which we now assign the structure [(Tpi-Pr,MeCo)2(kt-O)2] is presumably the reactive species attacking the C-H bonds of the ligand's alkyl substituents. The intervention of mononuclear species featuring terminal oxo ligands is ruled out by the large calculated dissociation energy for the Co202 core. For similar reasons, reversion to a li-peroxo intermediate on the path to ligand attack is energetically unfavorable. A closer examination of the abstraction reaction is of interest. The rate of the reaction in CH2C12 showed clean first order dependence on the concentration of [(Tpi'Pr,MeCo)2(I.t-O)2]. There was no appreciable solvent effect upon the rate constants (toluene, dichloromethane, acetonitrile). The activation parameters were AH~ = 16.4(5) kcal/mol and AS ~ = - 12(1) cal/deg.K. Most interestingly, the reaction exhibited a large kinetic isotope effect; i.e. substitution of D for H in the tertiary position of the isopropyl groups gave kH]kD = 22(1) at 281 K. One possible explanation for such a large isotope effect would be a contribution of quantum mechanical tunneling of the hydrogen atoms. Further support for this notion was provided by differences in the activation parameters of the H- versus D-abstraction (AAH~ = 2.8 kcal/mol and AH/AD = 0.13).[21] On the basis of these observation we have proposed that the intramolecular hydrogen atom abstraction suffered by [(Tpi-Pr,MeCo)2(I.t-O)2] involves tunneling. In this context it is interesting to note that both oxidation enzymes and related model systems have shown similarly large isotope effects. For example, Lipscomb has reported kH/kD = 50100 for the oxidation of methane by methane monooxygenase,[22] and Tolman et al. have reported large isotope effects in the decomposition of their Cu2(kt-O)2 model complexes.[23] The exact origin of these large isotope effects is currently unclear, but may well be tunneling also. Hydrogen bonding interactions (C-H...O)between ligand C-H bonds and the kt-oxo ligands have been observed in the copper complexes;[18] these may be partially responsible for the narrow activation barriers thought to be required for quantum mechanical atom tunneling. 3. C O N C L U S I O N S We have shown that rationally designed transition metal complexes can be used to "activate" 02, i.e. to transform it into a highly reactive oxidant capable of activating C-H bonds. The reactive species generated in this reaction sequence is a bis(I.t-oxo) complex resulting from cleavage of the O-O bond of a dioxygen ligand. The resulting M2(I.t-O)2 diamond core structure is emerging as a common stuctural feature in late transition metal dioxygen chemistry (e.g. for Mn, Fe, Co, Cu). There still remain several obstacles which must be overcome to permit application of this chemistry in a catalytic cycle. Chief among these is the attack of the reactive species on the ligand. To prevent this, we are preparing modified tris(pyrazolyl)borate ligands, which do not present abstractable hydrogen atoms to the oxo-ligands. In anycase, the promise of homogeneous catalysts for air oxidations of a wide range of organic substrates, including alkanes, warrants further investigation of this system.

1088 REFERENCES

1. Lippard, S. J.; Berg, J. M. Principles of Bioinorganic Chemistry; University Science Books: Mill Valley, CA, 1994; p. 302 ff. 2. Mayer, J. M. Comments Inorg. Chem. 1988, 8, 125. 3. Cotton, F. A.; Wilkinson, G. Advanced Inorganic Chemistry; 5th ed.; Wiley: New York, 1988; p 735 ff. 4. Byrne, E. K.; Theopold, K. H. J. Am. Chem. Soc. 1989, 111, 3887. 5. Trofimenko, S.; Calabrese, J. C.; Thompson, J. S. Inorg. Chem. 1987, 26, 1507. 6. Trofimenko, S. Chem. Rev. 1993, 93,943. 7. Kitajima, N.; Tolman, W. B. Prog. lnorg. Chem. 1995, 43, 419. 8. Egan, J. W.,Jr.; Haggerty, B. S.; Rheingold, A. L.; Sendlinger, S. C.; Theopold, K. H. J. Am. Chem. Soc. 1990, 112, 2445. 9. Reinaud, O. M.; Theopold, K. H. J. Am. Chem. Soc. 1994, 116, 6979. 10. Detrich, J. L.; Konecny, R.; Vetter, W. M.; Doren, D.; Rheingold, A. L.; Theopold, K. H. J. Am. Chem. Soc. 1996, 118, 1703. 11. Zhang, X.; Loppnow, G. R.; McDonald, R.; Takats, J. J. Am. Chem. Soc. 1995, 117, 7828-7829. 12. Fujisawa, K.; Tanaka, M.; Moro-oka, Y.; Kitajima, N. J. Am. Chem. Soc. 1994, 116, 12079. 13. Baldwin, M. J.; Root, D. E.; Pate, J. E.; Fujisawa, K.; Kitajima, N.; Solomon, E. I. J. Am. Chem. Soc. 1992, 114, 10421. 14. Kitajima, N.; Fujisawa, K.; Moro-oka, Y. J. Am. Chem. Soc. 1989, 111, 8975. 15. Kitajima, N.; Koda, T.; Iwata, Y.; Moro-oka, Y. J. Am. Chem. Soc. 1990, 112, 8833. 16. Hikichi, S.; Komatsuzaki, H.; Kitajima, N.; Akita, M.; Mukai, M.; Kitagawa, T.; Morooka, Y. Inorg. Chem. 1997, 36, 266. 17. Halfen, J. A.; Mahapatra, S.; Wilkinson, E. C.; Kaderli, S.; Young, V. G.; Que, Jr.,L.; Zuberbtihler, A. D.; Tolman, W. B. Science 1996, 271, 1397. 18. Mahapatra, S.; Halfen, J. A.; Wilkinson, E. C.; Pan, G.; Wang, X.; Young, Jr.,V.G.; Cramer, C. J.; Que, Jr.,L.; Tolman, W. B. J. Am. Chem. Soc. 1996, 118, 11555. 19. Shu, L.; Nesheim, J. C.; Kauffmann, K.; Miinck, E.; Lipscomb, J. D.; Que, Jr.,L. Science 1997, 275, 515. 20. Reinaud, O. M.; Yap, G. P. A.; Rheingold, A. L.; Theopold, K. H. Angew. Chem., Int. Ed. Engl. 1995, 34, 2051. 21. Kwart, H. Acc. Chem. Res. 1982, 15, 401. 22. Nesheim, J. C.; Lipscomb, J. D. Biochemistry 1996, 35, 10240. 23. Mahapatra, S.; Halfen, J. A.; Tolman, W. B. J. Am. Chem. Soc. 1996, 118, 11575.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1089

D e s i g n i n g Industrial R e d o x C a t a l y s t s for Selective A u t o x i d a t i o n s o f Hydrocarbons - A New Paradigm Divakar Masilamani Research and Technology, AlliedSignal Inc., P.O.Box 1021, Building CRL-150A, 101 Columbia Road, Morristown, NJ 07962-1021 1. INTRODUCTION Bulk chemicals are an important class of compounds that link the hydrocarbon feedstock to the finished products. Many of these bulk chemicals are produced by oxidizing hydrocarbons. Current processes for hydrocarbon oxidations are inefficient forming a broad spectrum of products requiring costly separations, recycle and waste disposal. Retro-fitting the current operating facilities to control pollution will only add to the processing cost without economic benefits. Modifying the process parameters will at best provide an incremental improvement in efficiency which is not sufficient to enhance the profit margin. 1.1 Idea Our idea is to develop large-scale catalytic processes which will use oxygen or air to convert hydrocarbons efficiently and selectively to oxygenated products. The process will be run under relatively mild conditions in a continuous mode, avoiding the use of solvents. High product-selectivity and atom-utilization will minimize waste and disposal cost. High process efficiency will cut operating cost. The new processes will be safe and profitable at the same time. 1.2 Rationale Many oxidation processes still use stoichiometric reagents such as permanganates, dichromates etc. Such processes produce more waste than product. "E-factor" (which is defined as Kg of waste / Kg of product) is often used to compare wastes 1 generated by different processes. Those that use stoichiometric reagents show high "E-factor" ranging from 25 - 100, while those using catalysts range between 0.1 - 5. Chemical processes must therefore avoid using stoichiometric reagents and switch to catalytic approach to minimize waste.

1090 2. SELECTIVE AUTOXIDATIONS OF CYCLOHEXANE Selective conversion of hydrocarbons to alcohols and ketones has been a problem for chemists for more than a century. Cyclohexane was selected as a model hydrocarbon for our study for two reasons. First, it is a symmetric hydrocarbon and oxidation leads to a single isomer. Secondly, the oxidation products (cyclohexanol and cyclohexanone) are important bulk chemicals used extensively for producing a large number of specialty chemicals and polymers. For example, cyclohexanone is a key intermediate in the production of nylon 6. AlliedSignal produces 500 million pounds of cyclohexanone every year.

2.1 Current technology for producing cyclohexanone Several approaches are currently used to produce cyclohexanone. Two autoxidative processes which are relevant to this presentation will be discussed here. First, we will discuss the AlliedSignal process. In this approach, cumene (isopropylbenzene) is reacted with air at 1050 at 5 psig to form cumene hydroperoxide. The conversion is slow but the selectivity is high approaching 100%. The hydroperoxide when heated at 770 with sulfuric acid, rearranges to form equimolar amounts of phenol and acetone. Phenol is then reduced at 1550 by hydrogen at 80 - 220 psig in the presence of Pd on carbon to form cyclohexanone. About 10% of cyclohexanol is also formed. The direct conversion of benzene to phenol is seldom selective. The route through cumene provides the needed selectivity. But it comes with a price. It is a three-step process starting from cumene. It produces an equimolar quantities of acetone as a byproduct. The DSM process (which is more commonly practiced) converts cyclohexane to cyclohexanone in two steps. In the first step, cyclohexane is contacted with air at high temperatures (165 ~ and pressures (45 - 150 psig). A variety of radical initiators are used as catalysts. The process require a residence time of 5.5 hours. The conversions are usually low (4 - 23%). The major products are cyclohexanol and cyclohexanone (formed in 2:1 ratio) along with a large number of byproducts. (At higher conversions, the byproducts accumulate at the expense of cyclohexanol and cyclohexanone.) The unreacted cycohexane is separated from the products by fractional distillation and recycled. In the second step, the cyclohexanol/cyclohexanone fraction is dehydrogenated between 2500 - 4000 in the presence of a catalyst to produce cyclohexanone in 75% yield. Both AlliedSignal and DSM processes involve free radical chain mechanism. In the case of cumene, the tertiary free radical generated is stable and therefore easy to control at lower temperatures. The selectivity is high. Cyclohexane on the other hand, produces a secondary free radical at higher temperatures. Once produced, they are highly reactive and selectivity suffers and large number of byproducts are formed. An ideal approach to selective hydorcarbon autoxidation is to avoid free radical processes altogether.

2.2 Biological oxidaion of cyclohexane Biological systems are extremely efficient and selective in converting hydrocarbons to alcohols. In the human liver, for example, the P-450 enzyme converts cyclohexane and oxygen to cyclohexanol and water selectively in 100% yield at ambient temperatures and

1091 pressures. The catalyst in this system is an iron (Fe III) porphyrin which by itself is inactive towards oxygen and cyclohexane. However, when reduced to the Fe Hstate by electrons, it reacts with oxygen and cyclohexane to form cyclohexanol and water as shown in Equation 1. C6H12 + 02 + 2e- + 2H + Cyclohexane

> C6HllOH + H20 Cyclohexanol

(1)

A continuous supply of electrons and protons to the catalyst makes the biological oxidation highly selective and efficient. The electrons are provided by nucleosides such as NADH and NADPH. The nucleosides donate a pair of electrons to co-factors such as flavins. The flavins are 2 electrordl electron switches that transfer electrons one at a time to the Fe llI porphyrin catalyst which is enclosed in a protein matrix. The reason why the P-450 system is so selective, efficient and elegant is due to its ability to cleave the O-O bond of oxygen. On receiving an electron, the Fe III porphyrin is reduced to its Fe H state. The catalyst is now activated and binds an oxygen and a proton to form FeW-O-OH. On receiving a second electron and a second proton, it is converted to FeV=O (the "oxo" species) and water. The O-O bond has been broken. The heat of formation of water (54.63 kcal/mol) provides the necessary energy to break the O-O bond. The FeV=O species is in essence an atomic oxygen (with 6 electrons) partially stabilized by iron in various oxidation states 2 as shown in Equation 2 below. FeV=O

~

FeIV-o.

~

Fe III [.O. ]

(2)

The "oxo" species abstracts a hydrogen from cyclohexane to form FeIV-oH and cyclohexyl radical which rebounds at rates 3 as high as 101~ to form cyclohexanol and Fem porphyrin is regenerated 4. Even though free radicals are involved, the oxidation is not a chain reaction and it does not involve alkyl peroxide radicals as in the commercial processes described earlier. Though elegant, biological oxidations occur within the cell in micromolar quantities in aqueous medium. They are too complex and totally unsuited for large-scale industrial productions. 3. DESIGNING NOVEL INDUSTRIAL REDOX CATALYSTS

3.1 Approach Our approach is to adopt the beautiful chemistry of P-450 in an innovative way to large-scale autoxidation of hydrocarbons using conventional reactors. This would require that we design a totally new composite catalytic system and identify inexpensive electron sources that will activate the catalysts continuously during autoxidations. In other words, we are searching for a new paradigm in catalyst design.

1092 3.2 Objective The work presented here does not provide a specific approach to large-scale selective oxidation of cyclohexane. The objective is to probe several new approaches to hydrocarbon oxidations building on the collective experiences of chemists and chemical engineers with large-scale industrial processing and adopt the elegant chemistries of biological oxidations. Technology is currently available for immobilizing appropriate microbes on solid supports and use them in conventional reactors to convert hydrocarbons to their oxidized products. While such processes are highly selective, they would need an aqueous medium in which the hydrocarbon is present at high dilution. Further, the microbes may have to be replaced every few weeks. Such a process will not be efficient. Scale is an important issue in the manufacture of bulk chemicals. It is the Achilles heel in biochemical processing. Our approach is to adopt conventional large-scale reactors (preferably operating in a continuous mode) and use novel sturdy composite catalytic systems which are different from biological catalysts but will still carry out the elegant chemistries of the P-450 enzyme. This approach will represent the launching of an evolutionary process to achieve the next level of sophistication in catalyst design. When we study and generalize the P-450 oxidations, three important issues stand out. (1) We need an inexpensive electron source for continuously activating the redox catalyst; (2) we need a redox transition metal catalyst capable of forming the "oxo" species and (3) we need an electron mediator, usually selected from noble metals, capable of transferring electrons from its source to the redox catalyst. These three issues are critical in adopting biological oxidation chemistries to large-scale hydrocarbon autoxidations. 3.3 Electron sources and reactor configurations

The electron source will dictate the reactor configuration. In equation (1), the two electrons and protons (needed for activating the catalyst) are equivalent to one mole of hydrogen. (Equation 3.) 2H + + 2e

)

(3)

H2

The two protons, the two electrons and oxygen taken together are equivalent to a mole of hydrogen peroxide. (Equation 4.) 2H + + 2e- + 0 2

)

H202

(4)

Hydrogen is the least expensive and highly concentrated source of electrons and protons. However, it forms an explosive mixture with oxygen and therefore is not safe to be mixed with it. Hydrogen peroxide and its related reagents such as HOC1 and t-butylhydroperoxide are convenient to use. More importantly, they help to overcome oxygen transport problems inherent in autoxidations. Further, peroxides are known to convert redox catalysts such as Fe +3, Mn +3, Ti +4 etc directly to their "oxo" species. The electron mediator is no longer needed. Using peroxides, simplifies the design of the reactor.

1093 Isopropanol (and other secondary alcohols) are also excellent electron sources when used with Ru, Pd and Rh as electron mediators. After donating two electrons and two protons (ie. a hydrogen equivalent) the secondary alcohols are converted into ketones. The ketones are readily reduced back to the secondary alcohols by catalytic hydrogenation and recycled. Using secondary alcohols will therefore add an extra step to the process but they are safe in the presence of oxygen. Thus secondary alcohols can be considered as safe sources of hydrogen. 3.4 Catalytic systems

A composite catalytic system will consist of a redox transition metal catalyst and an electron mediator. The redox catalyst is selected from metal ions such as V +5, Ti +4, Mn+3 Fe +3 etc based on their ability to form the "oxo" species. The electron mediator is selected from Ru, Rh and Pd based on their well established ability to remove and transport electrons from hydrogen and secondary alcohols. 4. RESULTS 4.1 Extracting electrons from hydrogen with catalytic membranes

We developed a membrane system to demonstrate the feasibility of using hydrogen as an electron source and at the same time preventing it from coming in contact with oxygen. The membrane was prepared by mixing a solution of silicone rubbers with 5% Pd on carbon and spreading this mixture over a polysulfone membrane which is not permeable to hydrogen but will allow electrons and protons to pass through. The coated membrane was cured at 100 ~for an hour. Hydrogen at 10 psig pressure was converted to electron and protons and transferred to a buffer solution (at pH 8.5) containing methylene blue on the other side of the membrane. The methylene blue was reduced to its colorless leuco form. On exposing to oxygen, the blue dye was regenerated and oxygen was converted to hydrogen peroxide. However, Mn +3 tetraphenylporphyrin was reduced preferentially over oxygen to its Mn +2 form as evidenced by UV spectrum. These experiments show that Pd is capable of extracting electrons from hydrogen and transporting them through the membrane without hydrogen contacting the redox catalyst or oxygen on the other side of the membrane. However, the membranes developed leaks over a period of a week and they were not compatible with organic solvents. 4.2 Peroxides as source of electrons and oxygen

Peroxides are easy to handle. A large number of papers report using hydrogen peroxide, HOC1 and t-butylhydroperoxide in hydrocarbon oxidations. As discussed earlier, the peroxides are sources of electrons, protons and oxygen circumventing the need for an electron mediators. The peroxides directly convert redox catalysts to their the "oxo" species. We used 30% hydrogen peroxide to oxidize cyclohexane (50 mM) dissolved in dichloromethane containing 5 mM Mn +3 tetraphenylporphyrin as the redox catalyst. The biphasic system was stirred for 24 hours at room temperature. Cyclohexane was converted

1094 to 48% of cyclohexanol and 10% of cyclohexanone. The catalyst was completely destroyed during the course of the reaction. t-Butylhydroperoxide was more effective since it is soluble in organic solvents. A mixture (5 mL) of cyclohexane and benzene (80:20 volume ratio) containing 1 mM Mn +3 porphyrin was stirred at ambient temperatures. Different amounts of t-butylhydroperoxide were added slowly over several hours using a syringe pump. The products were analyzed after 24 hours. About 60 - 80% peroxide was consumed. Cyclohexanol and cyclohexanone were formed in 9:1 ratio along with t-butanol. Trace amounts of side products were also identified by gc/MS to be cyclohexyl t-butylperoxide, cyclohexyl t-butylether and di tbutylperoxide. Even though the reaction is selective at ambient temperatures, it is slow and separation of the catalyst from the reaction mixture was difficult. Further, about 20 - 25% of the catalyst was destroyed in 24 hours. The recovery of t-butanol and its conversion to tbutylhydroperoxide is a costly operation.

4.3 Use of secondary alcohols as electron donors Secondary alcohols are excellent hydrogen donors. In the0Presence of RuH2(PPh3) 3 secondary alcohols such as isopropanol produce hydrogen 5 at 150 with high turn-over numbers. In this process, the alcohols are quantitatively converted to the ketones. Ru, Rh and Pd catalysts are excellent in "transfer hydrogenation" reactions using secondary alcohols as hydrogen sources to reduce other substrates such as the imines. We were convinced that the above catalysts are also potential electron mediators in transferring electrons and protons from secondary alcohols to the redox catalyst. We were able to demonstrate that Pd and Ru are excellent electron mediators that converted cyclohexanol (a secondary alcohol) in 98% selectivity to cyclohexanone. Four electron mediator catalysts (Pd on carbon, PdC12, RuEC1E(p-cymene)2 and a trinuclear Ru carboxylate) were tested. The conversions were poor with Pd catalysts even though the selectivities were high. The RuEC12(p-Cymene)2 was more efficient 6. However, it required MnO 2 as an electron sink. THF was used as the solvent. Cyclohexanone was formed in 36% yield in 60 hours at 70 ~ The most promising results were obtained with the trinuclear Ru carboxylate prepared by Wilkinson and used by Drago 7. Cyclohexanone was formed in 60% yield in 6 hours at 400 in 98% selectivity. In this reaction, cyclohexanol was used neat and oxygen was used as the electron sink and was converted to hydrogen peroxide. This was a surprising result since Tang s has shown that the oxidation of cyclohexanol by oxygen in the presence of RuC13 is extremely slow.

4.4 Composite redox catalysts for selective hydrocarbon autoxidations Titanium silicalite has been recognized as an efficient redox catalyst in a number of industrial processes. Enichem has a process in which Ti silicalite catalyzed the conversion of cyclohexanone by ammonia and hydrogen peroxide to cyclohexnone oxime 9. The mechanism appears to involve the "Ti oxo" species. A number of procedures are available from literature to make the Ti silicalite catalyst and it will be easy to incorporate electron mediators such as Ru or Pd into the silicalite matrix. The performance of such composite

1095 catalysts in selective autoxidations can be evaluated using cyclohexane containing a small amount of cyclohexanol. The autoxidation will be initiated by cyclohexanol by transfering its electrons to activate the composite catalyst. In this step, cyclohexanol will be converted to cyclohexanone. The activated catalyst will react with oxygen and cyclohexane to form water and generate more cyclohexanol thereby starting a self-perpetuating reaction sequence in which the net reaction is the conversion of cyclohexane to cyclohexanone and water occurring in the same reactor. The reaction sequence is summarized in Scheme 1 below. Scheme 1 Cyclohexane autoxidation Step 1. C6HllOH Cyclohexanol Step 2.

Net reaction

C6H100 + H2/Catalyst Cyclohexanone

H2/Catalyst + 02 + C6H12 ~ Cyclohexane C6H12 + 02 Cyclohexane

C6H11OH + H20

) C6H100 + H20 Cyclohexanone

AGo= +6.49 AS~ 29.89 (kcal/mol) ( eu ) AGo= -90.41 AS~ -28.06 (kcal/mol) (eu) AG~

AS~ (kcal/mol)

1.83 (eu)

In Scheme 1, the free energy and entropy for step 1, for step 2 and for the net reaction were calculated from thermodynamic data l~ It is clear that step 1 is endothermic with a positive entropy. Higher temperatures will favor this step. Step 2 is overwhelmingly exothermic; but the entropy is unfavorable. Higher temperatures will affect this step only slightly. The entropies of these two steps counteract each other and as a result the free energy of the net reaction will not change much with temperature. At 25 o and 150 ~ the free energies for the net reaction are respectively -83.92 kcal/mol and -81.83 kcal/mol. Based on this analysis, we decided to run the reaction at 150~ Clerici and Bellussi 11 have shown that hexane in methanol can be selectively oxidized to 2-hexanol, 3-hexanol, 2-hexanone and 3-hexanone using a mixture of oxygen and hydrogen at 25 o - 30 ~ The reactions were run in the presence of HC1 for 20 - 24 hours. Several titanium silicalite catalysts containing Pd (0.01 mol ratio to TiO2) were prepared and used in these reactions. Presumably, hydrogen plays the role of an electron-donor activating the Ti catalyst. However, no explanations were offered. We prepared composite catalysts containing Ru in Ti silicalite (molar composition SiO2:0.025TIO2 : 0.001 Ru). The catalyst was suspended in 3 mL of a mixture (1:4 weight ratio) of cyclohexanol and cyclohexane in a 100 mL steel autoclave and stirred under oxygen at 150 o for 24 hours. Cyclohexane was converted to cyclohexanone to an extent of 15 - 20%. The reaction was highly selective for cyclohexanone. Products other than unreacted cyclohexane and cyclohexanol were not observed. More than the calculated amount of oxygen was consumed. The experimental set up was similar to the one described by Clerici and Bellusi ll.

1096 The use of secondary alcohol for activating redox composite catalysts in hydrocarbon autoxidations is attractive. Isopropanol is excellent as an electron source. Acetone will be produced as the byproduct in these autoxidations. It is easily separated and reduced back to isopropanol by catalytic hydrogenation and recycled. Several composite redox catalysts must be prepared with various combinations and ratios of electron mediators (such as Ru, Pd, Rh etc) with redox catalysts (such as V +5. Ti § Mn +3 , Fe+3 etc) and evaluated for their ability to selectively oxidize hydrocarbons. 5. C O N C L U S I O N A new paradigm in the design of catalysts for selective autoxidation of hydrocarbons is presented. The catalyst will consist of two components: a redox component and an electron mediator. The redox component is selected from transition metal ions. These ions are capable of breaking the O-O bond of oxygen to form the "oxo" species provided they are continuously activated by a supply of electrons. The second component of the catalyst is an electron mediator (selected from noble metals) which help to transfer electrons from a suitable source to the redox component. Hydrogen and secondary alcohols are the best sources of electrons. For safety reasons, the mixing of hydrogen and oxygen is not desirable. This complicates the reactor design. However, secondary alcohols are convenient and safe to use. Buttheir regeneration will add an extra step. Peroxides can serve as sources of electrons, protons and oxygen. Further, they help to overcome oxygen transport problems inherent in autoxidations. However, peroxides also destroy the catalysts particularly at higher temperatures. We have shown that cyclohexane can be selectively converted to cyclohexanone. In this reaction, the intermediate (cyclohexanol) serves as the electron source. What is important is that this selective oxidation was achieved in a conventional batch reactor. We have a long way to go. But we have made the start. We have managed to "lift" the elegant chemistry of the P-450 enzymes out of the confines of biological cells into a larger arena of conventional man-made reactors. REFERENCES

1. R.A.Sheldon, Chemistry & Industry (1992) 903. 2. D.H.R.Barton, F.Halley, N.Ozbalik and E.Young, New J. Chem. No.3 (1989) 177. 3. V.Bowry, J.Lusztyk and K.U.Ingold, Pure & Appl. Chem. Vol. 62, No.2 (1990) 213. 4. M.Hirobe, Pure & Appl. Chem. Vol. 66, No.4(1994) 729. 5. J-E.Backvall, R.L.Chowdhury and U.Karlsson, J.C.S., Chem.Comm. (1991) 473. 6. U. Karlsson, G-Z.Wang and J-E. Backvall, J.Org.Chem.Vol.59 (1994) 1196. 7. C.Bilgrien, S.Davis and R.S.Drago, J.Am.Chem.Soc. Vol. 109 (1987) 3786. 8. R.Tang, S.E.Diamond, N.Neary and F.Mares, J.C.S. Chem.Comm. (1978) 562. 9. S.Tonti, P.Roffia and V.Gerasutti, Ammoimation, US Patent No.5 227 525 (1993). 10. D.R.Stull, E.F.Westrum and G.C.Sinke, The Thermodynamics of Organic compounds, Wiley, New York, 1969. 11. M.G.Clerici and G.Bellussi, Oxidating Paraffins, US Patent No. 5 235 111 (1993)

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1097

Selectivity of Active Sites on Oxide Catalysts C. Batiot, F.E. Cassidy, A.M. Doyle and B.I~ Hodnett Dept. of Chemical and Environmental Sciences, University of Limerick, IRELAND.

ABSTRACT The selectivity of active sites on oxide catalysts have been assessed by comparing their ability to selectively activate a C-H bond in a reactant rather than a similar C-H or C-C bond in a selective oxidation product. Active sites on oxide catalysts are capable of activating target bonds in reactants that are up to 30-40 kJ mole 1 weaker than similar bonds in the selective oxidation product. Good selectivities are always recorded provided that selective oxidation reactions attempted do not exceed the discriminating capacity of the active site. Evidence is also presented that C-C bonds, which are generally weaker than C-H bonds, are protected from rupture by steric factors. 1. INTRODUCTION The essential value of a selective oxidation catalyst can be represented by a simple selectivity-conversion plot, the best catalysts being those that give the highest selectivity at a given conversion [1]. Other factors such as activity and deactivation are less important, because the former can always be boosted by increasing the W/F ratio or the temperature and generally deactivation phenomena in oxidation catalysis are not severe. Htmdreds of examples of selective catalytic oxidation have appeared in the open and patent literature, but to date we do not know the factors which determine the upper limits of selectivity that can be attained for a given reaction. One factor which is well known is that a given selective oxidation catalyst has to be tailored for a specific reaction and generally a great deal of effort is required, in the first instance by way of screening a wide range of materials and, secondly, in fine tuning the best of these with additives, promoters, dopants and supports to arrive at the best catalyst for a particular reaction and reactor configuration. The concept of active sites on heterogeneous catalysts was first introduced by H.S. Taylor [2] and has been widely used since, although not always fully defined. A meaningful assessment of the selectivity of active sites on oxide catalysts should be carried out on a comparative basis. The ultimate goal in the design of heterogeneous catalysts is to attain a specificity or selectivity typically achieved by enzymes. Hence figure 1 presents a comparison of enzymes and heterogeneous catalysts on the basis of four criteria common to both: These are (i) operating temperature, (ii) molecular mass of the active site, (iii) turnover frequency and (iv) selectivity. The normal definitions of terms (i), (iii) and (iv) have been used for figure 1 and

1098 the molecular mass of the active site for the enzyme was estimated from molecular mass of the enzyme, assuming just one active site per enzyme molecule [3]. The lower limit of active site size for heterogeneous catalysts has been estimated at the size of a single metal atom on a support whereas the upper limit is calculated from the concentration of acid sites in high silica to alumina ratio zeolites. This measure is intended to determine the number of active sites per unit volume that could be fired into a reactor so that the contribution by the support or the underlying bulk is taken into account in making the calculation. Realistically, on this basis, the active site size range in oxide catalysts lies in the range 103-104 amu, the lower value referring to unsupported catalysts and the latter to supported or diluted systems. Temperature (~

Selectivity (%)

Molecular Mass of Active site (Daltons)

Turnover Frequency

(s-b

106

100

1000

200,000 95

95 60,000

100

102

20,000

37 30 m

10.4

-20

10-4

100 H.C.

ENZ

H.C.

ENZ

H.C.

ENZ

H.C.

ENZ

Figure 1: Comparison of active sites in hetergeneous and enzymatic catalysis. (H.C. = Heterogeneous catalysis; ENZ = Enzymatic catalysis). The general overview presented in figure 1 is sufficient to indicate that in the majority of cases active sites on heterogeneous catalysts can be bigger than conventionally represented and overlap to some extent the size of corresponding sites in enzymes. The normal operating temperature range for enzymes is restricted by comparison with heterogeneous catalysts but the turnover frequencies that can be achieved on enzymes far outperforms that achieved by heterogeneous catalysts. In addition the selectivity achieved is normally very dose to 100% for enzymes and in addition the transformations that can be achieved are chemically much more complex than over heterogeneous catalysts [3]. Literature descriptions of active sites on oxide catalysts are oiten speculative and very often just generate a picture of the surface active site by extrapolation of the bulk structure. In general they envisage approach of the starting material to the active site in a preferred orientation without any indication of how the preferred orientation is established. In addition, the description of the active site is usually restricted to a small number of molecules. For example, the vanadium phosphorus oxide catalysts used for n-butane oxidation to maleic anhydride is based on the vanadyl pyrophosphate structure and an active site architecture is

1099 often presented which involves just two vanadium ions surrounded by a first coordination sphere of oxygen anions, without si,~nificant reference to the role of the pyrophosphate species [4]. Other active sites on oxide catalysts have been described where the site is strongly diluted on a molecular scale by a non-reactive species [5]. The central question is how selective can we expect heterogeneous catalysts to perform in the light of their rather simple chemical nature and the simple chemical transformations that they are designed to mediate. In the past attempts have been make to determine genetic factors which appear to be important in all selective oxidation reactions. Some of these factors are dearly related to catalyst structure, such as the requirement to dilute the active site at the surface and the primal role of lattice oxygen. Other factors such as the need to kinetically isolate the reaction products relate more to reactor engineering [5]. This work concentrates on another genetic factor which is related to the structure of the species appearing in the gas phase, emphasising the role of substrate structure, particularly the bond dissociation energies of the weakest C-H bonds in the reactants and products in determining selectivity in oxidation and ammoxidation reactions [ 1]. 2. RESULTS AND DISCUSSION Most selective oxidation reactions can be treated on the basis of a sirnple kinetic scheme as follows: Scheme I

HC

kl

~-

S.O.P

k2

=

COx

t A basic operating principle of selective oxidation catalysis is the need to minimise the contact time between the selective oxidation product and the catalyst to prevent conversion of the product, typically, into oxides of carbon. Whereas this aspect of selective oxidation catalysis is well recognised, it has never been put on a quantitative basis, so that the ability of a particular active site to activate a target bond in a reactant in preference to a similar bond in the product. Two aspects of scheme 1 will be addressed here, namely the factors which determine the limiting selectivity in terms of the ratio of k l to k2 and secondly the factors which determine the limiting selectivity in terms of the ratio of kl to k3. The selectivity at 30% conversion will be taken as measure of k l to k2 and the selectivity at zero to 10% conversion will be taken as a measure o f k l to k3. Towards these ends 14 selective oxidation reactions and two ammoxidation reactions have been evaluated through the use of selectivity-conversion plots, constructed from literature data [ 1]. Two examples of these plots are presented in figure 2 for ethylbenzene oxidation to styrene and methane oxidation to ethane. These selectivity-conversion plots were generated for a variety of catalysts for each reaction over a range of temperatures and space velocities. It should be stressed that the objective of this exercise was not to determine a reaction pathway or network, but simply to evaluate the best performance which has been achieved for any given reaction, hence the use of data from different catalysts and operating conditions.

1100

ethylbenzene - - > styrene 100 4: ' " ' ~

methane ---> ethane 100

"- ......

80 60 40

sel

,."

20

.,..:~.:~.%.'~

2o

0

0

t

I

I

I

20

40

60

80

100

0

20

40

60

80

100

conversion (o~

conversion (%)

Figure 2: Selectivity-conversion plots for the oxidation of ethylbenzene to styrene [6,7] and methane to ethane [8-25]. In all cases studied an upper limit could be identified beyond which experimental studies have not yet progressed and data points which fall below the upper limit are assumed to arise from operation with poor catalysts or in non-optimised conditions [ 1]. The first part of our analysis concentrates on the properties of the species appearing in the vapour phase, namely the reactants and the selective oxidation products and how well an active site distinguishes between these two species. A striking example is the oxidation of methane to ethane [8-25]. The structures of reactant and product are very similar as are their reactivities. Here we address the basis on which active sites typically present on oxide catalysts distinguish between almost identical bonds in reactants and selective oxidation products [26]. Generally selectivity in oxidation catalysis involves activation of the reactant through rupture of a C-H bond (the k l route in scheme 1), whereas diminishing selectivity is associated with rupture of any bond in the selective oxidation product (the k2 route in scheme 1). As a means of validating this hypothesis the upper selectivity limit, attained at a fixed conversion, in all 14 reactions used in this study was plotted against the function:

D~ where D~

(reactant)- D~

or C-C (product)

(reactant) is the bond dissociation enthalpy of the weakest C-H bond in the

substrate and D~

or C-C (product) is the bond dissociation enthalpy of the weakest bond

in the selective oxidation product.

1101

1

1cu3

.It 2

9e e ~ - " ~

.....

9 84 ,

80 9

A

oo

z=

11

40

2 20

t -70

-120

t

0

i

-20

13 ~ 1 4 ....l

30

D*Hc.I.I r,,,,~to,t- D ' H e . ,

130

80

0~ c-c ~..a,,.-t ( k J I m o l )

Figure 3 : Selectivity in product versus D~ C-H reactant" D~ C-H or C-C product at 30% conversion. 1, Ethylbenzene to Styrene. 2, 1-Butene to Butadiene. 3, Acrolein to Acrylic Acid. 4, Ethane to Ethylene. 5, n-Butane to Maleic Anhydride. 6, Propene to Acrolein. 7, Methanol to Formaldehyde. 8, Ethanol to Acetaldehyde. 9, Propane to Propene. 10, n-Butane to Butenes. 11, Propane to Acrolein. 12, Methane to Ethane. 13, Ethane to Acetaldehyde. 14, Methane to Formaldehyde [ 1]. The observed correlation (Figure 3) shows that there is a clear relationship between limiting selectivities and the nature of the weakest C-H or C-C bonds in the reactants and products. Active sites in oxide catalysts are capable of activating target bonds in the reactant that are up to 30-40 kJ mole 1 weaker than similar bonds in the selective oxidation product. When bigger differences in bond energies arise drastic reductions in selectivities result because the discriminating capacity of the active site has been exceeded. Scheme 2 presents a sequence of reactions starting with propane and leading to propene and acrolein. In the figure the percentage cited under each arrow is the limiting selectivity that has been reported in the literature for the reaction in question at 30% conversion. These data were generated on the same basis as those used for figures 2 and 3 above. The numbers above the arrows are the values which apply for the function: D~ (reactant) - D~

or C-C (product). Scheme 2

I [ [ ~C~C~C

I I I I

41 ldmole-1 >

14 ldmole-1

700/o --i--\

/

55 kJmole-1 40%

C~C~C~O

I

l

1102 This coherent reaction network clearly demonstrates the importance of the 30-40 kJ mole -] selectivity limit. When it is exceeded, as is the case with propane oxidation to acrolein, selectivity declines drastically. Similarly the accummulated data for propane and propene ammoxidation [27,28] to acrylonitrile indicate selectivities at 30% conversion of 50% and 85% respectively. These data are consistent with the 41 kJ mole .] difference in bond enthalpies shown in scheme 2 for propane and propene. Having established the discriminating limits of typical active sites in selective oxidation catalysts the second question to be addressed is the factors which determine the ratio of kl to k3 in scheme 1. The same methodology was followed here except that selectivities at 10% conversion were taken as a measure o f k l to k3. The basic approach is to assume that the kl route is favoured by activation or rupture of a C-H bond in the reactant whereas the k3 route is favoured through activation or rupture of a C-C bond. This part of the analysis will be restricted to the oxidative dehydrogenation of ethane, propane and n-butane. The corresponding selectivity conversion plots are presented in figure 4.

ethane ---> ethylene

propane-->

propene

100 9

60

9

9

".j

"

9 9 9

,,9 , . . "e

40

t

Oo 9 o

*,

9

"

9

sel 20

9

9

"#

"

40 : ' ~ r . ' , ' ~ . . .

"

""

sel 20 ~#~" : ee

0 I

I

I

I

t

0

20

40

60

8O

100

conversion (%)

0

.

o~

9

I

~

I

t

20

40

60

80

100

conversion (%)

n-butane ---> butenes

80 :':" 60

9 ., "

40

" .~;:..--

"" . " ' . sel

20

0

.

:~ e .

9t .

0

.

..

-.

**

Figure 4: Selectivity-conversion plots for

9

I

I

I

20

40

60

conversion (%)

100

ethane [29-38], propane [39-56] and n-butane [57-64] oxidative dehydrogenation.

1103 Generally C-C bonds in alkanes are much weaker than the corresponding C-H bonds, as shown in Table 1. Clearly then active sites do not distinguish between these bonds on a bond strength basis only and in view of the structure of these alkanes it is reasonable to suggest that steric factors must play a role here, since C-C bonds are generally more difficult to accommodate within an active site than C-H bonds.

Table 1: Bond Dissociation Enthalpies in Alkanes [26]

DC-HkJ molel(C-H3)

Alkane

Dc.c kJ mole "1

Ethane Propane n-Butane *(The value in brackets

376 420 368 417 360 (343)* 403 refers to the central C-C bond in n-butane)

DC-HkJ molel(C-H2) n/a 401 390

A more reasonable correlation emerges with the selectivity at 0-10% conversion when we examine the ratio of C-C to C-H bonds in a particular alkane, as shown in figure5.

100 --I-90

"-

Selectivity at 0% conversion Selectivity at 5% conversion Selectivity at 10% conversion

80 Sel 70 0

i

I

i

0.1

0.2

0.3

0.4

Ratio C-C to C-H Bonds

Figure 5: Influence of the ratio of C-C to C-H bonds in ethane, propane and n-butane on the selectivity in oxidative dehydrogenation at 0, 5 and 10 % conversion.

1104

3. CONCLUSIONS There was a clear upper limit in terms of selectivity-conversion beyond which experimental studies have not advanced for many selective oxidation reactions. These limits have been achieved through detailed catalyst design and reactor optimization. This work shows that active sites on oxidation and ammoxidation catalysts are capable of selectively activating, typically, a C-H bond in a reactant, rather than a similar C-H or C-C bond in the product provided that the bond dissociation enthalpy of the weakest bond in the product is no more than 30-40 kJ mole~ weaker than the bond dissociation enthalpy of the weakest bond in the reactant. When these limits are exceeded selectivity falls drastically. This work also indicates that primary activation of alkanes is through C-H bonds although the corresponding C-C bonds are much weaker. Cleavage of a C-C bond in the primary activation step leads directly to carbon oxide formation, but this step is less favoured because stedc factors make it diitictflt for the C-C bonds to be accommodated at the active site.

REFERENCES .

2. .

4. 5.

9

10 ll 12 13 14 15 16

C. Batiot and B.I( Hodnett, Appl. Catal. A, 137 (1996) 179. J.M. Thomas and W.J. Thomas, Principles and Practice of Heterogeneous Catalysis, Ed. VCH Weinheim, 1997. D. Voet and J.G. Voet, Biochemistry, J. Wiley & Sons Inc., 1995. B. Schiott and I~A. Jorgensen, Catal. Today, 16 (1993) 79. ILK. Grasselli in Surface Properties and Catalysis by Non-Metals, (Ed. Bonnelle, J.P. et al), D.Reider Publishing Company, Dordrecht, (1983) 273. M. Turco, G. Bagnasko, P. Ciambi, A. La Ginestra and G. Russo, Stud. in Surf. Sci. Catal, 55 (1990) 327. F.M. Bautista, J.M. Campelo, A. Garcia, D. Luna, J.M. Marinas and 1LA. Quiros, Stud. in Surs Sci. Catal, 82 (1994) 759. V.G. Roguleva, M.A. Nikiphorova, N.G. Maksimov and A.G. Anshits, Catal. Today, 13 (1992) 219. ICD. Campbell, Catal. Today, 13 (1992)245. G.J. Tjatjopoulos and I.A. Vasalos, Catal. Today, 13 (1992) 361. J.S.J. Hargreaves, G.J. Hutchings, IEW. Joyner and C.J. Kiely, Catal. Today, 13 (1992) 401. S.J. K o ~ J.A. Roos, J.M. Diphorn, l~H.J. Veehoi~ J.G. Van Ommen and J.1LH. Ross, Catal. Today, 4 (1989) 279. G.J. Hutchings, M.S. Scurrell and J.1L Woodhouse, Catal. Today, 4 (1989)371. C. Chevalier, P. Ramirez, M. Ceruso, A. Choplin and J.M. Basset, CataL Today, 4(1989) 433. A. Kiennemann, 1L Kieffer, A. Kaddouri, P. Poix and J.L. gehspringer, Stud. in Surf Sci. Catal., 55 (1990) 365. A.A. Kadughim~ O.V. Krylov, S.E. Plate, Y.P. Tulenin, V.A. Selezaev, A.V. Bolrov and Y.M. Kimelfeld, Stud. in Surf. Sci. Catal., 55 (1990) 447.

1105 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51

P. Kovacheva, N. Davidova and A.H. Weiss, Stud. in Surf. Sci. fatal., 82 (1994) 403. S.T. Brandao, L. Lietti, P.L. Villa, S. Rossini, A. Santucci, IL Millini~ O. Forlani and D. Sanfilippo, Stud. in Surf. Sci. fatal., 82 (1994) 443. T. LeVan, M. Che, M. Kermarec, C. Louis and J.M. Tatibouet, fatal. Lea., 6 (1990) 395. Choudhary, AICHE J., 37 (1991) 915. F.P. Larkins and M.IL Nordin, J.Catal., 130 (1991) 147. Yang, Bull. Soc. Chim Belg., 100 (1991) 5. Hinson, J.Chem. Soc. Chem_ Commun., (1991) 1430. Kiwi, J. Phys. Chem_, 96 (1992) 3, 1344. K o ~ Thesis Twente (1990). Handbook of Chemistry and Physics 61 st edition, R.C. Weast editor, 1980-1981 A. Corma, J.M. Lopez Nieto, N. Paredes and M. Perez, Appl. fatal., 97 (1993) 159. IL Burch and E.M. Crabb, Appl. fatal., 100 (1993) 111. Y.-C. Kim~ W. Ueda and Y. Moro-oka, Appl. fatal., 70 (1991) 189. ILK. Grasselli and J.L. Callahan, J. fatal., 14 (1969) 93. K. Seshan, H.M. Swaan, ILH.H. Smits, J.G. Van Ommen and J.ILH. Ross, Stud. in Surf. Sci. fatal., 55 (1990) 505. J. Barrault and L. Magaud, Stud. in Surf. Sci. fatal., 82 (1994) 305. I. Matsuura and N. Kimura, Stud. in Surf. Sci. fatal., 82 (1994) 271. B. Grzylowska, P. Mekss, IL Grabowski, K. Wcislo, Y. Barbaux and L. Gengembre, Stud. in Surf. Sci. fatal., 82 (1994) 151. J.G. Eon, P.G. Pries de Oliveira, F. Lefebre and J.C. Volta, Stud. in Surf. Sci. fatal., 82 (1994) 83. P.G. Pries de Oliveira, J.G. Eon and J.C. Volta, J. fatal., 137 (1992) 257. L. Magaud, Thesis Poitiers (1994). H.H. Kung, US Patent No. 4 777 319 (1988). IL Burch and R. Swarnakar, Appl. fatal., 70 (1991) 129. IL Burch and E.M. Crabb, Appl. fatal., 97 (1993) 49. S.S. Hong and J.B. Moffat, Appl. fatal., 109 (1994) 117. J.C. Vedrine, J.C. Le Bars, J.C. Vedrine and A. Auroux, Appl. fatal., 88 (1992) 179. A. Erdohelgi and F. Solymosi, Appl. fatal., 39 (1988) Lll. G.C. Colorio, B. Bonnetot, J.C. Vedrine and A. Auroux, Stud. in Surf. Sci. fatal., 82 (1994) 143. S. Bordoni, F. CasteUani, F. Cavani, F. Trifiro and M.P. Kulkarni, Stud. in Surf. Sci. fatal., 82 (1994) 93. J.G. Mc Carty, A.B. Mc Ewen and M.A. Quinlan, Stud. in Surf. Sci. fatal., 82 (1994) 405. M. Merzauki, B. Taouk, L. Monceaux, E. Bordes and P. Courtine, Stud. in Surf. Sci. fatal., 72 (1992)165. P.M. Michalakos, M.C. Kung, I. Jahan and H.H. Kung, J. fatal., 140 (1993) 226. E. Morales and J.H. Ltmsford, J. fatal., 118 (1989) 255. L. Mendelovici and J.H. Lunsford, J. fatal., 94 (1985) 37. T. Hayakawa, A.G. Andersen, H. Orita, M. Shimizu and K. Takehira, fatal. Lett., 16 (1992) 373.

1106

52 53 54 55 56 57 58 59 60 61 62 63 64

O. Desponds, ILL. Keiski and G.A. Somorjai, Catal. Lett., 19 (1993) 17. L. Mendelovici, Chem_ Lett., (1982) 1469. S. Bordoni and F. Trifiro, J. Che~L Soc. Faraday Trans., 90 (1994)19, 2981. H.M. Swaan, Thesis Twente (1992). A. Kaddouri, European Congress (1993). M.A. Chaar, D. Patti, M.C. Kung and H.H. Kung, J. Catal., 105 (1987) 483. M.A. Chaar, D. Patti and H.H. Kung, J. Catal., 109 (1988) 463. D. Patel, P.J. Andersen and H.H. Kung, J. Catal., 125 (1990) 132. L. Owens and H.H. Kung, J. Catal., 148 (1994) 587. R.G. Rizayev, R.M. Talyshinskii, J.M. Seifitllayeva, E.M. Guseinova,Y.A.Panteleyeva and E.A. Mamedov, Stud. in Surf. Sci. Catal., 82 (1994) 125. A. Corma, J.M. Lopez Nieto, N. Paredes, A. Dejoz and I. Vasquez, Stud. In Surf.Sci.Catal., 82 (1994) 113. H.H. Kung, US Patent No. 4 777 319 (1988). D. Bhattacharyya, S.~ Bej and M.S. Rao, Appl. Catal., 87 (1992) 29.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1107

A novel computer-aided technique for the development of catalysts for propane ammoxidation to acrylonitrile X.-Q. Wu, Q.-X. Zhang, Q.-L. Dai, Z.-Y. Hou and D.-W. Lu Chemical Engineering Department, Zhejiang University Hangzhou 310027, People's Republic of China

In this paper, a new computer-aided technique was presented, with which the experimental procedure of developing catalysts is scheduled sequentially. In each sequential step the neural networks model and multi-objective optimization are used to determine optimal design for the next experiment. The sequential method proved very efficient in developing catalysts for propane ammoxidation to acrylonitrile. And the yield of acrylonitrile corresponding to the best catalyst was up to 58.9%.

1. INTRODUCTION The work of developing multicomponent catalysts, such as catalysts for selective oxidation of hydrocarbon, is very complicated, which often takes much time and spending. Conventionally, developing this kind of catalysts has been a repetitious trial-and-error procedure under the directions of qualitative analyses and experiences. To a multicomponent catalyst, many factors, e.g., preparation conditions, active components, structure and reaction conditions will influence its catalytic performance. This influence behaves generally in a synergistic way. The conventional method or trial-and-error procedure can not disclose and describe such a complicated influential relationship, so the efficiency ofdeveloping catalysts is low. In order to improve the developmental method, researchers have introduced computer-aided means, e.g., artificial intelligence, optimization and graphics, etc., and the usefulness of the means has been proved in some cases [ 1-3]. The process for direct acrylonitrile (ACN) synthesis from propane ammoxidation: C3H8 + NH3 + 202 > CH2CHCN+4H20 AH= - 151.13 kcal/mol (1) with its commercial potentiality, has been widely investigated in the recent years. Various catalytic systems have been reported for the reaction, yet the selectivity and yield of acrylonitrile were unsatisfactory. Moreover, it seems that V-Sb-A1 mixed oxides represent the most promising candidate components for the catalysts because of their ability to selectively convert an alkane to an unsaturated nitrile [4-6].

This work is partially supported by the National Natural Science Foundation of China.

1108 In the paper, we presented a new computer-aided technique, which takes the catalyst development as an iterative procedure, and each iteration includes four steps, i.e. distributing experimental points, carrying out experiments, modeling the catalytic relationship and forecasting the optimal design. The technique was applied to developing catalysts for propane ammoxidation to acrylonitrile.

2. METHODOLOGY

The key to improve the method of developing catalysts is to set up some quantitative catalytic relationships, with which one can make the developmental procedure become a combinational one of qualitative analyses, quantitative predictions and experiments. For some simple catalytic systems, without any pre-determined experiment, quantum chemistry can be used to estimate catalytic properties quantitatively. Unfortunately, up to now it is difficult to apply this method to the multicomponent catalytic systems. A certain amount of experiments is necessary to develop the multicomponent catalysts. But, how to decrease the amount of experiments noticeably is a matter to which researchers have paid great attention, it is also the problem to be solved in this paper. Hence, we proposed a novel computer-aided technique, by which the procedure of developing catalysts is transferred to an iterative or sequential one. 2.1. Principle of the technique In the technique, the orthogonal design, neural networks(NN) and multi-objective optimization are adopted to construct the iterative procedure. The principle and steps of the technique are detailed below:

1). Adopting orthogonal design to distribute a number of exploratory (or original) points evenly within the whole experimental space; 2). Performing the experiment at each located point and getting the data; 3). Using all current data (including new and old data) to compose the sample set for NN training and training NN to obtain the catalytic relationship model; 4). Carrying out multi-objective optimization based on the trained NN model to forecast optima (if exist several local optima); 5). Calculating the iterative precision with the predicted and tested results; if the precision unsatisfied, taking all optima as new experimental points, then turn to step 2; 6). If the precision satisfied, giving out the total optimum and stop the iterative procedure. Here the iterative precision is given as: = lira,-r

.ll/llr

.ll

(2)

where Yopt and Yoxp are the results of catalytic performance obtained by the optimization and experiment respectively. The diagram of the iterative procedure is shown in Figure 1. According to the strategy of the iterative procedure, it can be expected that, with the iteration

1109 going on, the precision of the NN model and optimization will be raised gradually. Especially, the precision in the sensitive regions will be raised more rapidly than others. 2.2. Distribution of the original experimental points It is important to determine the experimental space, since the prediction ability of the presented technique is limited to the space. If some factors are not included in the space, their effects will not exist in the relationship models. For a catalytic system, essential influential factors can be selected based onthe theoretical and empirical knowledge as well as the results of literature. Then the experimental space can be determined reasonably. In order to have the NN model in the first iteration be capable of describing the catalytic relationship approximately, the original points must be distributed evenly in the experimental space. For this purpose, the orthogonal design or other method should be used to schedule the original points. At these experimental points the catalysts are prepared and evaluated.

BEGIN)

I Distributing the original experimental points ]

I Performing experiment at each located point I

1

Modeling the catalytic relationships via neural networks with all current data

Forecasting each local optimum via multi-objective optimization

~'

Yes

Taking each local optimum as new experimental points during the next iteration

C_fixdng out the final optimal design

T (END)

Figure 1. The diagram of the iterative procedure in the proposed computer-aided technique for developing catalysts.

1110 2.3. Modeling of the catalytic relationship via neural networks It is very useful to extract the valuable information, both qualitative knowledge and quantitative catalytic model, from all current experimental data. Depending on the information, the experiment in the next iteration can be designed elaborately. In this way the iterative procedure can be led to a hopeful direction. As we known, the catalytic performance of a multicomponent catalyst will be affected by many factors. The correspondent theoretical model describing the catalytic relationship between performance and factors can hardly be built up. Regression method is often used to correlate low nonlinear multivafiable relationships, but for a high nonlinear or strongly synergistic relationship, its limitation arising t~om the determination of the structure of the regression model becomes very serious. However, the neural network model is a useful tool for correlating such a sophisticated relationship between the outputs data (results) and the inputs data (factors) [7]. Neural networks have been successfully used for a number of chemistry applications including correlation of structure-activity or structure-spectrum [8-9], estimation of acid strength of mixed oxides [ 10] and product distribution [ 11]. Although a broad range of NN architectures and learning paradigms are available [12], the back propagation algorithm for multilayer feedforward networks [13] is the most popular approach for engineering and chemistry applications. The multilayer feedforward NN shown in Figure 2 is composed of many interconnected processing units or neurons organized in successive layers. The network is called fully connected because each neuron distributes its output to each neuron in the next layer. The first or input layer is composed of fan-out units vx46ch do not perform any computation but simply distribute their input to all neurons in the next layer. The last layer is the output layer. Between the input and the output layers there can be several hidden layers but only one will be considered here for simplicity. It is assumed that all the hidden and output layer neurons are identical although other choices are possible.

11

02

I1

W~

O/

m 9

J

Is input layer

"X l,N2 )f~

m.. ,ffay r'"

inputs

outputs

output layer

"

Figure 2. Three-layer feedforward neural network.

Figure 3. Neuron i in layerj.

A neuron in the hidden or output layer can be represented as in Figure 3. The neurons perform the following computations. The/th neuron in layerj receives N inputs {I1,'",/N} from layer j-lwith connection strengths or associated weights { W~/,.-., W~ }. The neuron first ~omputes the weighted sum of the N inputs:

1111

S/'= ~ W Jklo, +b J,

(3)

k=l

where b / is a bias term, The bias term need not appear explicitly because it can be interpreted as a weight associated with a constant input of one. The output of the neuron is a nonlinear function of the sum in (3)" O[ = f (S/ ) (4) where f denotes the activation function. The activation function is often chosen to be a sigmoidal functionf(x)= 1/(1+ e -~ ). NN training consists of finding a parameter vector Wthat minimizes the mean square output error:

: ( w-) = T , _1_

'

If ;- f ,ll

(5)

where Wis comprised of all the elements w j M is the number of the data included in the training sample set. Y~ denotes the actual output, and Y~ is the corresponding network prediction. Using a gradient descent method, a solution of W is obtained with J( W ) reaching to its minimum. In this work, the input variables of the input layerX = (x~,x2,...,x,) r are catalytic influential factors, i.e., components of a catalyst and reaction conditions. The output variables of the output layerY=(y~,y2,...,yr,) r are results relating to the catalytic performance, such as catalytic activity and selectivity. Based on the sample set including all the current data, the NN is trained. Here the trained NN is written as: Y = F( X , W)

(6)

where F denotes the NN model function of the catalytic relationship. 2.4. Prediction of the optimal design via multi-objective o p t i m i z a t i o n

With the help of the trained NN model, the effects of factors on the catalytic performance can be simulated numerically. On the other hand, the optimal design for influential factors to gain the best catalytic performance can be also predicted by using the NN model. This will be realized by solving the following optimization problem with multi-objective: Opt{yl,Y2,''',ym} I7 = F ( X , W )

t

s.t.

(7)

X ~R

where R is the feasible region for X . To solve the problem (7), we should transfer it from a multi-objective problem to a single-objective one. A case of the transformation is as below:

M~

1 ~ J= m-E. (Y,- Y7)2 ;=

? = F (2 ,O7 ) s.t.

XeR

(8)

1112 where Y7 is the ideal value of the ith sub-objective. A number of conventional optimization techniques, complex method for instance, should be used with respect to solving the single objective problem The result is an effective (or Pareto) solution of the multi-objective problem and just the optimal design for the experimental procedure.

3. EXPERIMENTAL Depending on the sequential method mentioned above, in our work of developing catalysts for propane ammoxidation to acrylonitrile, we also focused efforts on the catalytic active system of V-Sb-AI mixed oxides. The work was divided into three parts, i.e., optimizations of the composition of promoters and supporters, the composition of main components as well as reaction conditions.

3.1. Catalyst preparation In a stirred flask equipped for heating under reflux, the solid of N]-I4VO3 was dissolved in hot distilled water, to the hot solution the solid of Sb203 was added The slun~ was maintained under reflux for 18 hours. Then the solution was added to a prepared soft and homogeneous A1203-SiO2 gel, and followed by the addition of (]~]A)5I-Is(H2(WO4)6), SnO2, (]N[I-I4)H2PO4, CrO3, (NH4)6Mo7024"4H20, etc.. With stirring the resulting slun~ was evaporated. The thick material was dried in an oven at 114~ overnight, then ground and dried in air at 350~ for 5 hours. After cooled, the dried material was pelletized then screened and the 30 to 60 mesh particle size was collected, the screened material was finally heat treated at 610~ in air for 3 hours.

3.2. Catalyst evaluation The catalyst evaluation mentioned in section 4.1 and 4.2 was carried out in a tubular stainless steel fixed bed reactor. The gaseous feed components were metered through mass flow controllers into the top of the reactor at atmospheric pressure. Steam was introduced by air through a water bath. Air was taken as the source of oxygen. For the evaluation of catalysts, the reaction temperature was 500~ the molar feed ratios was C3HdNH3/Oz/H20 = 1.0/2.2/2.1/3.0, and the reaction contacting time is 14.0 h'gC3Hdgcat The catalyst was activated by introducing air at 610~ for two hours and NH3 at 5300C for 30 minutes. The analyses of the products was performed with on-line GC. The catalyst used in the expeiiment of optimizing reaction conditions is the best one obtained in the work of optimizing the main components. In order to observe the effect of each condition on the catalytic performance, the ranges of conditions were broadened deliberately.

4. RESULTS AND DISCUSSION

4.1. Optimization of the composition of promoters We selected the contents of promoters, such as P, K, Cr, Mo, the content of V-Sb and the weight-ratio of A1203/SIO2 in the supporter as six influential factors. Originally, 19 sample points were distributed, at which the catalysts were prepared and evaluated, thus all of the data were collected to compose the original training sample set. Among the data, the ranges of the catalytic performance were, the conversion of propane Xp 5 3 . 3 % - 82.1%, the selectivity of acrylonitrile SACS2 3 . 5 % - 50.5% and the acrylonitrile yield YAcs 14.3% ~ 34.9%.

1113 Based on the sample set, a NN configured in 6-20-12-2, that is six units (six influential factors) in the input layer, twenty units in the first hidden layer, twelve units in the second hidden layer and two units (the conversion of propane Xp and selectivity of acrylonitrile SACN)in the output, layer, was trained. The typical error curves of NN self-learning is shown in Figure 4. The fitting ability of the trained NN was tested with some data outside the sample set. We can see from Figure 4, up to 130000 times of the NN learning, the overfitting (or overlearning) problem is forthcoming. At this point the NN learning process should be finished. 0.09 9

~

T

:3 TestingPatteras I 0.07 0.06 . . . .

~li_

0

l~r~tlllUlg ratL~lns I 0.08

_d -~--'--~

,

~

1 0"04 o.o ~.~.~,j, 0.02 O.Ol o 50000 100000 150000 200000 LearningTimes

Figure 4. Root-mean-square error curves for NN training.

Yacn% 60 50 40 30 20 10 0 ori.

[] opt. m Exp.

1st

2nd

3rd

Figure 5. The best ACN yields in each iteration for optimi~ng promoters.

Then the sequential steps mentioned in section 2 were repeated till the precision was satisfied, and the catalytic results corresponding to the optimal composition were XP 80.0%, SAC~53.7% and YAC~43.0%. For this case the iteration was undertaken thrice, in each one four local optimal composition were predicted and then tested experimentally. The best acrylonitrile yields (including both results of calculation and experiments) in each iteration are shown in Figure 5.

4.2. Optimization of the composition of main components Based on the best composition of promoters, the composition of main components was optimized. The contents of V, Sb, W, Sn and supporter were taken as five influences, and XP, SAC~ also as the outcomes. 20 Original sample points were distributed, at which the catalysts were prepared and evaluated. The ranges of the results were, Xv 11.6%~88.2%, SAC~ 15.0% 44.4% and YAC~ 2.4% - 28.9%. Accordingly, A 5-20-12-2 network was adopted in the sequential iterative procedure. The results of the optimization and evaluation in the first and second iteration are listed in Table 1 and 2. Table 1 Results of the optimization and evaluation in the first iteration Optimization (%) Catalysts Evaluation (%) No. XP SACN XP SACN 96.987 58.041 98.216 50.984 C-11 94.424 57.046 96.039 50.651 C-12 82.344 65.999 76.096 58.131 C-13 C-14 86.560 57.960 80.986 45.598 ,..

Relative errors (%) XP

SACN

1.254 1.682 -8.211 6.883

-13.840 -12.626 -18.595 -21.329

1114 From Table 1, we can find that the highest YACNof evaluations is 50.1%, better than the best result in the work of optimi~g promoters. But the relative errors of the acrylonitrile selectivity between the optimization and evaluation are big. Therefore, the data of these four points were added to the original ~rnple set for training the NN model. Thus the trained NN model became more accurate, so did the optimization. This fact can be seen from Table 2. Table 2 Results of the Catalysts No. C-21 C-22 C-23 C-24

optimization and evaluation in the second iteration Optimization (%) Evaluation (%) Relative errors (%) XP XP SACN XP SACN SACN 95.085 67.635 93.696 59.762 -1.461 -11.64 81.246 64.126 80.348 66.722 -1.105 +4.048 94.206 66.431 83.760 67.647 -11.09 +1.830 86.016 64.889 84.994 62.882 -1.188 -3.093

The results in the second iteration showed that the relative errors between optimization and evaluation were less than those in the first iteration, and YACNreached to 55.0%. After performing the third iteration, the best catalyst of the optimization was similar to the best one tested in the 2nd iteration. It meant that the catalyst better than the best one obtained did not exist in the experimental space. So the iterative procedure Was finished, and the results relating to the best composition of main components were Xp 93.7%, SAcN 59.7% and YAcN 56.0%. The best YACNin each iteration is shown in Figure 6.

4.3. Optimization of reaction conditions It is knportant to consider the effects of reaction conditions, such as reaction temperature, ratios of ammonia to propane(N/C), oxygen to propane (O/C) and steam to propane (H/C), on the catalytic performance. To investigate and optimize these effects, we used a 4-12-7-2 network to correlate the relationship between the catalytic performance and conditions. Yacn% 70 60 50 40 30 20 lO l, 0 ori.

/

~Opt. mExp.

Yacn% 60

iji!i - i i lil

40

[!ii!' i"i-iiili"i] iiii"i 1st

2nd

3rd

Figure 6. The best YACN in each iteration for optimi~ng main components.

OOpt. IExp.

30 10 0 off.

1st

2nd

3rd

Figure 7. The best YACN in each iteration for optimi~g reaction conditions.

The original sample set consisted of the data of 16 experimental points. As above, the iterative procedure was undertaken again to search for the best reaction condition. Six experimental points were tested in three following iterations. Only twice was the procedure

1115 repeated, the iterative precision was satisfied. The final best results were: Xp 85.2%, SACN 69.2% and YAcs 58.9%, which corresponded to the conditions of reaction temperature 524.8~ and C3Hs/NH3/Oz/I-I20 = 1.0/3.0/4.4/6.0. At these molar feed ratios, a lower concentration of propane in the feed gas led to a lower space time yield. However, the results exhibited the synergistic effects of reaction conditions on the catalytic performance. By adding some desired constraints on the influential variables in problem (7), the iterative procedure will make out more reasonable reaction conditions for the commercial meaning. Furthermore, the best ACN yields in each iteration are shown in Figure 7. 4.4. Discussion

In three parts of work to optimize the composition of promoters and main components, as well as the reaction conditions, totally, 81 experimental points were tested. The best ACN yields of each part are shown in Figure 8. The best YAcs was raised sharply from initial 34.9% to final 58.9%. This fact demonstrates the proposed technique is very effective for developing catalysts.

60

Yacn% m

50 40 30 20 10 0

m 1 iiiiiiiiI

m m iiiiiiiil !iiiiiii!i

I iiiiiiii/iiiiiiil/iiiiiiii , m Init.

Prom.

Main

Cond.

Figure 8. The best ACN yields in each iterative procedure for optimizing promoters, main components and conditions. Otherwise, there still existed some problems to be considered in our work. First, the method of distributing the original experimental points should be capable of making the distribution reflect the tendency of catalysis sufficiently with the least points. The more the points is distributed, generally the more evident the tendency is, but the larger the amount of experimental work is. On the other hand, a small data sample set is not capable of disclosing the catalytic tendency adequately, then result in an overfitting problem in NN learning. Therefore it is an interesting question for the technique how to avoid the dilemma. As a successive research of this work, we are having a try by using uniform design to take the place of the orthogonal design. Secondly, for the development of multicomponent catalysts, the influential factors are too many to be included wholly in an iterative procedure. As we done in this work, dividing a whole catalytic system into three parts, some synergistic effects of the factors might disappear. Possibly this difficulty may be overcome by combining the theoretical or heuristic knowledge and the modeling of NN. Under the guidance of heuristic analyses, the experimental space will be determined strategically.

1116 5. CONCLUSION To develop a multicomponent catalyst, the information, the quantitative relationship between the catalytic performance and influences particularly, extracted from the data is very useful. It is necessary and crucial to work out an interactive strategy for controlling the whole process from data acquiring, correlating the data to forecasting the optimal design. For this sake, in the paper, a new computer-aided technique was proposed and then used to develop catalysts for propane ammoxidation to acrylonitrile. Under a certain reaction condition, an optimal catalyst with high-performance of Xp 85.2%, SACN 69.2% and YACN58.9% was obtained. This example demonstrated that the technique can raise the efficiency of developing catalysts obviously.

REFERENCES

1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13.

H.C. Foley, E. E. Lowenthal and X.-D. Hu, Computer Aided Innovation of New Materials II, Elsevier Science Publishers B.V., (1993) 1101. R.A. Van Santen, Chem. Eng. Sci., 45 (1990) 2001. E.R. Becker and C. J. Pereira (editors), Computer-aided Design of Catalysts, Marcel Dekker, Inc. 1993. G. Centi, R. K. Grasselli and F. Tfifiro, Catal. Today, 13 (1992) 661. R. Nilsson, T. Lindblad and A. Anderson, J. Catal., 148 (1994) 501. A.T. Guttmann, R. I~ Grasselli and J. F. Brazdil, US Patent, No. 4 746 641, (1988). X.-Q. Wu, B.-Y. Li and D.-W. Lu, Proceedings of the 7th National Conference on Chemical Engineering, Beijing, China, (1994) 1136. J. Zupan and J. Gasteiger, Anal. Chim Acta., 248 (1991) 1. P.C. Jurs, CICSJBulletin, 11(5)(1991) 2. S. Kito, T. Hattori and Y. Murakami, Ind. Eng. Chem. Res., 31 (1992) 979. S. Kito, T. Hattori and Y. Murakami, Appl. Catal. A: General, 114 (1994) L173. B. Widrow and M. A. Lehr, Proc. IEEE, 78 (1990) 1415. D. E. Rumelhart and J. L. McCleUand, Parallel Distributed Processing: Explorations in the Micro structure of Cognition, MIT Press, Cambridge, MA, 1986.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1117

Catalysts by Rational Design: Prediction and Confirmation of the Properties of the Co/Ce/Br liquid-phase autoxidation Catalyst Based on the Kinetic Similarity to the Co/Mn/Br Catalyst Rob K.Gipe a and Walt Partenheimerb,c a7081 93rd Ave. SE, Mercer Island, WA 98040, USA b Central Research and Development, DuPont Experimental Station, Wilmington DE 19803-0262, USA CWork performed at the Amoco Research Center, Naperville, IL 1. ABSTRACT The kinetics of m-chloroperbenzoic acid (MCPBA) with Co(II)/Mn(II)/Br(-I) and Co(II)/Ce(III)/Br(-I) mixtures in acetic acid is reported. MCPBA first reacts with Co(II) to give Co(Ill), then Co(III) oxidizes Ce(III) to Ce(IV), which is followed by Ce(IV) oxidizing Br(-I) to dibromine, see figure 1. This is similar to the observed reactions of MCPBA with the Co(II)/Mn(II)/Br(-I) combination. We therefore predicted that the Co/Ce/Br combination would be a similar autoxidation catalyst to Co/Mn/B ,i.e.,there is a synergistic interaction when either Mn or Ce is added to a Co/Br catalyst as well as a lowering of the rate of carbon dioxide formation. This was subsequently observed. The reaction of toluene with dioxygen, catalyzed by Co/Br, Mn/Br, Ce/Br, Co/Mn/Br, and Co/Ce/Br mixtures is reported. Problems encountered in determining a synergistic interaction are delineated and include 1) the co-oxidation of a number of intermediates, 2) having different compositions of intermediates at different times, and 3) that fact that a substantial amount of the initially added bromide exists as benzylic bromide. Hammett studies of the Co/Mn/Br and Co/Ce/Br catalysts give similar rho values of-1.2 and -1.4, respectively, indicative that both catalysts operate via a similar mechanism. 2. INTRODUCTION The discovery of metal/bromide catalysts in 1954 [1] evolved into one of the major industrial processes producing oxygenates from hydrocarbons using dioxygen as the primary oxidant: metal/bromide hydrocarbon + 0 2 ;, oxygenate + H20 ( 1) HOAc The primary example is the production of terephthalic acid from p-xylene which is used in the production of polyethylene(terephthalate). A recent review describes the oxygenation of at least 251 hydrogens to give 279 products using at least 35 different catalyst combinations. There are a number of reports which include cerium as one of the elements

1118 of the catalyst package, see table 1. Addition of cerium to a Co/Br catalyst has been reported to increase the rate of conversion of 2-methylnaphthalene [3]. This information led one of the authors to extend our kinetic studies to include interactions of cerium with cobalt and bromide. This kinetic study led to the prediction that 1) cerium would be synergistic with Co/Br in the same way that manganese is, and 2) reduction in the formation the wasteful by-products, carbon dioxide and monoxide, would occur.

MCPB

~2"

Co(HI

Mn(HI

Br CH3 CHa

MCBA

Co(ID~

MCPBA-~

Co(

- Mn(ll) ~

~

Br"

O CHs ~

Ce(IV)

2

Br

CHI

MCBA ~-~

Co(H)~,/

~

Ce(III)

Br

~3

CHa

Figure 1. Most favorable kinetic pathways when MCPBA is added to Co/Mn/Br and Co/Ce/Br mixtures in 10% water/acetic acid 3. EXPERIMENTAL

3.1 Kinetic studies These were determined as previously reported [4]. 3.2 Autoxidation of toluene Measurements were made in a glass cylindrical reactor as previously described [5]. Initial conditions were 1.09M toluene in 100.0 ml anhydrous acetic acid at 98-100~ The concentrations of metals in the Co/Br, Mn/Br, Ce/Br catalysts is 0.0200M and 0.0100M for each metal in the Co/Mn/Br and Co/Ce/Br catalysts. The bromide/metals ratio is always 1.00 mol/mol. Initial salts used were cobalt(II) and manganese(H) acetate tetrahydrates, cerium(m) acetate hydrate, and sodium bromide. Rates of oxygen uptake and carbon monoxide and carbon dioxide formation were calculated from the GC data knowing the flow rate of gases through the reactor and the composition of air.

1119 Table 1 Reported Metal/Bromide Catalyzed Autoxidations containing Cerium Catalyst a Co/Ce/Br 68 b Co/Ce/Br 76 c Co/Ce/Br d Co/Ce/Br

reagent/product tetralin/acetoxytetralin 2,6-dimethylnaphthalene/2,6-dicarboxynaphthalene p-cresyl acetate/4-acetoxybenzoic acid

solvent/yld Ac 20 HOAc Ac20

90

decahydronaphthalene/mixture of acetates and ketones Ac20 2,6-diisopropylnaphthalene/2,6-dicarboxynaphthalene o-toluic acid/o-phthalic acid 4,4-dimethylbiphenyl/4,4'-dicarboxybiphenyl 2-methylnaphthalene/2-naphthoic acid 2-methylnaphthalene/2-naphthoic acid benzylacetate/benzaldiacetate HOAc 45 toluene/benzylacetate Ac20 15

e f g h i j k

Co/Ce/Br Co/Ce/Br Co/Ce/Br Co/Ce/Br Co/Ce/Br Co/Ce/Zn/Br Co/Ce/Zn/Br

1 m n o p q

p-nitrotoluene/p-nitrobenzoic acid Co/Ce/Zr/Br 2,6-diisopropylnaphthalene/2,6-dicarboxynaphthalene Co/Mn/Ce/Br 1,1 -bis(3-methylphenyl )ethane/1,1 -bis(3,3'-benzoic)ethane Co/Mn/Ce/Br o-toluic acid/o-phthalic acid Co/Mn/Ce/Br Co/MoJI'i/Ce/Br 1,2,4-trimethylbenzene/ 1,2,4-tricarboxybenzene 2,6diisopropylnaphthalene/2,6-dicarboxynaphthalene Co/Mn/Ce/Br

a. Jap. Pat. 56118042 (1981) b.T. Maki and Y. Asahi, Jap Pat 61210052 c. K.Yazu,M.Saito, K.Ukegawa,T. Nakayama, S. Tadao, T. Suzuki, M. Ono, J. Miki,N.Takei,K.Tate,Nippon Kagaku Kaishi, 1,(1991)92-96. d. K.Yazu,T.Wakabayashi,T. Nakayama, Chem Lett, (1986) 1409-1412. e. S.Hayashi,T.Matsuda,A.Sasawaka, Ger Pat DE 3707876 A1, (1987). f. T. Nakayama, E. Nakamura, K. Koguchi,Nippon Kagaku Kaishi, 4, (1982), 650 g.H. Mami et al, Jap Pat 310846, (1988). h. Y. Ogawa, T. Yamada, US 93,339,202(1993) i. Y. Ogawa, T. Yamada, Aromatikkusu,46(3/4),80-3 (1994)(Japan)(CA121(3):34977f) j. Jap. Pat. 56104845 (1981) k. Jap. Pat. 56104846 (1981) 1. Jap. Pat. 59080637 (1984) m. P.A.Sanchez,D.A.Young,G.E.Kuhimann.W. Partenheimer W P Schammel, US Pat 4,950,786, (1990). n. H. Mami et al, Jap Pat 310846, (1988). K.Yazu,T.Nakayama, Nippon Kagaku Kaishi,3 (1988)304-310.(CA 109(21): 189633s) o. Jap. Pat. 47042639 (1972) to K. Nakaoka and S. Wakamatsu p. C. Fumagalli, L. Capitanio, G. Stefani,EP513835, (1992). (CA118(8):60707a) q.T. Yamada, K. Sugiura, Y. Doko, K. Maeda, R. Minami,Y. Nagao, JP4330039

1120 3.3 Hammett studies These were performed in the same reactor using identical procedures as those described above. Two or three reagents were simultaneously oxidized and the first order rate constants of their disappearance were calculated. Each experiment was repeated twice and the averaged rate constants were used in figure 2. M-xylene, at a concentration of 1.00 M, was used as the kinetic standard in each experiment, i.e., the rates of all the other reagents were compared to it. The concentration of all the other reagents were 0.150M. ..................................... ~-.~5

~.. D-OCH3

4-

p-t-t)utyf~t~

m

E -0.5

1

0.5

~m CH3 (

0 ~5

m-Cl "%

n.-C'.t~9~H.':I

<. p-N02

9

Iog(k/ko) Figure 2. Hammett plot for a Co/Ce/Br catalyst in anhydrous Acetic acid at 80~ 4. RESULTS AND DISCUSSION

4.1 The Kinetic Comparison of Reactions of MCPBA with Combinations of Co(ll)/Mn(ll)/Br(-l) and Co(ll)/Ce(lll)/Br(-l) The rate constants are summarized on table 2. We will omit the ligands and the polynuclearity in the following discussion because, except for Co(II) acetate, they are unknown. In work initially reported by Jones [6,7] and confirmed by ourselves [4], MPBA very rapidly reacts with Co(II) acetate in acetic acid/water solutions to form Co(III)a [8]: 30 C Co(II) 2 + MCPBA :> 2 Co(III)a + MCBA (2) HOAc/H20

t 1/2 =0.016 sec.

1121 In contrast to Co, Ce(III) acetate reacts quite slowly with MCPBA. The reaction of 0.0029M MCPBA with 0.0250M cerium(m) acetate results in a colorless solution slowly turning yellow at room temperature. The yellow color is due to an increase in absorbance in the 350-400 nm region due to the formation of cerium(IV): 23 ~ 2 Ce(III) + MCPBA

;, 2 Ce(IV) + MCBA

(3)

HOAc/H20 t 1/2 = 190min When a solution of Ce(III) (0.005M) and Co(II) (0.005M) is added to a solution of 0.0580 M MCPBA, the UV-VIS indicates the presence of only Ce(IV) and Co(II) after 4 sec (the time of mixing). Since MCPBA does not react that quickly with Ce(III), rxn 3, but does react quickly with Co(II), rxn 2 , we surmised that the MCPBA first reacted with the Co(II) to form Co(III)a and then Co(III)a reacted with Ce(III). This was confirmed using the stopped-flow apparatus. At 600nm, there was an initial increase in absorbance due to formation of Co(III)a which was followed by a decrease due to the reaction of Co(III)a with Ce(III): 23 ~ Co(III)a + Ce(III)

:~, Co(II) + Ce(IV)

(4)

HOAc/H20 t 1/2 = 0.042 sec The chemistry of cerium(Ill) acetate is similar to that of Mn(II) acetate in that both do not react as fast with MCPBA as Co does, but do react quickly with Co(III)a.

Ce(IV) (0.00125M) was prepared by mixing MCPBA, Co(II), and reacting the mixture with 0.0125M KBr:

Ce(III), and then

23 ~ Ce(IV) + Br-

- .......

- > Ce(III) + 1/2 Br 2

(5)

HOAc/H20 t 1/2 = 38 min. The rate of reaction 5 is only 0.73 that observed for Mn(III) + Br- when corrected for the slightly different concentrations of the metals and bromide and assuming a first order dependence of Ce(III) and Br-, see table 2. Co/Ce/Br are illustrated in figure 1.

The kinetic similarity of Co/Mn/Br and

There are at least three different forms of cobalt(III) in acetic acid which have been dubbed Co(III)a, Co(III)s and Co(III)c by Jones [6]. Co(III)a reacts further to form

1122 Co(HI)i, a change easily observed in the visible region, and this has a half-life of 1.7 min at 30 ~ C. Co(III)a, 'cobalt(Ill) active', is a more reactive form than Co(HI)i, 'cobalt(llI) stable'. Table 2 illustrates that Co(III)i reacts 2-3 orders of magnitude more slowly than Co(III)a toward Mn(II) or Ce(III) (when corrected for the differences in concentration). The spectrum of Ce(IV) does not change thereafter for 20 minutes suggestive that there are not different forms of Ce(IV) as there are for Co(l/I). Table 2 Rate Constants for Selected Reactions for MCPBA, Co(g), Co(III), Ce(III),Ce(IV), Mn(II),Mn(III), and Bromide in 10% Water/Acetic acid a 9

Reactants

~

Products

k,s- 1b

temp.C comments

MCPBA + Co(II

MCBA + Co(l/I0 66(2)

30

MCPBA + Ce(III) MCPBA + Mn(II) MCPBA + KBr

MCBA + Ce(IV) 6.2(.2)x10 -5 MCBA + Mn(III) 0.017 MCBA + KBr 3 0.08

23 23 23

(c) (d)

MCPBA Co(III)a Co(III)s + Ce(III) Co(III)s + Mn(II) Co(III)a + Ce(III) Co(III)a + Mn(II) Ce(IV) + KBr

MCBA 2x10 -6 Co(III)s 0.0070(.0002) Co(II) + Ce(IV) 0.00067 Co(II) + Mn(III) 0.0015 Co(II) + Ce(IV) 16.3(1.6) Co(II) + Mn(III) 6.6(0.1) Ce(III) + KBr 3 0.00031 (0.1)

25 25 23 23 23 23 23

(e)

Mn(III) + NaBr

Mn(II) + NaBr 3 0.0066(.0004)

23

(f)(g) (f)

(h)

a. [MCPBA]o=0.0005 M and all others 0.0100M unless otherwise stated. All data have been measured in our labs except the stopped flow measurements which were performed at Purdue University together with Bill Schleper in Dale Margarum's lab. b. Standard deviation, in parenthesis(), based on at least three independent measurements. c. Initial concentrations are Ce(I11)=0.0025 M, MCPBA=0.0029M d. Autocatalytic reaction, rate refers to fast part of S curve e. Rate of thermal decomposition reported by C.F. Hendriks, H.C.A. van Beek, and R.M. Heertjes, Ind. Eng. Chem. Prod. Res. Dev., 18 (1979) 38 for perbenzoic acid. A number of other peracids give approximately the same rates. f. Initial concentrations are Co(III)s=0.00025, Ce(III) or Mn(II)=0.0025M g. Using a sample of Co(Ill) prepared via ozone. h. Initial concentrations are Ce(IV)=0.00125M, KBr=0.0125M.

4.2 The Synergistic Interactions in Co/Ce/Br and its Similarity to Co/Mn/Br The synergistic interaction for the Co/Mn/Br catalyst has been previously reported by Ravens [9] based on the replacement of some of the cobalt by manganese in a Co/Br catalyst and by us [ 10] based on the fact that the sum of the activities of Mn/Br and Co/Br catalysts is less than the Co/Mn/Br catalyst.

1123 We will define the synergy factor (SF) as the rate of reaction (R) in the presence the catalyst with components X1/X2/X3... divided by sum of the rates of reaction of the individual components X1,X2,X3... (or sums of these components). For Co/Mn/Br and Co/Ce/Br catalysts, respectively, we have: SF = RCo/Mn/Br/(RCo+RMn+RBr+RCo/Br+RMn/B r)

(6)

SF = RCo/Ce/Br/(RCo+RCe+RBr+RCo/Br+RCe/B r)

(7)

A SF value > 1.00 indicates a synergistic interaction and a value < 1.00 indicates an antagonistic interaction while a value of 1.00 would indicate the absence of synergy. Cobalt(H) acetate itself is an autoxidation catalyst [11 ], as is manganese(H) acetate [ 11 ], and as are bromide compounds [12]. Under the mild conditions employed in these experiments (approximately 1 atmosphere of air and 100~ Co catalyzed oxidation of toluene has an induction period of about 1 hr, the rate of oxygen uptake is roughly 0.3 ml O2/min and hence this will be subsequently ignored (RCo =0.0). We do not observe oxygen uptake with manganese(H) acetate, cerium(Ill) acetate or sodium bromide catalysts with the conditions employed in these experiments, hence their rates are zero (RMn=RBr=Rce=0.0). Values of SR > 1.0 are always observed for both the Co/Mn/Br and Co/Ce/Br catalysts. We also find a higher degree of synergy in the Co/Mn/Br catalyst than the Co/Ce/Br one, i.e., SFCo/Mn/Br > SFCo/Ce/Br. For example from the rates of oxygen uptake we have: Table 3 Results from the Autoxidation of Toluene using Co/Br,Mn/Br, Ce/Br, Co/Mn/Br and Co/Ce/Br Catalysts rate oxygen uptake,ml O2/min time, hr Co/Br Mn/Br Ce/Br Co/Mn/Br Co/Ce/Br SFCo/Mn/Br SFCo/Ce/Br 1.00 2.00 3.00 4.00 5.00

1.52 1.21 1.02 1.12 1.01

1.20 1.12 0.78 0.60 0.51

0.69 0.66 0.65 0.35 0.32

7.20 7.18 7.49 3.86 3.02

4.26 3.43 3.04 3.33 3.63

2.64 3.08 4.16 2.24 1.99

1.93 1.83 1.82 2.26 2.73

The SF can also be calculated in other ways. For example, the yield of benzoic acid for Co/Br, Mn/Br, Ce/Br, Co/Mn/Br, and Co/Ce/Br at 5.0 hr is 4.6, 1.4,0.0, 40.4 and 20.0 mol %. This gives a SFCo/Mn/Br =6.7 and SFCo/Ce/Br = 4.3. Figure 3 gives the activity of the catalysts, based on the conversion of toluene, where Co/Mn/Br > Co/Ce/Br > Co/Br > Mn/Br > Ce/Br. Synergy factors can be calculated at a given value of conversion. From figure 3 at 3.00 hr one obtains SFCo/Mn/Br =1.38 and SFCo/Ce/Br = 1.21. Using

1124 the data on figure 3, one obtains the first order rate constant for toluene disappearance as 1.57, 1.43, 0.86, 5.60 and 3.22 (xl03,s -1) for Co/Br, Mn/Br, Ce/Br,Co/Mn/Br, and Co/Ce/Br catalysts, respectively. From these rate constants one obtains SFCo/Mn/Br = 1.86 and SFCo/Ce/Br = 1.32. The latter are probably the most meaningful synergy factors. The SF values on table 3 should be measured at conditions where the concentration of the oxidant and catalyst are identical. This is very difficult to do for the following three reasons: 20 90

8O 7O

~r

6o

0 so W ,-

-r

Co/Br

-a-

Mn/Br

j.~r"

Ce/Br

/~

=

40

'0 9 ,.m

>m

~ "0

--x- Co/Mn/Br

/X------- N

18

M

\

14 12

,,,o .c

C o / C e / B r .-- x j

8 20

.

Co/Br

Ibe~

[]

"

A

Mn/Br

m 2

10

0

0 0

2

4

time,hr

Figure 3. Activity of Selected Catalysts

0

20

40

\

C /Br

•

Co/Mn/Br

=

C o )/ C e / B r 60

80

Conversion,%

Figure 4. Benzaldehyde Formation for Selected Catalysts 1. One is not observing only the autoxidation of toluene in these experiments but rather the co-oxidation of toluene with benzaldehyde, benzyl alcohol, benzyl acetate, and the benzyl bromide. Benzaldehyde, benzyl alcohol, and benzyl acetate are all more reactive than toluene [13]. 2. We do not know if the distribution of these intermediates will remain constant for a given conversion of toluene with a given catalyst. Figure 4 indicates otherwise. Unfortunately, the concentrations of the benzyl alcohol and acetate were not measured. The benzaldehyde yield, at a given value of conversion, varies considerably from catalyst to catalyst. 3. The amount of catalytically active bromide is different for each due to the variation of the concentration of (z-bromotoluene during the experiments. (z-Bromotoluene is an inactive form of bromine in these reactions [ 14]. The appropriate manner to express the yield of benzylic bromides is on the initial sodium bromide added rather than on the initial toluene basis since sodium bromide is the limiting reagent. The benzylic bromide yields, on a sodium bromide basis, range from 22% for the Mn/Br catalyst to 93 for the Co/Mn/Br catalyst, see figure 5. Thus the effective bromide concentration varies from 7 to 78%

1125 Interestingly, the more active catalysts contain the highest amount of inactive bromide, i.e., the highest yield of o~-bromotoluene, see figure 5 ! 4.3. The Steady State Oxidation States of the Metals and Reduction in the Formation of Carbon Oxides (CO, CO2).

Observation of the colors during the metal/bromide oxygenation of toluene are consistent with very low steady state concentrations of the metals in their higher oxidation states. Previously we had estimated based on UV-VIS studies that only 0.6% of the cobalt in a Co/Mn/Br catalyzed oxidation of p-xylene consisted of Co(III) [10]. In the absence of 100 90 80

_~ r

9

70 60

4O

r

20

X

9 Co/Br

Xlx'X'

9 "0

A& A= A

5O O I-J IZl

x

9 9 9

.

9Mn/Br 9

=A

9Ce/Br

m

a

9

9 Co/Br

[]

&e

~2

9M n / B r

3O

x Co/Mn/Br

9Ce/Br

' v

9

x

Co/Ce/Br

X

X

x Co/Mn/Br

10

x 0

20

40

Co/Ce/Br

60

80

Converslon,%

Figure 5. Formation of o~-Br-Toluene For Selected Catalysts

o 100

~" 0

20

40

60

80

converslon,%

Figure 6. Formation of Carbon Dioxide for Selected Catalysts

bromide, the UV-VIS are consistent with substantial amounts of Co(III) and Mn(III). The observed colors of the oxygenations of toluene, during the first hour, are: Catalyst Observed During Rxn Colors of Metal Acetates in Acetic acid Co green Co(II), pink Mn brown Co(III), green Co/Br pink Ce(III), light yellow Mn/Br colorless Ce(IV), colorless Ce/Br colorless Mn(II), colorless Co/Ce/Br pink Mn(III), brown Co/Mn/Br pink The replacement of some of the cobalt by manganese reduces the steady state concentration of Co(III) even lower during the catalytic reaction which results in a lower rate of carbon dioxide formation [10]. This is also observed when some of the cerium is replaced by cobalt i.e. in a Co/Ce/Br catalyst, see figure 6.

1126 4.4 Hammett Studies of Co/Mn/Br and Co/Ce/Br catalysts Hammett studies of metal/bromide catalysts have been previously reported by Kamiya [15] and ourselves [10]. The results for a Co/Ce/Br catalyst are given on figure 2. Six different reagents are reported using a range of sigma+ substituent constants from 0.79 (p-nitrotoluene) to -0.78 (p-methoxytoluene). The rho value of the Hammett equation (the slope of figure 2) is -1.4(0.12) which is similar to that of-1.2(0.11) [17] determined for a Co/Mn/Br catalyst. The negative value of the rho values indicate a build-up of positive charge in the aromatic ring in the transition state [ 16]. In cobalt only oxidations, which utilize Co(III) as the oxidant, the rho value is -1.8, consistent with a radical cation mechanism [ 17]. The lesser values of the metal/bromide rho values reported here, and by Kamiya [ 15] indicate a different transition state than a radical cation. The value of rho for the benzylic abstraction of hydrogen atoms by the bromide atom of-1.82 (in acetic acid) [18] is also not consistent with the values of-1.2 to -1.4 reported here as is benzylic abstraction by benzyl peroxy radicals in toluene which give a value of-0.63 [19]. The transition state may therefore involve bromine atom abstraction but with the bromide partially or wholly bonded to the manganese(II) metal as originally suggested by Kamiya [ 15] and supported by the authors [2,4,10]. REFERENCES

1. R. Landau and A. Saffer, Chem. Eng. Prog., 48,(1968)20. 2. W. Partenheimer, Catalysis Today, 231 (1995). 3. Y. Agawa, T. Yamada, Aromatikkusu 46(1994)80 (Japan), CA121(3):34977fS. 4. T. Oyama, and J. W. Hightower, "Catalytic Selective Oxidation", Amer. Chem. Soc., 1993, chapter 7 by W. Partenheimer and R.K. Gipe. 5. W. Partenheimer, J. Mol. Catal., 67(1991)35. 6. Jones, G.H., J. Chem. Research (S) (1981) 228-229. 7. Jones G.H.,J. Chem. Soc., Chem. Commun.,(1979)536;Jones, G . H . , J. Chem. Research (M), (1981)2801 ;Jones, G.H.J. Chem. Research (S) (1982)207. 8. There are at least three different forms of cobalt(Ill) in acetic acid. Following Jones notation, we have labeled these Co(III)a, Co(III)i, and Co(III)c. 9. D.A.S. Ravens, J. Chem. Soc.,55(1959) 1768. 10. D.W. Blackburn, "Catalysis of Organic Reactions", Marcel Dekker,Inc., 1994, chapter 14, by Walt Partenheimer. 11. S. Carra, E. Santacesaria, Catal. Rev.-Sci.Eng., 22(1980)75. 12. J.E. Mclntyre and D.A.S. Ravens, J. Chem. Soc., (1961) 4082. 13. R.A. Sheldon and J.K. Kochi, "Metal-Catalyzed Oxidations of Organic Compounds", Academic Press, New York, N.Y. 1981, p 23. 14. K. Sakota, Y. Kamiya, and N. Ohta, Bull. Chem. Soc. Japan, 41 (1968)641. A.S. Hay, and H.S. Blanchard, Can J. Chem.,41(1965)1306. See also reference 2. 15. Y. Kamiya, J. Catalysis,44(1974)480. 16. T.H. Lowry, K.S.Richardson, "Mechanism and Theory in Organic Chemistry, Harper and Row, 3rd Edition, page 143. 17.C.F. Hendriks, H.C.A. Beck, P.M. Heerjes, Ind. Eng. Chem. Prod. Res. Dev.,17(1978)266. The latter report a value of-1.9 but did not account for the number of equivalent hydrogen atoms. We have taken the values from this reference

1127 and plotted them vs. sigma + values to obtain value of- 1.8. 18. J.R. Gilmore, J.M. Mellor, Chem. Comm., (1970),507. 19. G.A. Russell, J. Amer. Chem. Soc.,78(1956), 1047.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1129

T h e K i n e t i c s of the Partial O x i d a t i o n of M e t h a n e to F o r m a l d e h y d e over a S i l i c a - S u p p o r t e d V a n a d i a Catalyst A.W. Sexton and B.K. Hodnett Department of Chemical and Environmental Sciences, University of Limerick, IRELAND. ABSTRACT A kinetic study has been carded out on the partial oxidation of methane to formaldehyde over a silica-supported vanadia Catalyst. The results indicate that oxygen was adsorbed on the catalyst and took part in the reaction in an Eley-Rideal or Mars-van Krevelen manner. The nature of the interaction with the catalyst was dependent on whether the reaction took place in methane rich (PcH4 - 80 kPa) or lean (PcH4 = 4 kPa) conditions. A reaction mechanism for the partial oxidation of methane to formaldehyde is proposed, which is consistent with the data reported here. Methanol oxidation experiments over this catalyst suggested that it was not an intermediate under the conditions employed during this study. INTRODUCTION In recent years there has been much interest in the conversion of methane to value added products, such as ethane/ethylene [1], methanol [2], formaldehyde [3-5] and synthesis gas [6]. Many studies have been carded out on the partial oxidation of methane to formaldehyde over silica [7], and over molybdena [8,9], vanadia [5,10] supported on silica, or FeNbB-O [11]; with nitrous oxide [7-9] or oxygen [7,10] as the oxidant.

Table 1 Published Kinetic Findings on the Partial Oxidation of Methane to Formaldehyde with Oxygen Catalyst

Rate Law

SiO 2 MoO3/SiO 2

kpo20pCH41

V205/SiO 2

kpo2~

FeNbB-O * - Refers to methane conversion

HCHO Ea / kJ mol "l 250* (507- 527~ 117* (527- 597~ 189 227 256

Ref. [7] [14] [15] [11]

Several authors have presented data on the activation energy of the reaction of methane to formaldehyde with N20 [7-9,12,13] or 02 [7,14,15] as oxidant. Those using oxygen are

1130 summarised in Table 1. Rate laws, measured under various conditions, have also been proposed in the literature [7-9,12-15]. Here the results of a new kinetic study of the partial oxidation of methane to formaldehyde, using air/oxygen are presented, along with a proposed mechanistic scheme consistent with the data.

EXPERIMENTAL The catalyst studied was 1 wt% vanadium (as metal) supported on Cab-O-Sil (M5), hereafter referred to as 1V-cabosil. The catalyst was prepared by the wet impregnation method described earlier [5]. Catalytic testing was carried out in a fixed bed quartz microreactor with on-line analysis of the reaction products, as described previously [16]. The catalyst reached steady state within minutes throughout these experiments. Products were analysed every 30 minutes, for 2 hours at each temperature. The influences of W/F, methane and oxygen partial pressures in methane rich (up to 80 kPa CH4 in the feed) and methane lean (up to 4 kPa CH4 in feed) conditions were examined at 550 and 600 ~ Negligible reaction was observed at lower temperatures. The reactant partial pressures were varied, keeping the total flow rate at 25 ml min ~, by adding helium ballast to the system. Oxygen (02) was the chosen oxidant. The oxygen partial pressure was varied in methane rich (Pcm = 81 kPa) and methane lean ( P c H 4 -- 20 kPa) conditions. The CH4 partial pressure was varied in the range 0-81 kPa (keeping the oxygen partial pressure at 20 kPa). The effects of W/F (catalyst mass used/reactant gas flow rate) were assessed, using air as the oxidant. This was done in methane rich conditions, using 0.1 g 1V-cabosil while varying the feed gas flow rate from 6.25 - 100 ml min 1. Methanol oxidation experiments were carried out in order to determine if methanol was an intermediate in the production of formaldehyde from methane. To this end a methanol saturator was placed upstream of the reactor. The saturator was submerged in an ice/acetone bath (at -16 to - 20 ~ keeping the saturated methanol partial pressure at 5 kPa. This was approximately equivalent to the total carbon containing products generated during standard reaction conditions. The gas feed stream to the saturator consisted of 81 kPa helium and 20 kPa air. The flow rate was varied from 6.25 - 100 ml min "~. RESULTS In order to determine the catalyst stability two experiments were carried out. Firstly, the catalyst was tested in 10 ~ intervals, from 450-600 ~ The catalyst temperature was than lowered in 10 ~ intervals back down to 450 ~ No significant change in activity or product distribution was noted on the downward cycle. Secondly, the catalyst was tested for 50 hours in methane rich conditions. Again with no significant change in activity or selectivity. Hence, the catalyst was deemed to be stable under the experimental conditions employed here. If the methane-air system was being used in a kinetically controlled regime the observed reaction rates would be independent of reaction mixture delivery rate to the catalyst bed. The methane conversion rate increased up to 25 ml min ~ at 550 ~ and with almost all flow rates

1131 at 600 ~ These data indicate that at 600 ~ the reaction is under diffusion control. However, at 550 ~ the reaction was seen to be operating under kinetic control, over the range of flow 25-100 ml min ]. Figure 1 illustrates the selectivity-conversion relationship between the various products obtained at 550 and 600 ~ during these experiments. It can be clearly be seen that HCHO was a primary product. Indeed as conversion increased HCHO selectivity decreased, with an analogous increase in CO selectivity. Further increase in CH4 conversion lead to the onset of increased CO2 selectivity, indicating a sequential reaction from methane to products as follows: C H 4 '-)

HCHO 41, CO r CO2 100

100 80

80 -

~60

.~60.>, 40

.v..~

40

CO

r~20

~20

--

0

1 : 2 Conversion7 %

3

1.5

i

i

v"

2.0 2.5 Conversion / %

Figure 1. Selectivity as a Function of Methane Conversion at (a) 550 ~ 0.1g 1V-cabosil, 81 kPa C H 4 and 20 kPa air

3.0

and (b) 600 ~

The Arrhenius plot for methane activation and formaldehyde production are shown in Figure 2. As can be seen from the plot shown in Figure 2, two distinct regions were noted in the Arrhenius plot for methane activation. There was a linear region which corresponded to a lower temperature range of 490-550 ~ with a second from 550-600 ~ The activation energy for methane conversion, in the temperature range 490-550 ~ was 323 kJ mol l. The activation energy for methane activation decreased towards zero over the temperature range 550-600 ~ indicating that the reaction was limited by diffusion at the highest temperatures studied. Over the same temperature range (490-550 ~ the activation energy for formaldehyde was 242 kJ mol l. These values are consistent with those previously reported by Otsuka and Hatano [ 11] (see Table 1). The plots for methane activation and formaldehyde formation were noted to be coincident at the low range of temperatures examined (480-500 ~ while some divergence between the two sets of data was noted at higher temperatures. This was due to the fact that formaldehyde selectivity was high at low temperatures, while sequential reactions became more important as the temperature was increased. The influences of reactant partial pressures on formation rates of the various products are illustrated in Figure 3.

1132

7

1.1

1.2 1.3 1/T x 10-3 / K "1

1.4

Figure 2. Arrhenius Plot for Temperature Range 480-600 ~ W / F = 8 . 9 6 x 10 "2 k g r a i n m o l "1, 81 k P a CH4 a n d 2 0

kPa air 3.0

a

HCHO+CO

CO

3.0

L~ 2.0

9

1.0

,-

'~

~

.,~

2.0-

The maximum rate ever observed, during this study, for C O 2 production was 300 ~tmol kg 1 s~ while the maxima for formaldehyde and carbon monoxide were 1300 ~tmol kg ~ s~ and 3400 ~tmol kg 1 s "1 respectively. Hence, combination of the rates for HCHO and CO was the important factor in determining the overall shapes of the total product formation curves and CO2 will not be discussed further here. The rate of evolution of formaldehyde as a function of methane partial pressure (Pen4) in the range 20-81 kPa is shown in Figure 3a. Formaldehyde and carbon monoxide rates increased fairly linearly with increasing Pen4. In methane lean (Pcm = 4 kPa) conditions overall reaction rates were 10 times lower, but very little CO or CO2 was produced (Figure 3c). The formaldehyde production rate decreased slightly for Po2 in the range 20-

~

0

20

40

60

80

100

A

CO

1.0 0.0

0.0

t []

0

'

I

5

'

I

'

10

I

15

'

t j

20

Po2 / kPa

PCH4 / k P a

0.30 "7

% 0.20 O

Figure 3. Product Formation Rates as a Function of Reactant Partial Pressures (a) 20 kPa 02 (b) 81 kPa CH 4 (c) 4 kPa CH4, T = 550 ~ W/F = 8.96 x 10.2kg min mol~.

~0.10 0.00

II

0

i

B I

25

,--

i

B

i

,--

50 75 Po2 / kPa

,

m

100

1133 81 kPa and fell more rapidly above 81 kPa. This suggested that the rate determining step, in methane lean conditions, may have involved the contact of methane with the catalyst active site, which was inhibited by the large excess of oxygen in the system. For methane rich feeds, formaldehyde and carbon monoxide production increased with increasing oxygen partial pressure (Figure 3b). An attempt was made to ascertain the reaction orders for each of the three products, namely HCHO, CO and CO2, with respect to both methane and oxygen partial pressure in methane rich conditions and with respect to oxygen in methane lean conditions. The reactions were found to be neither first nor second order with respect to either of the reactants. The difficulty in determining reaction orders may have been due to the fact that in methane rich conditions almost total oxygen conversion was achieved, resulting in a considerable gradient in oxygen concentration in the catalyst bed. The data were then subjected to the test plots for Langmuir-Hinshelwood, Eley Rideal and Mars-van Krevelen mechanisms, allowing for both associative and dissociative desorption cases. Very poor correlation coefficients were - 0.0015 obtained with all Langmuir-Hinshelwood o reaction test plots. Figure 4 illustrates a sample test plot for a Mars-van Krevelen 0.0010 (MvK) reaction mechanism, with an 550~ associative oxygen adsorption step. It was difficult to distinguish between 0.0005 o associative and dissociative oxygen adsorption with MvK reactions. It is, 0.0000 ' I ' I ' however, clear that oxygen adsorption was an 0.00 0.10 0.201 0.30 important factor in the reaction in methane 1/POE / kPa rich conditions 9 In methane rich feed conditions the data slightly favoured a model Figure 4. Sample Mars-van Krevelen Test which entailed dissociative oxygen adsorption Plot for Associative Oxygen Adsorption at as the slow step in Methane Rich Conditions For methane rich conditions the data W/F = 8.96 x 10-2kg min mol~., 81 kPa CH4 indicates that dissociative adsorption of oxygen was an important factor, based on the test plots for an Eley-Rideal mechanism. Conversely, associative oxygen adsorption appeared to be important in methane lean conditions. There was little conclusive evidence, in the form of these correlation coefficients, to chose between a Mars-van Krevelen or an Eley-Rideal type mechanism. However, our observations indicate that the most appropriate forms of the rate equations for HCHO production are the following:

6oo

METHANE RICH CONDITIONS

9 __ RHCHO kPcH4

{

PO2 I+KPo2 0.5

po 2

METHANE LEAN CONDITIONS" RHCHO = kPeH 4 1+ Kpo2

(Eqn. 1)

}

(Eqn.2)

1134

100

80

~,

~

60

-

40

r~

20

CO

o

'

30

,,~

40

50

'

,

'

60

,

'

70

I

'

,

80

90

'

100

Conv./% Figure 5. Selectivity as a Function of Methanol Conversion at 550 ~

0.1 g 1V-cabosil, 5 kPa CH3OH,21 kPa Air in feed

The role of methanol as a possible intermediate is examined in Figure 5. Figure 5 shows that methanol is readily converted to HCHO in the standard reaction conditions over 1V-cabosil. However, a number of factors indicate that methanol is not an intermediate in the reaction. Firstly, methanol was never detected under TAP reaction conditions [4,15]. Secondly, McCarthy has indicated that the rate constants for a large number of fundamental reactions of methane and its derivatives with surface oxygen species may be calculated using the extended version of the Arrhenius equation indicated in Eqn. 3. (Eqn. 3)

.

t

McCarthy has reported rate constants for the following reactions [19]" CH4 +O(surO #- CH 3 + OH(surf)

(Rxn. 1)

CH3OH + O(surOd-- CH3OH + OH(surf)

(Rxn. 2)

The values for the extended Arrhenius equation are given here [19]" ko(cH4)= 1.73 x 1012mol cm s k0(cmo~= 1.35 x 1012mol cm s n(CH4) = 0

n(CmOH) = 0

Ea(CH4) = 63.73 kJ mol 1

E~cmor0 = 46.99 kJ mol ]

These data predict more methanol than formaldehyde in the product stream during methane oxidation; this clearly is not the case. DISCUSSION

The results indicate only sequential reactions occur in the partial oxidation of methane to formaldehyde, with no evidence for the existence of parallel reaction pathways to COx. An Eley-Rideal/Mars-van Krevelen type of mechanism was found for the partial oxidation of methane to formaldehyde. The differences in the rate equations were due to differences in the amounts of oxygen present in methane rich and lean conditions. Methanol was not an intermediate in the reaction. Methanol oxidation experiments indicated that methanol was oxidised sequentially to formaldehyde, carbon monoxide and hence to carbon

1135 dioxide. However, no evidence was observed to suggest that any methanol was present during the partial oxidation of methane over the catalyst. Results from the kinetic experiments for the partial oxidation of methane indicated that a complex series of fundamental reactions leading to formaldehyde took place. Two different kinetic expressions have been derived to describe the partial oxidation of methane in rich and lean feed conditions. An attempt is made here to elucidate why this might be so. During the partial oxidation of methane to formaldehyde, methane reacted with some form of active oxygen at the catalyst surface. In methane lean experiments the reaction mixture contained up to 97 kPa O2 (96% of the total reactants in the system). In methane rich conditions the maximum amount of O2 present was 4 kPa (4% of the total). In each instance oxygen had to absorb at the catalyst surface before methane oxidation could occur. Since the chance of encountering another oxygen molecule before reaction with surface electrons, was much lower in methane rich feed conditions the rate of reaction was governed by Eqn. 1 (i.e. dissociation of oxygen occurred as indicated in Rxn. 3a). Conversely, in methane lean conditions oxygen arrived at the catalyst surface faster than electrons could be delivered, hence, the associative adsorption of oxygen was favoured (Rxn. 3b); Eqn. 2 describes the system behaviour under these conditions. This rationale affords an explanation as to why two different rate equations were found to apply in methane rich and lean feed experiments. These data are consistent with the following reaction scheme:

O2(g) + 2e~ur0 # 2Oiads )

(Rxn. 3a)

O2(g) + e(surf) e O2(ads)

(Rxn. 3b)

2CH4(g) + O (ads) ~ CH ~ + OH iads)

(Rxn. 4a)

CH4(g ) + O~ds) ~ CH~ + OHiads )

(Rxn. 4b)

2CH ~ 4- O (Latt) @ CH 3O" + 2e~ur0

(Rxn. 5)

CH30"~ HCHO + H"

(Rxn. 6)

Proposed Reaction Scheme for Formaldehyde Production The evidence to date suggests that methane does not chemisorb on the catalyst surface [4,15]. The partial oxidation of methane has been studied in a temporal analysis of products (TAP) reactor, over 1V-cabosil. A feature of the TAP system is that comparisons of residence times of various components in the reactor, and hence on the catalyst surface, can be made. Kartheuser has shown that methane and an inert gas with molecular mass - 16 g mol ~ had the same residence times in the TAP reactor, over 1V-cabosil, implying that methane did not adsorb on the catalyst surface [15]. This is consistent with the Eley-Rideal and Mars-van Krevelen mechanisms. Hence, methane from the gas phase (Rxn. 4) reacted with surface oxygen to form CH 3" radicals. It must be noted that both forms of Rxn. 4 describe the conversion of methane, but under different conditions. Reaction 4a predominated in methane lean cases, while Rxn. 4b was more relevant to methane rich conditions. Radicals generated

1136 via Rxn. 4 could further react with the lattice oxygen anions to form CH30"radicals (Rxn. 5). These radicals could subsequently decompose giving formaldehyde (Rxn. 6). ACKNOWLEDGEMENT

This work was supported in part by the EU Joule programme, contract number JOUF-0044-

c(yr) REFERENCES

[1] M. Makri, Y. Jiang, I.V. Yentakakis and G.Y. Vaneyas, Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 101 (1996) 387 [2] K. Katja, X.M. Song, and M. Huuska, Catal. Today, 21 (1994) 513 [3] F. Arena, F. Frusteri, D. Miceli, A. Parmaliana and N. Giordano, Catal. Today, 21 (1994) 505 [4] B. Kartheuser, B.K. Hodnett, H. Zanthoffand M. Baems, Catal. Letters, 21 (1993) 209 [5] M. Kennedy, A. Sexton, B. Kartheuser, E. Mac Giolla Coda, J.B. McMonagle and B.K. Hodnett, Catal. Today, 13 (1992) 447 [6] K. Heitnes, S. Lindberg, O.A. Rokstad and A. Holmen, Catal. Today, 21 (1994) 471 [7] S. Kasztelan and J.B. Moffat, J. Chem. Soc., Chem. Commun., (1987) 1663 [8] H.-F. Liu, R.-S. Liu, K.Y. Liew, R.E. Johnson and J.H. Lunsford, Jr. Am. Chem. Soc., 106 (1984)4117 [9] M.M. Khan and G.A. Somorjai, J. Catal., 91 (1985) 263 [10] N.D. Spencer and C.J. Pereira, J. Catal., 116 (1989) 399 [11] K. Otsuka and M. Hatano in "Methane Conversion by Oxidative Processes Fundamental and Engineering Aspects", E.E. Wolf (Ed.), Van Nostrand Reinhold, New York (1992) [12] M.D. Kennedy, Ph.D. Thesis, University of Limerick, Ireland (1992) [13] K.J. Zhen, M.M. Khan C.H. Mak, K.B. Lewis and G.A. Somorjai, J. Catal., 94 (1985) 501 [ 14] N.D. Spencer and C.J. Pereira, AIChE J., 33 (1987) 1808 [ 15] B.J. Kartheuser, Ph.D. Thesis, University of Limerick, Ireland (1993) [16] E. Mac Giolla Coda, M. Kennedy, J.B. McMonagle and B.K. Hodnett, Catal. Today, 6 (1990) 559 [17] G.C. Bond, "Heterogeneous Catalysis - Principles and Applications", Oxford Chemical Series, Oxford University Press, Oxford (1974) [ 18] A.W. Sexton, Ph.D. Thesis, University of Limerick, Ireland (1995) [19] J.G. McCarthy in "Methane Conversion by Oxidative Processes - Fundamental and Engineering Aspects", E.E. Wolf (Ed.), Van Nostrand Reinhold, New York (1992)

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1137

Catalytic Destruction of Volatile Organic Compounds on Platinum/Zeolite A. O'Malley and B.I( Hodnett Dept. of Chemical and Environmental Sciences, University of Limerick, IRELAND.

ABSTRACT A range of platinum (Pt) exchanged zeolites were tested for the catalytic destruction of volatile organic compounds (VOC's) in air. The range of zeolites included ZSM-5 and [3zeolite in their acidic forms and the results are compared with conventional Pt/AI203, Pt/TiO2 and Pt/SiO2 catalysts. Catalysts typically contained 0.5-6.5wt% platinum and dispersions, as measured by hydrogen chemisorption, were in the range 39-65%. Toluene and ethylbenzene, prime indoor sources of VOC's, originating from paints, adhesives and combustion products were fed to the reactor in the concentration range 200-5000ppm with excess oxygen. Typically in our reaction conditions, 100% conversion of toluene or ethylbenzene could be achieved over Pt/~-zeolite and Pt/ZSM-5 below 175~ 1.1NTRODUCTION Volatile Organic Compounds, VOC's, are an important class of air pollutants, commonly found in the atmosphere at ground level in all urban and industrial centres. Strictly speaking, the term volatile organic compounds refers to those organic compounds which are present in the atmosphere as gases, but under normal conditions of temperature and pressure would be liquids or solids. Those VOC's which are present in the atmosphere as a result of human activities, arise mainly from motor vehicle exhausts, evaporation of petrol vapours from motor cars, solvent usage, industrial processes, oil refining, petrol storage and distribution, landfilled wastes, food manufacture and agriculture. Natural biogenic processes also give rise to substantial ambient concentrations of organic compounds and include emissions from plants, trees, animals, natural forest fires and anaerobic processes in bogs and marshes [ 1]. Emission inventories are now available for VOC's for most European countries. European emissions of low molecular mass volatile organic compounds from human activities amounted to about 24 million tonnes/year in 1989. This total is comparable with levels of sulphur dioxide and nitrogen oxides (as NO2), each of which are of the order of 20 million tonnes/year for Europe [ 1]. The US Clean Air Act of 1990 calls for a 90% reduction over the next nine years in emissions of toxic chemicals, 70% of which are volatile organic compounds [2]. A protocol, signed by most of the European countries, calls for a 30% cut in VOC emissions by 1999 relative to the 1988 levels [3]. For VOC destruction, catalytic oxidation is one of the most important air pollution control techniques. In this process VOC's are oxidised over a catalyst at temperatures much lower than those required for thermal oxidation. The temperature for 100% conversion of a given

1138 VOC to CO2 and H20 varies depending on the nature and concentration of the VOC and the catalyst. Typical literature reports include a 5%COO catalyst on a carbonaceous support which totally oxidised 600ppm toluene at 210~ and 1300ppm toluene at 250~ [4]. Pt/monolith completely oxidised 500ppm toluene at 230~ [5]. Zeolites also showed a good performance for the complete oxidation of VOC's at low temperature. For example 140ppm ethene was totally oxidised at 200~ using ZSM-5 and 150~ using Cu/ZSM-5 [6]. In this work two aspects of the destruction of VOC's are investigated. The first relates to the performance of a new series of platinum exchanged zeolites and the second elucidates the structural factors which determine the ease of oxidation of two VOC's, ethylbenzene and toluene. 2.EXPERIMENTAL

2.1 Catalyst Preparation A series of Pt/A1203 catalysts containing 0.5-6.5wt% Pt were prepared by wet impregnation. Briefly, the required amount of chloroplatinic acid, H2PtCI6.6H20, (Aldrich), was dissolved in excess ethanol and added to pre-calcined Al203 (supplied by Rhone Poulenc, particle size 212-500nm). The prepared mixture was stirred constantly for 5 hours. The excess liquid was removed under vacuum at 80~ in a rotary evaporator. The resulting solid was dried in air at 100~ for 12 hours and calcined at 450~ for 4 hours [7]. 2wt% Pt supported on TiO2 (Aldrich) and 2wt% Pt supported on SiO2 (Carbosil) were also prepared by wet impregnation. ~Itte range of Pt/zeolite catalysts were prepared by ion exchange [8]. The zeolites, NH4+/ZSM-5 and NH4+/[3-zeolite (supplied by P.Q.), were converted to their acidic form by heating to 427~ in air. The required amount of Pt(NH3)4C12.H20 precursor (Johnson Matthey) was added to a suspension of H20 and the zeolite. This was stirred at reflux for 6 hours. The sample was filtered and washed three times with hot water, then dried for 12 hours at 120~ in air and calcined for 4 hours at 450~ 2.2 Characterisation The platinum contents of the prepared catalysts were measured by Atomic Absorption Spectroscopy using a Varian Spectra AA 400 plus. In preparation for analysis, all samples were dissolved in aqueous HF and only measured platinum loadings will be referred to below. Surface area measurements at 77K were performed using a Micromeritics Gemini lIl 2375 surface area analyser with nitrogen as the adsorbing gas. All samples were outgassed at 200 ~ C before analysis. Thermogravimetric analyses were carried out on a Stanton Redcrofi TG770 thermobalance with a Bausch and Lomb Omnigraphic 2000 XY Recorder. Approximately 8mg of catalyst was heated at a rate of 10~ rain1 in a flow of 30 mls rainl air from 20~ to 900~ All samples were crushed to a fine powder and analysed by X-ray diflffaction from 20 = 10% 70 ~ using a Philips Diffractometer with nickel filtered CuKa radiation (~,=1.5406A~ Platinum dispersion was measured by hydrogen chemisorption. A dynamic pulse technique was used similar to that reported by Heck et al. Briefly, 100mg of catalyst was introduced into a quartz glass vertical reactor and kept in position by 2 plugs of quartz wool. A pulse of

1139 hydrogen (1%H2 in argon) was injected into an argon stream and hydrogen uptake was monitored using a TCD. 2.3 Catalyst Testing The catalytic destruction of toluene and ethylbenzene was carried out using a continuous flow system All gases (helium, 99.99% and 29.7% oxygen in helium) were supplied by BOC gases and the flow rates were controlled by mass flow controllers (Tylan General model FC2900). Toluene was introduced into the vapour phase by passing helium through a saturator. The saturator consisted of a stainless steel tube filled with molecular sieve (4A) soaked in the organic liquid. The partial pressure of toluene in the helium stream could be controlled by adjusting the saturator temperature or by diluting the toluene/helium stream leaving the saturator with oxygen/helium from the mass flow controllers. The reactant stream consisted of 5000ppm toluene or 1000ppm ethylbenzene, 12vo1% O2 and the remainder helium, at a total flow of 50ml/min, unless otherwise stated. In all experiments 0. l g catalyst was tested using a quartz glass vertical reactor tube. The catalyst was kept in position by 2 plugs of quartz wool. The feed stream entered a series of 3-way valves which allowed the reactor to be bypassed or not as required. A K type thermocouple was inserted into the catalyst bed. The temperature of the furnace was controlled by a Eurotherm Before testing, the catalyst was pre-treated for 90 minutes in a stream of helium or hydrogen at 30ml/min at 400~ Atter pre-treatment the catalyst was cooled to the desired temperature and the reaction was started by introducing the reactant stream Product gases were analysed by on-line gas chromatography (Varian model 3400 GC) using a porapak T column and a Thermal Conductivity Detector. 3. RESULTS AND DISCUSSION An investigation into catalyst deactivation was carried out under typical reaction conditions. Briefly, the reactant stream containing 5000ppm toluene, was passed over 2.5wt% Pt/[3-zeolite at 156~ Under these conditions toluene conversion was 30% and no loss of activity occurred over a 7 hour period, in spite of the fact that the used catalyst was found by thermogravimetry to contain 10wt% coke. Figure 1 illustrates the conversion of toluene with reaction temperature for a series of platinum supported alumina catalysts. A1203 presents little activity over the temperature range studied. At 400~ a conversion of 20% was observed. Impregnating the support with platinum resulted in a significant increase in catalyst activity. A 2wt% Pt/AI203 converted 100% toluene at 250~ Increasing the platinum content decreased the temperature for 100% conversion. 100% conversion oftohene was achieved over 6.5wt% Pt/A1203 at 220~

1140

=

100-

0

9 6.5wt% PtJAI203

1,=.

W

|>

5.5wt% Pt/AI203 9 4.5wt% Pt/AI203 x 2wt% Pt/AI203

75-

O~ o .~ @

D

v

|

50-

0

AI203

25-

m

0

I-

loo

3oo Temperature

400

(oC)

Figure 1. Conversion of toluene with reaction temperature using the indicated Pt/A1203 catalysts. (PTol,=e=5000ppm, Po2=12vol%, W/F=0.06g s/ml). A similar trend was observed when ZSM-5 was used as support for various platinum contents as shown in figure 2. ZSM-5 showed no activity below 225~ and reached 25% conversion at 400~ Exchanging the ZSM-5 with 0.5wt% platinum resulted in 100% conversion of toluene at 227~ and a further improvement in activity could be achieved by increasing the platinum content to 2wt%. c 0 Itl

100

v

75-

[] 2wt% Pt/ZSM-5 9 1.9wt% Pt/ZSM-5 x 0.5wt% Pt/ZSM-5

50-

o ZSM-5

.m

L.

r0 e,,

25m

0 I--

0 100

1,.,1,~1,

I

200

v

31~0

4( )0

T em perature (oC)

Figure 2. Conversion of toluene with reaction temperature using the indicated Pt/ZSM-5 catalysts. (PTol,=e=5000ppm, Po~= 12vo1%, W/F=0.06g s/ml). A slightly different pattern was observed with the series of Pt/13-zeolites as seen l~om figure 3. 13-zeolite exhibited minimal activity even at 400~ Total conversion of toluene at 170~ was achieved by exchanging 0.5wt% Pt into the zeolite. Increasing the platinum content presented no further improvements in activity.

1141 tO W

100

o m

z~ 2.3wt% Pt/13-zeolite 9 2wt% Pt/13-zeolite x 0.5wt% Pt/13-zeolite

75

> C 0

o 13-zeolite

50

e~

25 o p.

OI 100

o 2()0

p 300

Temperature

~

400

(oC)

Figure 3. Conversion of toluene with reaction temperature using the indicated Pt/13-zeolite catalysts. (Pxol,=e=5000ppm, Po~= 12vo1%, W/F=0.06g s/ml). Figures 4(a) and (b) summarise the effect of platinum content on the temperature for complete conversion and for 50% conversion of toluene, respectively. Increasing the platinum content of alumina results in a decline in the temperature required for complete conversion (figure 4(a)). This is also true for platinum exchanged ZSM-5. However, varying the platinum content on 13-zeolite does not si~ificantly alter the temperature for complete conversion. A similar dependency on platinum content was observed for the temperature required for 50% conversion. (a) (b) 250

250

225-

2250 o o

~

tO

o Pt/ZSM-5

200-

x

~'~ 200-

9 Pt/AI203 PUp-zeolite

1-

175

150

o

i

~

~

~

Pt (wt%)

~

6

~

o Pt/ZSM-5

9PvA~o3

175150

o

X PtJl3-zeoUte

i

~,

3

~,

~

6

~

Pt ( w t % )

Figure 4. The effect of platinum content on the temperature for (a) complete conversion and (b) 50% conversion of 5000ppm toluene, respectively. (Po~=12vol%, W/F=0.06g s/ml). Table 1 compares the platinum dispersion of a 6wt% Pt/Al203 and 2wt% Pt/[3-zeolite before and after reaction. Also included is the coke content of the t~esh and used catalysts. The 6.5wt% Pt/Al203 presented a platinum dispersion of 39% while the 2wt% Pt/13-zeolite was significantly lower at 21%. Following reaction, this dispersion remained unchanged, however 10wt% coke was detected on 6.5wt% Pt/Al203 and 12wt% coke on 2wt% Pt/13zeolite.

1142

Catalyst

Dispersion (%)

6.5wt% Pt/A1203(flesh) 6.5wt% Pt/A1203(used) 2wt% Pt/13-zeolite (flesh) 2wt% Pt/~-zeolite (used)

39 39 21 21

Table 1

Coke (wt%) 0 10 0 12

Two aspects of these results will be discussed here, namely the low platinum dispersion and the presence of si,,~ificant amounts of coke on alumina and zeolite catalysts after testing. The latter finding is probably due to a combination of test temperatures below 250~ and the propensity for toluene to form coke. In any event, coke formation did not seem to inhibit the oxidation reaction to any significant extent. The low platinum dispersions, confirmed by XRD measurements, indicate that the active species may reside outside the zeolite micropores on Pt/[3-zeolite. This phenomenon was also illustrated by Smimitios et al [10]. The role of the zeolite support then needs to be addressed if its function is not to supply the dispersing medium for the platinum phase. Figure 5 compares various platinum supported catalysts for converting 1000ppm ethylbenzene. Platinum supported on [3-zeolite, TiO2 and SiO2 present similar activities, completely converting ethylbenzene at 170~ 173~ and 179~ respectively. However 2wt% Pt/AI203 required a temperature of 233~ for complete oxidation. The supports alone showed no si~ificant activity in the temperature range tested. 100

N

9 =

O

,=

"-

,.r

c

M.I

O O

"9

9 2wt%

Pt/TiO2

9 2wt% Pt/SiO2

75 50

Pt/AI203

o

2wt%

x

0.5wt% Pt/~-zeolite

25 0 100

200

Tem perature

3()0

400

(oc)

Figure 5. Conversion of ethylbenzene with reaction temperature using the indicated catalysts. (PF~hylb~e= 1000ppm, Po~= 12vo1%, W/F=0.06g s/ml). Figure 6 presents a comparison of ethylbenzene and toluene conversions over 2wt% Pt/A1203.

1143 100 C 0 W t....

75-

q)

50-

C 0

25-

(.1

0 100

150

200

2,50

300

Temperature (oC) Figure 6. Conversion of(o) 5000ppm ethylbenzene and (o) 5000ppm toluene with temperature using 0.1g 2wt% Pt/AI203. (Po2=12vol%, W/F=0.06g s/ml). This data indicates that ethylbenzene is more easily destroyed by oxidation than toluene. This finding is consistent with many literature studies which indicate variable destructibility depending upon substrate structure. Here it is postulated that destructibility is related to the bond dissociation enthalpy of the weakest C-H bond in the substrate. For ethylbenzene this correspond to the methylene C-H bond on the ethyl group with a bond dissociation of 332.3kJ/mol. For toluene the weakest C-H bond is on the methyl group with a value of 368.2kJ/mol [11]. The hypothesis is that the slow step in the VOC destruction is the generation with the aid of a catalytically active site of a radical l~agment according to H

I R2-- CmR1

I R;

R2m(2mR1 + H"

I

R3

Thereafter total oxidation proceeds via gas phase radical chemistry. The conversiontemperature profile for a given substrate depends upon the activity of the catalyst and the inherent destructibility of the substrate. Here we propose that the lower the bond dissociation enthalpy of the weakest bond the more readily the substrate can be activated by the catalyst. At the basis of this hypothesis is the suggestion that active sites in total oxidation catalysts distinguish between C-H bonds on a bond strength basis only. REFERENCES 1. 1LG. Derwent in Volatile Organic Compounds in the Atmosphere, Ed. 1LE. Hester and R.M. Harrison, Env. Sci. Tech. (1995). 2. J.N Armor in Environmental Catalysis, Ed. J.N. Armor, ASC Symposium Series 552, (1994) 298. 3. Chemistry and Industry, (Dec 2,1991) 855. 4. M.T. Vandersall, S.G. Maroldo, W.H. Brendley, Jr.I~ Jurczyk and 1LS. Drago in Environmental Catalysis, Ed. J.N. Armor, ASC S)anposium Series 552, (1994) 339.

1144

5. B.L. Mojet, M.J. Kappers and J.T. Miller, Stud. Surf. Sci. Catal. 10 (1996) 1165. 6. L.M. Parker and J.E. Patterson in Environmental Catalysis, Ed. J.N. Armor, ACS Symposium Series 552, (1994) 301. 7. J.F. Le Page in Applied Heterogeneous Catalysis: Design, Manufacture and Use of Solid Catalyst, Ed. Technip(1987). 8. P.K Alm, S.Nishiyama, S. Tsuruya and M. Masai, Appl. Catal. A:General, 101 (1993) 207. 9. 1LM. Heck and 1L Farrauto in Catalytic Air Pollution Control, Ed. Van Nostrand Reinhold (1995). 10. P.G. Smirnitios and E. Ruckenstein, Appl. Catal. A:General, 117 (1994) 75. 11. C.Batiot and B.KHodnett, Appl. Catal. A:General, 137 (1996)179.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1145

High temperature propane oxidation to reducing gas over promoted Ni/MgO catalysts. Role of impregnation condition and promoter on properties of catalysts M.V. Stankovi~ and N.N. Jovanovi6* Institute of Chemistry, Technology and Metallurgy, Center for Catalysis and Chemical Engineering, Njegogeva 12, 11000 Beograd, Yugoslavia

ABSTRACT The selectivity of variously promoted magnesia supported nickel catalyst, prepared under different impregnation conditions, has been investigated for the reaction of propane oxidation by air to CO and He. Promoted Ni/MgO catalysts have been prepared by multiple successive impregnation to give catalysts between 1 and 5 wt% nickel loadings. The catalysts samples were prepared by soaking of the MgO support in aqueous solutions of nickel nitrate and the corresponding nitrates of the metals used as promoters, followed by drying, calcination and reduction. Study of the oxygen chemisorption on the catalyst samples showed that the catalyst obtained from impregnation solution of the lowest concentration have the smallest Ni crystallites, the Ni crystallite size increases with the number of impregnation steps, and that the effectiveness of promoters to produce small Ni crystallites decreases in the order AI203>MgO>CaO. From the reaction study at temperatures in the range from 690 to 950~ performed over promoted Ni/MgO catalysts, it has been observed that selectivity towards CO+H2 formation decrease in the presence of promoter used in the order AI203>MgO>CaO, and that Ni loading in the catalyst appears to be arround 3 wt% to achieve the highest selectivity to CO and 1-/2 under experimental conditions studied. A correlation between the mean Ni crystallite size and the selectivity to CO and 1-/2 formation was established: the smaller Ni crystallites were more selective than the larger ones. 1. I N T R O D U C T I O N Reducing gas used for the treatment of metals in industry can be produced from propane in the presence of air or Oz over a supported metal catalyst with high selectivity towards CO+He formation. This process is performed at temperatures higher than 800~ There are several different reactions for high temperature propane oxidation in the presence of air [1]. The main reaction which occurs during propane oxidation with Oz (12.3% C3Hs in air) is the reaction of carbon monoxide and hydrogen formation. The theoretical

*The work was partly supported by the Serbian Ministry of Sciences and Technology.

1146

conditions for the formation of CO and 1-12from C3Hs are shown by equation: C3Hs + 3/2 02 ~ 3CO + 41-12 Aft=-227 kJ/mol

(1)

Propane can react with 02 to form complete combustion products (4% C3Hs in air): C3Hs + 502 ~ 3CO2 + 4H20

Aft=-2043 kJ/mol

(2)

Finally, C3H8 can crack to form CI-I4 and coke: C3Hs ~ 2CH4 + C

A//'=-47 kJ/mol

(3)

It is generally accepted that the overall process folows the reverse water-gas shift reaction: CO + H20 r CO2 + H2 A/-/'= -41 kJ/mol (4) As a result of these reactions a mixture of CO+H2+CO2+CI-I4+H20 is obtained which complies with thermodynamic predictions, and tend to effect complete equilibrium among all the components of the product gas. Conversions close to the equilibrium values can be achieved with considerable ease over supported Ni catalysts. To favour propane oxidation according to reaction (1) a selective catalytic material must be used. For practical purposes, nickel is usually impregnated on a suitable porous support which provides thermal stability at working temperatures [2]. But selectivity of a catalyst may depend on various other factors like composition, concentration of active component, physical and structural parameters. The effect of these parameters on the behaviour in propane oxidation of the Ni supported on mullite has been studied in our previous papers [3,4]. With this objective the present work was undertaken to investigate the effects of some parameters of the impregnation process on the selectivity of promoted Ni/MgO catalysts for the reaction of propane oxidation by air to CO and H2. The properties of catalysts relevant for their selectivity such as the Ni-loading in produced catalyst samples, the nickel surface area and mean Ni-crystallite size as well the pore size distribution of catalyst samples are presented. 2. EXPERIMENTAL 2.1. Preparation of catalyst samples Various promoted Ni/MgO catalysts have been prepared by multiple successive impregnation of the MgO support. The support used, which was calcined at 1450~ had the following characteristics: macroporous spherical granules having diameter of about 20 mm, BET surface area of 0.25 m2g 1 and specific pore volume of 0.145 cm3ga. Aqueous solutions of Ni-nitrate and nitrate of the corresponding promoter, with atomic ratio 10:1, respectively, were used. All the impregnation steps were performed at 25~ with the ratio of solution volume to support mass of 3 cmaga, and the solution/support contact time of 30 min. After each impregnation the catalyst precursors were subsequently dried at ll0~ for lh and calcined at 400~ for 2h to convert nitrate salts onto oxides.

1147 2.2. Physico-chemical characterization of samples The content of nickel in catalyst samples was determined by a standard chemical analysis, using dimethylglyoxime. The content of promoters in catalyst samples were determined by atomic absorption spectroscopy (Varian AA 775 series). The specific surface area, SB~r, of the support and of the prepared samples was evaluated by the BET method from the nitrogen adsorption isotherms determined at -196~ in a high vacuum apparatus. The specific pore volume, Vp, and the pore size distribution for the support as well as the catalyst samples were determined by mercury porosimetry (Carlo Erba, Porosimeter Model 2000). Surface area of Ni, SNi, and the mean Ni crystallite size, dNi, of the prepared catalyst samples were calculated from the oxygen chemisorption data, which were obtained by the pulse chromatographic method [5]. All the catalyst samples were previosly reduced by hydrogen at 300~ for 2 h and finished at 450~ for 1 h. The Ni surface area was calculated assuming that the chernisorption of oxygen onto the supported nickel correspond to one on the pure nickel metal. The Ni surface area was calculated assuming a chemisorption stoichiometry O/Ni~un=l and surface nickel atom average area of 0.067 nm 2. The mean Ni crystallite size was derived according to the relation: 5" 103 dNi=~ nm (5) ~Ni " SNi

where YNi is specific density of nickel, g cm 3, and S~ is surface area of nickel, m z gNi-1. 2.3. Catalyst selectivity tests The high temperature propane oxidation by air, catalysed by prepared samples, was studied in the temperature range from 690- 950~ in a flow fixed-bed quartz reactor on-line connected with an analytical system. In all the catalytic run almost equally amount of catalysts of about 40 g, and the catalyst fraction granulated from 2 to 3 mm were used. Propane and air mixture with a volume ratio of 1:7.14, were passed over the catalyst at a gas hourly spave velocity (GHSV) of 300 h a, and at atmospheric pressure. Before the catalytic tests the catalyst samples were carefully reduced in situ with propane and air mixture, at a volume ratio of 1:9.6, respectively. Analysis of C3H8, 02, CO, H2, CO 2 and CH 4 in gas mixture in the inlet as well as in the outlet of the reactor was performed using a Perkin Elmer gas chromatograph (columns: 4 m x 3 mm I.D. 60/80 Poropak Q, and a 2 m x 3 mm I.D. 60/80 Molecular Sieve 5A, both at 150~ A calibration mixture (Messer-Griesheim) was used as the reference in quantitative analysis of the product samples. The water content in the reaction products was determined by on-line connected Prolabo-hygrometer. Mass balance accurate to + 1% were obtained for all the analyses. The conversion of propane was defined as moles of propane converted per mole of propane introduced according to: n ~ - ner X= (6) o ner

1148 In relation (6) n~ denotes the molar flow-rate of propane to the reactor and nw represents the molar flow-rate of propane at the outlet of the reactor. The selectivity for each product is calculated by equation: ui "~ d n i / d t

SF = 100

%

(7)

E Di"1 dni/dt

where ui is stoichiometdc coefficient of product "i", ni is amount of the product 'T', and the quantity dni / dt is called the rate of formation of the product 'T' [6].

3. RESULTS AND DISCUSSION 3.1. Physico-chemical properties Some properties of the prepared magnesia supported nickel catalyst samples are summarized in Table 1. Table 1 List of samples studied Sample Ni loading

Promoter

SBL-r

Vp

SNi

dNi

No.

Designation*

wt%

wt%

m 2 g-1

cm 3 g-1 m e gNi"1

nm

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16

M NiAI/M-1-2 NiAI/M-1-3 NiAI/M-1-4 NiAI/M-1-5 NiAI/M-1-6 NiAI/M-2-1 NiAI/M-2-2 NiAI/M-2-3 NiAI/M-2-4 NiAI/M-3-1 NiAI/M-3-2 NiAI/M-3-3 NiAI/M-3-4 NiAI/M NiMg/M NiCa/M

0 1.10 1.62 2.13 2.63 3.13 1.13 2.12 3.04 3.87 1.79 3.19 4.15 4.95 3.31 3.24 3.32

0 0.10 0.14 0.19 0.23 0.27 0.10 0.18 0.26 0.34 0.16 0.28 0.36 0.43 0.29 ** 0.30

0.25 0.6 0.8 1.0 1.2 1.3 1.1 1.2 2.1 2.2 1.5 1.6 1.7 1.9 1.4 1.4 1.3

0.145 0.138 0.134 0.130 0.124 0.116 0.136 0.134 0.128 0.117 0.132 0.126 0.113 0.103 0.130 0.125 0.136

0

0 15.1 12.8 12.6 11.6 11.2 9.0 8.7 8.4 8.1 6.3 5.7 5.5 5.2 15.7 11.1 7.7

37 44 45 48 50 62 65 67 69 89 100 103 107 36 51 73

*Designation of the samples: The first chemical symbol behind Ni indicates the used promoter, the letter M refers the MgO-support; the f'n~t number denotes the nickel concentration in the impregnation solution and second number refers the number of successive impregnation steps. ** no determined.

1149

The Ni-loading for samples listed in Table 1, depending from Ni concentration in the impregnation solution and the impregnation steps used, varies in the range from 1.10 wt% to 4.95 wt%. Comparing data in Table 1 for the catalyst samples obtained by equal impregnation steps, but with various Ni concentration in impregnation solution, exhibit that the Ni-loading increased roghly proportional to the Ni concentration in the solution. The achieved Ni-loading within the pores of MgO-support is in the range from 66 to 54 %, in relation to the theoretical Ni-loading, which is calculated by taking into account the Ni concentration in impregnation solution and assuming the pore volume impregnation mechanism in absence of Ni2*-ions adsorption on the magnesia support surface area. Comparing data in Table 1, it is obvious that increasing the number of impregnation steps raises the Ni-loading in the catalyst samples, reduces the pore volume and increases the BET surface area. Compared with the MgO support, an increase in the BET-surface area and a decrease in the pore volume of the obtained catalyst samples can be explained by the porosity of nickel oxide deposit within the support. From data presented in Table 1, can be observed that the Ni surface area, SNi, and the mean Ni crystallite size are affected by both the Ni concentration in impregnation solution and the number of impregnation steps. Increasing the number of impregnation steps with constant Ni concentration in solution (samples No. 1 to 5; No. 6 to 9 and No. 10 to 13) decreases the Ni surface area and increases the mean Ni crystallite size in the produced catalyst samples. It can be explained by growth of the Ni crystallites as a result of deposit of Niz+ from impregnation solution on the Ni crystallites formed in the previous impregnation steps. Raising the Ni concentration in impregnation solution (see samples obtained by equal impregnation steps in Table 1) decreases the Ni surface area and increases the mean Ni crystallite size in the obtained magnesia supported nickel catalyts. Comparing data in Table 1 at the approximately constant Ni-loading in the produced A1203- promoted Ni/MgO catalysts of 3.11_+0.07 wt% (the samples No 5, 8 and 11) it is obvious that the sample No. 5 produced by 6step impregnation in the solution of 1 mol Ni2* dm 3 has the most developed Ni surface area of 11.2 m 2 gNi"1, and the smallest mean Ni crystallite size of 50 nm, howewer the sample No. 11 obtained by 2-step impregnation in solution of 3 mol Ni 2+ dm 3 have the smallest Ni surface area of 5.7 m 2 gN(,1 and the largest mean Ni crystallite size of 100 nm. The influence of promoter used on the Ni surface area and the mean Ni crystallite size in the promoted Ni/MgO catalyst samples can be observed from the values collected in Table 1 (samples No.14 to 16). These samples were prepared by soaking the magnesia support in the aqueous solutions concentration of 0.75 Mnickel nitrate and 0.075 M nitrates of corresponding promoter by required impregnation steps to provide the desired Ni-loading of about 3 wt%. In the presence of the promoter used, the Ni surface area decreases and the mean Ni crystallite size increases according to the following order: AI203>MgO>CaO. Taking into account that the particle radius of A1203 is 6.3 nm, of MgO is 20.8 nm and of CaO is 65.6 nm [7], this nonreducible promoters effect on the mean Ni crystallite size is in a good agreement with an increasing particle size of the promoter used.

1150

The derivative of the cumulative pore volume curves with respect to pore diameter for MgO support and obtained AlzO3-promoted Ni/MgO catalyst samples are depicted in Figure 1. The magnesia support has a porous structure with prevalent pore diameters in the range of macro pores approx. 8000 nm. The diagrams in Figure 1 for three series of prepared catalyst samples show that the peak amplitude in the pore diameter range of about 8000 nm decreases by increasing the number of impregnation steps at constant Ni concentration in impregnation solution, i.e with the increasing of the achieved Niloading. The pore size distribution curves for catalyst samples show a significant broadening to the pores with smaller diameters. It is interesting to note that the observed changes in the pore structure of catalyst samples are more intense for the samples prepared from impregnation solution with greater Ni z+ concentration than for the ones prepared from the diluted solution.

~

Figure 1. Derivative of the cumulative pore volume curves with respect to pore diameter for the support and catalyst samples. The number of curves corresponds to the number of samples in Table 1.

3.2. Catalyst selectivity The conversion for propane oxidation in air at temperatures in the range studied for the catalyst samples listed in Table 1 was total. CO, H2, CH4, CO2 and H20 were the only detectable products on all catalysts. Results on the study of high-temperature propane oxidation by air, catalysed b y prepared catalyst samples, are demonstrated in Figures 2 and 3. The diagrams in Figure 2 present the selectivity for products formation of main reaction in propane oxidation by air on prepared catalyst samples listed in Table 1. These results demonstrated that the selectivity for the samples prepared by multiple impregnation in 1 M Ni2+ solution increase with increasing of the Ni-loading

1151

up to 3.13 wt%. The selectivity as a function of temperature on the samples obtained by 2-step and 3-step impregnation changes rapidly up to final investigated temperature, whereas the selectivity on the samples obtained by 5-step and 6-step impregnation at the temperature of 790~ reaches about 94%, and at higher temperatures change is small. Comparing the selectivity for the samples obtained by multiple impregnation with 2 M Ni 2+ solution it can be observed that the maximum selectivity is achieved on the catalyst with 3,04 wt% Ni, i.e for the catalyst with the medium Ni-content.

1 O0

7" +

1 O0

Owl

7" +

90

90

B

B B

~, 8o

B" so ti1

~ if)

~

70

~6o

m

8

NiAI/M-2-2 NiAVM-2-3 NiAI/IVI-2-4

-e- , NiA,VM-;1-6 '

700

800

'

Temperature,~

~o

' 1 000

1 O0

90

7§

~,oo

~

900 TemperabJre,~ l

960

I O0

+

70

B 2t B'

/

"

7oo

,

"

.

l

ooo

,

1 000

,

90

8o

I1)

~

oo

70

~ o3

-3 -2

7(1

// -B- NiA~-3-3

N;o~

~

700

NiAaA-3-4 800

900

Temper,al~r'e, oC

9 - o - NiAVM ---m- NiMg~

NiCB/M

Iooo

%0

860 8~o 960 ~o 1ooo Temperature, o C

Figure 2. Dependence of selectivity of CO + H2 formation in propane oxidation on the temperature for the catalyst samples listed in Table 1.

1152 1 00

25,

20

(%1 "r 90 +

8

.B

B

B" ao

~ O3

10

70

~0

--e- NiAVM-I-6 ' 7~0 ' o~0 ' 9~0

Temperatur'e,~

"

20

I

- e - NiAVIVI-1-6 -49- NiAI/IVl-2-3 A NiAI/IVl-3-2

1

I

1ooo

~o

- e - NiAVIVl-1-6 -B- NiAI~-2-3 NiAI/lVl-3-2

03

~,[o. 7ao. o6o. o~o.1 000 Temperat;ure,~

1ooo

25

!

" i 10

mo ~oo Temperab.re,~

7m

20

- e - NiAI/IVI-1-6 -49- NiAVIVl-2-3 --A- NiAI/IVl-3-2

~1o (D (D

~]0

700

800

900

Tempera~re,~

1 000

Figure 3. Dependence of the selectivity of products formation in propane oxidation on temperature for the AlzO3-promoted Ni/MgO catalyst with Ni-loading of ~3 wt% and different mean Ni crystallite size. The samples correspond to those in Table 1.

On the catalyst samples obtained by immersion of magnesia support in solution of highest Ni-concentration, and studied in the propane oxidation, the best selectivity was exhibited on the sample prepared by 2-step immersion and with lowest Ni-loading of 3.19 wt%. On the other two samples, having greater Ni-loading, exhibit lower selectivity at all investigated temperatures. The selectivity towards CO+H2 formation in the propane oxidation on catalyst samples obtained by impregnation with more concentrated solutions was lower than

1153 the one observed on samples prepared from 1 M N i z+ solution, even in the case when the samples had greater Ni-loading. The diagrams in Figure 2 also show the effect of the promoter used on the selectivity of CO+H2 formation in propane oxidation by air. The best selectivity at all temperatures was obtained on the Al203-promoted catalyst. The selectivity for the products formation of the main reaction (1) in the presence of investigated promoter decreases in order to AI203>MgO>CaO, which is in the correlation with the previously mentioned effect of these promoters on the Ni surface area in magnesia supported nickel catalysts. The selectivity of each product formation in propane oxidation by air on the samples with Ni-loading of about 3 wt% is presented in Figure 3. These samples are chosen from each examined catalyst series ( see Figure 2) due to highest selectivity expressed towards products formation of the main reaction. The results show that the principal reaction in propane oxidation by air on these samples is reaction (1) and the other reactions occur only partially. The best selectivity is obtained on the catalyst prepared by 6-step imregnation in the least concentrated Ni2+-solution. From these results, it can be concluded that on this sample, in a rather small extent, all mentioned reactions occur: propane combustion (2), propane cracking (3) and water-gas shift (4). On this catalyst sample reactions (2) and (4) were more favoured than reaction (3), which occurs at temperatures below 900~ only. It is already observed that the selectivity to products formation of main reaction is significantly smaller on the catalyst sample prepared by immersion in solution of the highest Ni 2+ concentration. The obvious differences in selectivity among these three catalyst samples could be correlated with the Ni surface area in these samples (Table 1). Thus the selectivity to formation of products of main reaction is remarkably higher on the smaller Ni crystallites, having the greater surface area, than on the larger ones. 4. CONCLUSION Promoted Ni/MgO catalyst for the propane oxidation by air to produce reducing gas (CO and H2) is obtained by multiple successive impregnation of the low area magnesia support. The Ni-loading in the prepared catalyst is increasing with both concentration of Ni z+ in impregnation solution and the number of impregnation steps. The catalyst sample with the Ni-crystallite having the smaller size is obtained by the multiple impregnation in solution of the lowest Ni concentration. In the presence of the promoter used, the mean Ni crystallite size in the promoted magnesia supported nickel catalyst increases according to the following order: AIzO3> MgO>CaO. The selectivity to the main products formation on the catalyst in the presence of investigated promoter decreases in order to AIzO3> MgO>CaO. The selectivity to formation of main products is connected with Ni-crystallite size, i.e. on the smaller Ni-crystallite the selectivity is higher than on the larger ones.

1154 CO and H2 are formed with greater than 95% selectivity in the propane oxidation at temperatures higher than 800~ on the Al203-promoted catalyst prepared by 6-step impregnation of magnesia support in the 1 M NiZ§ REFERENCES 1. M. Huff, P.M. Tomiainen and L.D. Schmidt, Catal. Today, 21 (1994) 113. 2. W.J.M. Vermeiren, E.Blomsma and P.A. Jacobs, Catal. Today, 13 (1992) 427. 3. N.N. Jovanovi~ and M.V. Stankovi~, Appl. Catal., 30 (1987) 3. 4. N.N. Jovanovi~, M.V. Stankovi~ and G.A. Lomi~, in A.Andreev r al.(editors), Heterogeneous Catalysis, Proceedings 8th International Symposium, Vama, Institute of Catalysis, Bulgarian Academy of Science, Sofia, Bulgaria, 1996, Part 2, 553. 5. N.E. Buyanova, A.P. Kamaukhov L.M. Kefr I.D. Rather and O.N.Chemyavskaya, Kinet. Katal. 8 (1967) 868. 6. R.L. Burwell, Pure Appl. Chem., 46 (1976) 71. 7. V.V.Veselov, T.A. Levanyuk, P.S. Pilipenko and N.T. Meshenko, in Nauchnye osnovy kataliticheskoi konversii uglevodorodov, Akademiya Nauk Ukrainskoi SSR, Naukova dumka, Kiev, 1977, 84 ( in Russian).

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1155

Structural sensitivity of the oxidation reactions catalyzed by dispersed transition metal oxides: role of defect structure. V. A. Sadykov, S. F. Tikhov, S.V. T sybulya, G. N. Kryukova, S. A. Veniaminov, V. N. Kolomiichuk, N. N. Bulgakov, L. A. Isupova, E. A. Paukshtis, V. I. Zaikovskii, G. N. Kustova, L. B. Burgina. Boreskov Institute of Catalysis SD RAN, Novosibirsk, 630090, pr. Lavrentieva, 5, Russia

1. INTRODUCTION Structural sensitivity of the catalytic reactions is one of the most important problems in heterogeneous catalysis [ 1,2]. It has been rather thoroughly studied for metals, while for oxides, especially for dispersed ones, situation is far less clear due to inherent complexity of studies of their bulk and surface atomic structure. In last years, successful development of such methods as HREM and STM along with the infrared spectroscopy of test molecules has formed a sound bases for elucidating this problem in the case of oxides. In the work presented, the results of the systematic studies of the bulk/surface defect structure of the oxides of copper, iron, cobalt, chromium, manganese as related to structural sensitivity of the reactions of carbon monoxide and hydrocarbons oxidation are considered. 2. EXPERIMENTAL

2.1. Catalyst preparation As starting materials, nitrates, hydroxocarbonates, (oxo)hydroxides, oxalates, ammonium oxalatoferriate and ammonium dichromate of"pure for analysis" or "specially pure" grades were used. Oxides were usually prepared by precursors thermal decomposition at 300-400 ~ with a subsequent annealing in air or in He flow in the range of temperatures up to 1100 ~ Some details of the preparation procedures are given in [3-7]. Only samples which were found to be surface pure by SIMS and XPS were used in studies. To change defect structure of oxides in soft conditions, mechanical activation in the high-powered mill EI-2x150 was applied [8-9].

2.2. Catalysts characterisation Electron microscopy. Samples were examined in JEM-100 CX, JEM-200C and JEM-400 C microscopes. Specimens were deposited onto a carbon film supported on a copper grid [3-5]. XPD patterns were obtained with a URD-6 diffractometer (Germany) using Cu I~ radiation. The Polycrystall program was used to determine the structural parameters [6, 8]. lR-spectra of the lattice modes of oxides were obtained using M-80 spectrometer [8]. Relative densities of extended defects were evaluated by X-Ray Small Angle Scattering method (SAXS) using Cu I ~ radiation with a nickel filter and an amplitude analyzer [4,8]. Mossbauer spectra of 57Fe (including experiments with oxides doped with this ion tracer) were acquired using an NF-640 spectrometer in the temperature range 298-4.2 K [7, 10].

1156

Surface coordinatively unsaturated centers were studied by the infrared spectroscopy of adsorbed test molecules (CO, NO) using IFS-113V Bruker spectrometer [11, 12].

Heats of oxygen adsorption and amounts of a weakly bound oxygen were determined using a high temperature Calvet microcalorimeter, TPD and electrochemical method [ 13]. Catalytic properties in the reactions of carbon monoxide oxidation (all oxides) and butene oxidative dehydrogenation (iron oxides) were studied using a microreactor with the vibrofluidized bed of catalysts and pulse/flow kinetic installation [4]. Catalytic activities were characterized by the reaction rate W (molec. CO/m2s) in differential conditions and first-order rate constant K (dm 3 butene (STP)/m2.s.atm), respectively.

3. RESULTS AND DISCUSSION 3.1. Genesis of defect structure

Types of defects and their relative stabifity. Main types of defects are given in the Table 1. Table 1. Main types of defects found in dispersed transition metal oxides Oxide system

Predominant types of defects

CuO

(001 ) and (100) twins; screw dislocations along <010>, microstrains, misfit dislocation network at CuO/Cu20 interface, grain boundaries. ~-Fe203 Cation vacancies and interstitials; (0001) twins and stacking faults; (1120), (10~2) and (11~3) twins; <11~0>{0001 } screw dislocations, grain boundaries, surface steps, surface spinel precipitates. Co304 Cation vacancies and interstitials, (111) twins and stacking faults, grain boundaries, microstrains, misfit dislocation network at C0304/C00 interface MnO2 Dislocations and (100) stacking faults; intergrowth of e and 13phases. Fe304 -y-Fe203 Cations vacancies and superstructure; (110) stacking faults and twins CoO Clusters of point defects; (110) twins; surface steps, dislocations, spinel microinclusions, planar---defec--t-S---stabilizedbyi________mpurit:_____!_e. .......................... s: As dependent upon the genesis, three types of systems could be distinguished: 1. Low-temperature systems (calcination temperatures 400-500 ~ where anion residues or excess oxygen remain in the lattice, and such point defects as cation vacancies (up to several % of the regular sites occupancy) are generated due to electroneutrality requirement [5-7]. Among extended defects, twins, bulk or near-surface stacking faults and diclocations confined to the most densely packed planes predominate, suggesting their generation by the precursor phase topotactic transformation. Estimations of the extended defects density from the data of XPD combined with TEM and Mossbauer spectroscopy [3, 4, 6, 7, 10, 14] have shown that it varies from a 101 to a 10.5 per unit cell, reaching a maximum for samples prepared via thermal decomposition of oxohydroxides or nitrates. The density of the bulk planar defects was the lowest when carbonates, oxalates or ammonium dichromate were used as starting compounds [ 15, 16]. Since anion sublattice of such precursors is completely rearranged in the course of decomposition, this fact could be explained by the absence of any topotaxy in this case. The same was true when aging of the amorphous hydroxides occurs in the acid solutions, so that

1157 crystallization of oxide particles proceeds via dissolution -precipitation route. In this case, nearly perfect single crystal platelets of oxides (i.e. hematite) were obtained [ 12]. In addition to topotaxy, relative occurence of the extended defects seems to depend upon microimpurities, which are capable to effect their excess energy and, hence, stability. In such a way, residual hydroxyls are expected to diminish the effective charge of the anion sublattice, thus reducing electrostatic repulsion in the faulted regions and decreasing extended defects energies [8]. Similarly, some extended defects can appear due to oxygen deficiency, and crystallographic shear structures in ruffle and related oxides are well-known examples. For samples of Cr203 prepared by thermal decomposition of Cr (III) nitrate solution, we have revealed a new type of extended defects generated by the excess of oxygen. In this case, high oxidative potential due to nitrate anions decomposition, favors transition of up to 50% of chromium cations into 4+ state. As a result, oxide particles appear which are composed of slabs of ot-Cr203 (structural type of corundum) and CrO2 (structural type of ruffle) with thickness ca 200 A and the most developed surface faces parallel to the (101)R or (100)R planes. These slabs are stacked in rather thick (up to 1500 A) particles with the [211]R axes of symmetry (Fig.la), which have a great number of surface steps.

Fig. 1. A typical image of particle (a), HREM picture (b) and microdiffraction (c) for lowtemperature Cr203 samples prepared from nitrate. According to HREM data (Fig. l b), stacking is not coherent and possibly occurs along various directions favourable for epitaxy of these phases. Moreover, in the <111> direction, a superstructure with period ca 13 A was observed, that is rather close to ot-Cr203 cell length in the <0001> direction. Microdiffraction for such particles is very complex due to splitting of reflexes but remains of a single crystal type (Fig. 1c). After air annealing at 500 ~ CrO2 phase disappears without changing particles' shape and thickness. Particle's sizes in the (11~0) direction estimated from analysis of the halfwidths of the (11"~0) and (222[0) XPD peaks remain nearly the same as in the two-phase particles (ca 300 A). It means that (11~0) type stacking faults are formed with densities up to 0.1/unit cell. 2. Middle-temperature systems (up to 900 ~ Usually, increase of calcination temperature leads to annealing of the genetic planar defects or their reconstruction [3,4,7]. Simultaneously, for dispersed oxides, recfistallization proceeds, and extended defects of a new types are generated by sintefing, collapse of fine intraparticle pores or by reversible phase transition [6,7 ]. As a result, in this temperature region the overall density of the extended defects tends to increase (Fig. 2). Stability of extended defects was found to be greatly enhanced by microimpufities segregation in their vicinity revealed by EDX spectral analysis and Mossbauer spectroscopy of the 57Fe ion tracer [6,7,14]. As follows from the analysis of the extended defects structure [ 17 - 19], in their vicinity some cations

1158 are shifted into interstitial positions forbidden in the ideal structure. For Co304 samples with high density of the (111) stacking faults, up to 5% of cobalt cations were found to be transferred from the regular tetrahedrons into neighboring empty octahedra. For defect hematite samples, XPD data imply presence of up to 1-2 % of cations in the interstitial positions. In the IR spectra of the lattice modes of defect oxides, additional absorption bands appear. Thus, for hematite, absorption band at 570 - 580 cm1 typical to Fe in Td emerges [20].

400 O It% t~

O

O OO tt3

200

0 573

1

873

!

1173 T, K

Fig. 2. Integral density of extended defects estimated from SAXS data (1) and relative intensity of I R " defect" band (2) versus annealing temperature for hematite samples.

Fig. 3. A typical image of CoO particle with a surface layer ofCo304.

As follows from Fig. 2, the relative intensity of a such "defect" band indeed correlates with the integral density of bulk extended defects which dominate in the middle-temperature region. For Co304 with high density of defects in the (111) plane, in the 400-500 cm1 region absorbance is observed due to fragments with a local CoO structure, which is predicted for such defects. 3. High-temperature systems (ca 1000-1100 ~ where density of the bulk defects falls (Fig. 2). For some oxide systems, a great number of surface extended defects was detected. Thus, for (z - Fe203 , prismatic faces, which are atomically flat at moderate temperatures of calcination [21 ], in the high-temperature region undergo reconstruction forming steps [6]. Microimpurities segregation at the surface generates also patches with a local superparamagnetic spinel structure [ 10]. A great number of misfit dislocations are generated in particles of oxides after phase transition (Co304 - CoO, CuO-Cu20) [4,7]. 4. Soft ways to change defect structure. Mechanochemical activation was found to be very efficient in changing defect structure [8, 9, 15, 16]. For dispersed oxides, mechanochemical activation was found to generate a great number of point defects thus changing oxides stoichiometry. For such oxides as CuO, Co304, ot - Fe203, interaction of point defects with extended defects was found to "wash out" the latter via climb mechanism [8, 14, 15]. Another efficient route to change defect structure of the surface/near-surface layer is a reversible hydroxylation/dehydroxylation or oxidation/reduction of the oxides at moderate temperatures [22]. As a result, a great number of near-surface extended defects were generated. Fig. 3 demonstrates a typical moir6 picture arising due to oxidation of the surface layer of CoO particle into Co304, which shows a well-developed misfit dislocations network.

3.2. Structure of surface defect centers and bonding strength of reagents. IR spectroscopy of adsorbed CO~NO combined with isotope dilution experiments in the adlayer revealed both isolated and clustered coordinatively unsaturated cations on the surface.

1159 Discrimination between these centers was based on the next points [ 11,12]: 1) Clustered centers have a lower values of Vco and VNo due to a ligand effect; 2) Vibrations of isotopically identical CO (NO) molecules adsorbed on clustered centers are dynamically coupled, hence, clustered centers can be distinguished by the adlayer isotope dilution experiments. Integral coverages less than 5-10 % of monolayer used in our experiments proved that these centers are defects and not regular sites. 3) clustered centers are easily reduced by CO even at room temperatures giving rise to absorption bands typical to Me+/Me~ centers or subcarbonyls. A typical spectral picture illustrating isotope dilution experiment for Cr203 is shown in Fig. 4. 0,7

!

I

I

O-1 W.10

o-2

-16

D-3

B

I

K.10 5

15

30

l0

20 l0

I

2200 2150 Wavenumbers,

I

2100 cm

-I

Fig. 4. IR spectra in the carbonyl region for reduced Cr203. A- 12CO, B-13CO+12CO (7:1) mixture. 298 K.

460

600

860

10'00 T ~

Fig. 5. Rate constants of butene oxidation (1,2) and rates of catalytic CO oxidation (3) versus temperature of hematite samples calcination. Bu:O2 = 1:10 (1) or 1:1 (2).

Table 2 lists band positions for corresponding carbonyl complexes, heats of desorption estimated from the activation energies of desorption, and isotope shift values. In general, results obtained for NO as a test molecule and not shown here for the sake of brevity, correlate well with those obtained by using CO [ 12]. In most cases, for oxides pretreated in oxidizing conditions, no bands which could be assigned to complexes with the regular surface cations (Co 3+, Fe 3+, Cr 3+' Mn 4+, Cu 2+) were observed even at liquid nitrogen temperature [8, 11, 12, 14-16, 22]. It implies that cations in the regular positions of the most developed faces are effectively shielded by the oxygen anions. Isolated coordinatively unsaturated cations with decreased effective charge appear after weak reduction at moderate temperatures and/or vacuum pretreatment. The density of clustered centers was found to correlate with the density of bulk or nearsurface extended defects [8, 12, 22]. Hence, clustered centers can be assigned to surface steps including those at outlets of the bulk extended defects. Density of such centers estimated by using known values of the absorption coefficients is not higher than several percent of monolayer, broadly varying for different samples of the same oxide system [ 12, 14, 22]. Atomic models of surface centers and heats of oxygen adsorption. For main types of the most developed surface faces of oxides studied here, models of their atomic arrangement based upon minimization of the surface energy in the framework of the semiempirical Interacting Bonds Method in a slab approximation were proposed [4, 13, 23].

1160

Table 2. Characteristics of the oxides surface centers by infrared spectroscopy of adsorbed CO Oxide

Carbonyl band position, cm 1 2140 - 2150 2110 - 2120 2170 2120-2150 1980-2080

CuO CoOx

FezO3

Band assignment Cu+-CO isolated Cu + - CO clustered CoZ+-CO isolated Co1+(2+1 -CO clustered Co ~ - CO Fe z+ - CO isolated Fe z+ - CO clustered Fe 1+ - CO clustered Fe ~ - CO Cr 3+ - CO isolated Cr 3+clustered

2190 - 2200 2170 2100-2140 2060 2200 2170

CrzO3

AH of CO desorption, kcal/mol 20-25 15-17 19-20 5-7 25-30

Dynamic shift, cm 1 0 10-20 -

4-5 20 20 > 20 4-5 10-15

0 15-20 0 18

In Table 3, heats of oxygen adsorption calculated by this method are compared with the experimental values. Regular centers on the most developed densely packed planes are mainly covered by tightly bound bridged (Mz O) oxygen forms. On-top (M-O) oxygen forms are located at surface cations on more open planes where steric hindrances for MzO form exist. On-top forms also appear on the densely packed planes at surface steps including those at intersection of the bulk extended defects with the surface. MzO form can be converted into MO form by creating cation vacancy. As can be seen from the Table 3, the most weakly bound MO forms appear at cluster centers in the vicinity of a surface step comprised of surface cations in regular positions and subsurface interstitial cations. Formation of the Me-Me bonds between these cations during oxygen removal ("breathing bonds") decreases energy of the Me-O bond rupture. Estimated heats of oxygen adsorption agree rather good with the experimental values. For the majority of oxide systems, rather low amount of a weakly bound oxygen (usually, not higher than 10% of monolayer) was found. As dependent upon preparation conditions, for samples of the same phase it varies in a wide range, that agrees well with its assignment to defect centers. The situation is quite different for broadly nonstoichiometric Mn304+xand y-FezO3 spinel oxides, which have a rather big (up to 1 monolayer) amount of a moderately bound oxygen. This oxygen is assigned to MO form located at cations having neighboring cation vacancy. Table 3. Calculated and experimental values of the enthalpy of oxygen adsorption (kcal/mol). Oxide

AH calculated Regular centers

........

u/5 . . . . . . . . . . . . . . . . . . . .

Co304

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

..................

MzO: 130; MO (Oh): 40-50; MO (Td) 960-70. MO: 40-70; M20:130

Fe203 .

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

AH experimental

Defect centers

............

Centers at outlet of(110) stacking fault: 15.

Centers at outlet of(0001) fault: 18-30

.......

15; 40; 120.

......20; 37;___!__20.

1161

3.3. Structural sensitivity versus flexibility of the oxides surfaces Structure sensitivity. For most systems studied, for the same oxide phase, specific catalytic activity in the reaction of CO oxidation measured in conditions when effect of the reaction media on defect structure could be neglected, broadly varies depending upon genesis (Table 4). For iron oxides, rate constant of the butene oxidative dehydrogenation at 300 C varies from ca 0.1.10 -4 to 30 * 10.4 dm 3 Bu/m 2.s.atm, and catalytic activities in both reactions were found to correlate (Fig. 5). It means that for such oxide systems we have a clear case of structural sensitivity. For two spinel systems with a broad bulk nonstoichiometry - Mn304+x and y-FezO3, pretreated in oxidizing conditions, specific catalytic activity of various samples is rather close. In this case, apparent structural insensitivity clearly correlates with a big amount of a moderately bound oxygen (vide supra). For all other oxide systems, attempts to find any correlation between activity and the density of surface/bulk point defects were unsuccessful. Instead, specific activity was found to correlate with densities of the surface clustered defects. As an example, Fig. 6a shows such trend for the case when clustered defects concentration is characterized by normalized optical density of carbonyl complexes absorption band. Similarly, Fig. 6b illustrates correlation of specific activity with the amount of weakly bound oxygen located on clustered defect centers. In the latter case, correlation splits into three branches corresponding to various types of dominating extended defects for samples annealed in various temperature ranges [6]. For low- and middle-temperature samples, rather good correlations with the density of bulk extended defects were also observed [4, 22, 24]. Hence, structural sensitivity for the systems considered here is explained by a broad variation of the density of active centers located at surface extended defects including those at outlets of bulk defects. Flexibility of the bMk/surface structure and reaction media effect. For such systems as manganese oxides, copper oxides, spinel iron oxides Fe304-y-Fe203 [4, 5, 24, 25 ], reaction media effect at enhanced temperatures (up to 400 ~ ) and at Prolonged (up to 103 h) exposures in reaction mixtures was found to remove all initial differences in the phase composition and defect structure. All extended defects were washed out due to interaction with a flux of point defects created by reaction media. As a result, a constant level of the catalytic activity was achieved for these oxide systems demonstrating apparent structural insensitivity of the reaction of CO oxidation. Hence, in this case, great flexibility of the oxide bulk structure allows to reach the same true steady state of the catalyst. "~

"

lgw 17

/

/

.

9

20'

16 /

~ 10 ~i

15

~ P'S "103 Fig. 6a

I/z ~22

Coverage, %mon. Fig. 6b

Fig. 6. Correlations of catalytic activity of hematite samples with surface normalized intensity of the carbonyl band at 2110 cm 1 (a) and surface coverage by a weakly bound oxygen (b). 1 - low-temperature, 2 -middle -temperature and 3- high-temperature samples, respectively.

1162 In general, restrictions on flexibility are imposed by a low excess energy of extended defects, high lattice energy and a narrow homogeneity range, which hinder restructuring at typical temperatures of catalysis [24]. Such oxides as C o 3 0 4 , ot-Fe203 and ~-CrzO3 meet these criteria, and their steady-state catalytic activities vary nearly in the same limits as the initial ones. However, for these oxides as well as for all other oxides with clustered surface defects, a limited or partial flexibility of the surface layer was revealed [12, 13, 14, 22, 24]. Table 4. A scale of the rate of catalysis (W) variation in CO oxidation on dispersed oxides. Oxide

CuO

Co304

c~-Fe203 Fe304-

.............................................................................................................................................

W initial

MnOz

Mn203

Mn304

ot-Cr203

................................................................................................................................................................

4 "1016+ 5.10 TM + 1.1015 + 3"1016+ 5.1018; 3.1019; 1.1018; 2.1018; 185 ~ 25~ 227 ~ 227 ~

6"1017+ 3'1018; 140 ~

1.5"1017+ 1.4"1018; 140 ~

4'1016+ 2 "10151"1017; 8.5* 1016; 140 ~ 185 ~

W

steady- 4,1016, state 185 ~

2,1016 + 6,1014+ 6.1016+ 1.1017, 5"1017; 3"1017; 9.1016" 227~ 140 ~ ..........

227 ~ ~

.........

~ :

......

2.1017, 6.1016+ 4.1015 + 227~ 1"1017" 9"1016;

227 ~ ~ ~ .

.

227 ~ .

.

.

.__::

.

.

.

.

:::.

.

.

.

185 ~ .

.

.

.

.

.

~

The essence of this phenomenon is that properties of defect clustered centers and kinetic features of the oxidation reactions depend upon stoichiometry of the surface layer. For oxides studied, surface reduction is a topochemical type process and proceeds via spreading of the reduced zone from the extended surface defects accompanied by a cations redistribution between the regular and the interstitial positions. Reoxidation as well as hydroxylation/carbonization causes shrinkage of this zone. Table 5. Effect of oxides pretreatment on the kinetic features of low-temperature CO oxidation. ..............................................................................................

Sample

W1 a

ot-Cr203

w

E2 b

n CO 3 a b

a

b

2* 1014 3'1016

14

5-8

2* 1016 1" 10 TM

10

0

ic6;iS

.........

n 02 4 a b

W(CO) a b

0.5

0.2

0

0

1

2.4

1

0.5

0

0.5

1

100

75 ~ Co304

25 ~ a and b- pretreatment for 1 h at 400 ~ in O2 and He, respectively. For ct-Cr203, before pretreatment in He, 50 % of oxygen monolayer was removed by CO reduction at 300 ~ rate of catalysis (molec. CO/m2s); pulse regime, a steady state after 60 min contact with reaction mixture (1% CO + 1% O2 in He ) at a given temperature. 2 _ activation energy (kcal/mol); 3 and 4 _ reaction orders; 5 _ ratio of the rates of CO2 evolution in pulses of reaction mixture and 1% CO in He, respectively. 1

_

1163 The concept of a partial flexibility allowed to explain why activation energy of CO oxidation catalyzed by oxides of this type was higher when determined at a steady-state of the surface as compared with that found for a constant state of the surface (in pulse experiments) [25]. The main idea is that a temperature increase favors the surface reduction and, hence, increases the number of clustered defect centers. In some cases, variation of the properties of clustered defect centers with their degree of reduction was found to affect mechanism of the catalytic reaction of CO oxidation (transition from the oxidation-reduction scheme to a Langmuir-Hinshelwood (L-H) type). An example of this kind is given in Table 5 for Co304 and (z-Cr203. All typical features of the L-H type mechanism (high activity, low activation energies, fractional reaction orders, higher rate of CO2 evolution in pulses of CO+O2 as compared with that in CO pulses) are observed for reduced samples. 4. CONCLUSIONS Structural sensitivity manifestation for reactions of catalytic oxidation on transition metal oxides depends upon atomic structure of the surface planes, types and densities of the surface/bulk defects and structure flexibility. ACKNOWLEDGMENTS

The research described in this publication was made possible in part by Grants No RPVOOO and RPV300 from the International Science Foundation and the Russian Ministry of Science. REFERENCES

1. M. Boudart. J. Mol. Catal., 30 (1985) 27. 2. G.A. Somorjai. In" Annual Review of Physical Chemistry"', Vol. 45, p. 721. Annual Reviews, Palo Alto, CA, 1994. 3. G.N. Kryukova, V. I. Zaikovskii, V. A. Sadykov, S. F. Tikhov, V.V. Popovskii and N. N. Bulgakov. J. Solid State Chem., 74 (1988) 191. 4. V.A. Sadykov, S. F. Tikhov, G. N. Kryukova, V. V. Popovskii, N. N. Bulgakov and V. N. Kolomiichuk. J. Solid State Chem., 74 (1988) 200. 5. G.N. Kryukova, A. L. Chuvilin and V. A. Sadykov. J. Solid State Chem., 89 (1990) 208. 6. G.N. Kryukova, S. V. Tsybulya, L. P. Solovyeva, V. A. Sadykov, G. S. Litvak and M. P. Andrianova. Materials Sci. Eng. A, 149 (1991) 121. 7. V.I. Kuznetsov, V. A. Sadykov, V. A. Razdobarov and A. G. Klimenko. J. Solid State Chem., 104 (1993) 412 8. V.A. Sadykov, L. A. Isupova, S.V. Tsybulya, S.V. Cherepanova, G. S. Litvak, E. B. Burgina, G. N. Kustova ,V. N. Kolomiichuk, V. P. Ivanov, E. A. Paukshtis, A.V. Golovin and E. G. Avvakumov. J. Solid State Chem., 123 (1996) 191. 9. L.A. Isupova, V. A. Sadykov, I. A. Pauli, O.V. Andryushkova, V. A. Poluboyarov, G. S. Litvak, G. N. Kryukova, E. B. Burgina, L. P. Solovyeva and V. N. Kolomiichuk. In: "Proceedings, Int. Sem. on Mechanochemistry and Mechanical Activation, S. Petersburg, Russia, 1995.

1164

10. V. I. Kuznetsov, V. A. Sadykov, M. T. Protasova and G. S. Litvak. Izv. SO AN SSSR, Set. Khim. Nauk, 2 (1990) 112. 11. Yu. A. Lokhov, M. N. Bredikhin, S. F. Tikhov, V.A. Sadykov and A. G. Zhirnyagin. Mend. Commun., 1 (1992) 10. 12. S. F. Tikhov, V. A. Sadykov, V. A. Razdobarov, and G.N. Kryukova. Mend. Commun., 1 (1994) 69. 13. V. A. Razdobarov, V. A. Sadykov, S. A. Veniaminov, N. N. Bulgakov, V. V. Popovskii, G. N. Kryukova and S. F. Tikhov. React. Kinet. Catal. Lett., 37 (1988) 109. 14. S. F. Tikhov, V. A. Sadykov, G. N. Kryukova, E. A. Paukshtis, V. V. Popovskii, T. G. Starostina, V. F. Anufrienko, V. A. Razdobarov, N. N. Bulgakov and A.V. Kalinkin. J. Catal. 134 (1992) 506 15. L. A. Isupova, V. Yu. Alexandrov, V. V. Popovskii, E. M. Moroz, G. S. Litvak and G. N. Kryukova. Izv. SO AN SSSR, Set. Khim. Nauk, 1 (1989) 39. 16. L. A. Isupova, V. Yu. Aleksandrov, V.V. Popovskii, V. A. Balashov, A. A. Davydov, A. A. Budneva and G. N. Kryukova. React. Kinet. Catal. Lett., 31 (1986) 195. 17. P. Veyssiere, J. Rabier and J. Grilhe. Phys. Stat. Sol. (a), 31 (1975) 605. 18. Ph. R. Kenway. J. Am. Ceram. Soc., 77 (1994) 349. 19. L. A. Bursill, and R. L. Withers. Phil. Mag. A, 40 (1979) 213. 20. G. N. Kustova, E. B. Burgina, V. A. Sadykov and S. G. Poryvayev. Phys. Chem. Minerals, 18(1992),379. 21. G. N. Kryukova, A. L. Chuvilin and V. A. Sadykov. In: Materials Research Soc. Meeting Symp. Series, v. 295, p. 179-182, Materials Research Society, Pittsburgh, PA, 1993 22. V. A. Sadykov, Yu. A. Lokhov, S. F. Tikhov, G. N. Kryukova, M. N. Bredikhin, V. V. Popovskii, N. N. Bulgakov, L. P. Solovyeva, Ii P. Olenkova and A. V. Golovin. In: "Proceedings, 6th Intern. Symp. Heterogen. Catal.", Sofia, Bulgary, 1987, p. 359. 23. N. N. Bulgakov and V. A. Sadykov. React. Kinet. Catal. Lett., 58 (1996) 397. 24. V. A. Sadykov, S. F. Tikhov and V. A. Razdobarov. In: Unsteady-state Processes in Catalysis. Proc. Int. Conf., Novosibirsk, 1990. (Yu.Sh. Matros, Ed.). VSP, Utrecht, The Netherland, p. 407. 25. V. A. Sadykov and S. F. Tikhov. J. Catal., 165 (1997) 279.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1165

Oxidation of cyclohexane using polymer bound Ru(III). complexes as catalysts Jacob John, Mahesh K. Dalai and R. N. Ram* Chemistry Department, Faculty of Science, M. S. University ofBaroda, Baroda-390 002, India. Polymer bound Ru (III) complexes were synthesised by sequential attachment of chloromethyl group, 1,2-diaminopropane (DAP) as a ligand and metal chloride to styrene-divinyl benzene copolymer with 8% and 15% cross-linking. Synthesised catalysts were characterised by different techniques such as FTIR, reflectance UV'Vis spectroscopy, SEM, TGA, ESR, NMRand ESCA. Various physico chemical properties such as moisture content, bulk density, surface area and swelling behaviour in different solvents were also studied. The corresponding homogeneous complex [RuDAPCI2] CI was also synthesised. Catalytic activity of these catalysts was tested for oxidation of cyclohexane by varying the temperature of the system as well as concentration of substrate and catalyst. Values of energy of activation and entropy of activation have been evaluated from the kinetic data. A probable reaction mechanism has been proposed. I. INTRODUCTION The oxidation of hydrocarbons is an extremely important commercial reaction for functionalizing hydrocarbons to yield products that are either important in themselves or are intermediate enroute to other chemicals [ 1]. In fine chemicals, because of stringent ecological standards, more emphasis is given to oxidation by molecular oxygen or hydrogen peroxide in preference to non-.environmental friendly metal oxides. The catalytic activity of transition metal ions was reported to be low as such in homogeneous system while an enhanced activity was observed when they were heterogenized by supporting on to a solid support [ 1]. Research for viable polymer supported catalysts for laboratory and industrial oxidation reactions has received recent scientific interest [2-4]. The main problem is the leaching of the metal ion from the surface of the support when it is immobilized by the use of the monodentate or non-chelating ligands. We have reported catalytic activity of different heterogenized chelated metal complex catalysts for various hydrogenation reactions [5-10]. Covalently attached polymer bound multidentate amines could bevaluable starting material to synthesise polymer bound chelates and macrocycles [ 11]. Present study deals with synthesising the polymer bound Ru (III) complexes using 1,2-diaminopropane (DAP) as a ligand and oxidation of cyclohexane under mild reaction conditions using above catalysts. 2. EXPERIMENTAL 2.1 Materials and Equipments Styrene, divinyl-benzene (DVB), dioxane, methanol and cyclohexane were purified according to published methods [12] Styrene-divinylbenzene copolymer (XAD-8) was procured from Fluka

1166 AG, Switzerland. 1,2-dichloroethane and 1,2-diaminopropane were distilled before use. Aluminium chloride was purified by sublimation. RuCI3.3H20 (Lobachemie, Bombay) was used as received. Elemental analyses and TGA were carried out in our laboratory on a Coleman Analyser and a Shimadzu Thermal Analyser DT-30, respectively. The surface area of the support as well as the catalyst was measured using a Carlo Erba Strumentzion 1800. Swelling studies of the catalysts were carried out using polar and nonpolar solvents atconstant temperature. The detailed procedure has been described earlier [ 13]. UV-Vis reflectance spectra of the solid samples were recorded on a shimadzu UV-240 spectrophotometer with reference to nonabsorbing BaSO 4 as a standard and liquid sample in methanol. FTIR and NMR were recorded on a Perkin Elmer R-32 instrument. EPR spectra were scanned on a Breaker ESP 300k band spectrometer using a 100 KHz field modulation on powdered samples at 298 K in N 2 atmosphere. Scanning electron micrographs were recorded on a Jeol SJM T-300. ESCA was recorded on VG model ESCA-3 Mark (II) U.K. with A I K and MgK~ as the radiation sources.

2.2 Synthesis of the Catalysts Styrene-divinyl benzene copolymer with 15% crosslinking was synthesised by the suspension polymerization technique using benzoyl peroxide as an initiator [5]. After polymerization the beads were washed with distilled water, water ethanol (1:1) mixture and ethanol. It was finally, soxhlet extracted with ethanol benzene (1:1) mixture. In a separate series of experiment commercially available styrene divinyl benzene copolymer (XAD-8) was taken as a polymer support. Polymer beads were chloromethylated with HCI, para formaldehyde and acetic anhydride using 1,2 -dichloroethane as a solvent using AICI3 as a catalyst [ 14]. Chloromethylated XAD-8 copolymer was modified by introducing 1,2-diaminopropane.as ligand. The method is reported earlier in detail [5]. The functionalized polymer beads were kept in contact with an ethanolir solution of RuCI3.3H20 (0.4% w/v) for 7 days. There was a change in the colour of the supematent liquid from dark orange to light orange after complexation and beads turned to light grey indicating formation of metal complex on to the polymer matrix. The metal content was determined by refluxing the polymer supported catalyst with cone. HCI (AR) for 24 hrs and then estimating the metal concentration in the solution by spectrophotometfic method after complexation with a nitroso-R salt [ 15]. Synthesised catalysts were named as NPML where N=percent crosslinking, P = styrene-divinyl benzene copolymer, M = Metal (Ru) and L=Ligand (DAP). The following catalysts were prepared. Catalyst B = 15PRu(III)DAP Catalyst F = 8PRu(III)DAP 2.3 Kinetics of Oxidation The kinetics of oxidation of cyclohexane was studied at atmospheric pressure by measuring oxygen uptake using a glass manomatric apparatus. The initial rate was calculated from the slope of the plot of oxygen uptake at various interval of time. The detailed procedure and experimental set up are described earlier [ 12]. The products were analyzed by gas chromatograph .. Oxidation of cyclohexane produced two products, cyclohexanone (1) and cyclohexanol (El). Catalyst B gave 11.2% I and 20.4% Ill whereas catalyst F produced 12.3% I and 19.8% H.

1167 3. RESULTS AND DISCUSSION 3.1 Characterization of Catalysts Physicochemical properties of the supported catalysts are given in Tables 1-3. A decrease in the surface area was observed after loading the metal ions. This is in accordance with the earlier results [5, 9 and 12]. This might be due to blocking of pores of the polymer support after introduction of the ligand and the metal ions. The change in morphology of the polymer support after ligand and metal ion introduction was observed by SEM (Fig. 1) Successful functionalization of the polymer was confirmed by elemental analyses at different stages of preparation of the catalyst. Catalyst F (8% cross linked) was found to have more swellability compared to catalyst B (15% cross-linked). A decrease in swelling was observed as the nature of the solvent was changed from polar to non-polar. Methanol was found to be suitable swelling agent and employed for oxidation reaction because of better swellability with the catalyst and miscibility with the substrate. The UV-Visible reflectance spectra showed d-d transitions at 210 and 500 nmmi.'ght be due to Ru(III). ESCA studies of the polymer bound ruthenium catalysts gave peaks due to Ru (3d 3/2), Ru (3p 3/2) N (1 S), CI (2p 3/2) and C (1 S) for ruthenium - DAP, indicating +3 oxidation state of the metal ion. The unbound complex [Ru DAP C12]CI was characterized by ESR. The gr gH and gov values were found to be 1.897, 2.504 and 2.360 respectively indicating thereby the presence ofgu in low spin +3 oxidation state. The NMR spectrum of the ligand 1,2-DAP showed 2 peaks at 1.00 and 2.42 due to amino and methylene protons respectively. NMR of the unbound complex [Ru DAP C12] CI showed a shiR in peak by 0.2 due to methylene group while amino proton appeared in the same region with multiple splitting which indicates different electronic environment ofligand after complexation. The formation of metal complex on the surface of the polymer was confirmed by FTIR. The various infrared frequencies were assigned as shown in Table 4. TG analysis indicates that thermal stability of the polymer support increases on increasing cross linking of the polymer however no significant change .was observed when metal complex was loaded on the support. Initial weight loss observed might be due to moisture content. It was concluded that catalysts could be used safely upto 100~ (Fig. 2). A probable structure of the catalyst has been proposed based on the spectroscopic data (Scheme 1).

Table 1 Physicalproperties of the supported catalyst Physical property Pore volume (cm 3g-l) Surface Area (m~g-l) Apparent bulk density (g m3) Moisture content (wt%)

Catalyst B 0.398 (0.425) 26.842 (27.932) 0.50 1.28

Catalyst A 0.569 (0.525) 149.60 (160.00) 0.39 0.38

(Values for the respective polymer support are indicated in the parenthesis)

1168 Table 2 Elemental analysis at different stages of preparation of.catalyst B and F x C

87.21 84.28

Y

z

H

CI

C

H

N

C

7.35 7.04

8.9 4.8

84.21 83.33

7.40 6.98

2.59 1.77

83.92 82.88

H

N

7.36 6.77

2.53 1.53

Ru

3.25 x 10.2 1.57 x 10.3

x = after chloromethylation; y = after ligand introduction; z = after complex formation. Table 3 Swelling studies of the supported catalysts Swelling (tool %) Solvent

Catalyst B

Water Methanol Ethanol Dioxane DMF Acetone THF Benzene n-Heptane

Catalyst F

0.692 0.611 0.402 0.198 0.137 0.110 0.098 0.086 0.040

0.962 0.610 0.601 0.497 0.442 0.420 0.411 0.308 0.095

Table 4 IR frequencies (cm -~) Catalyst

Ru-CI

Ru-N

B

225

300

C-N 1072

3407 1598

1170

F

240

314

1069

3437 1633

1272

/CH3 Ci~ CI mm,-R . t t ~ CIj ~NH:Scheme 1

N-H

-CH2CI

1169

Fig. 1 Scanning electron micrographs of (a) P (S-DVB) with 15% cross-linking (b) Catalyst B

1170

3.2 Oxidation reactions The kinetics of oxidation of cyclohexane for polymer bound catalysts B and F were investigated. The stirring of the reaction mixture was maintained at an optimised rate (700 rpm) throughout the experiment to minimise diffusion. The reaction was carried out in a kinetic regime at atmospheric pressure and in the temperature range of 30-45~ The influence of various parameters on the rate of oxidation was studied (Tables 5-7). 3.2.1 Influence of Cyciohexane concentration The effect of substrate concentration on the rate of oxidation was determined in the range of 5.94 x 103 to 23.00 x 10-3 mol 1~ at 35~ and 1 atm. pressure at a constant catalyst concentration of 3.22 x 10.5 mol 1-! of Ru(III) for catalyst B and 1.55 x 10.6 mol 1-~ of Ru (III) for catalyst E It was observed that the rate of oxidation increases linearly with respect to substrate concentration. The order of reaction calculated from the linear plot~; of log (initial rate) vs. log [cyclohexene] was found to be fractional for both catalysts. 3.2.2 Influence of catalyst concentration The effect of catalyst concentration on the rate of oxidation was studied in the range of 1.61 x 10.5 to 6.43 x 10.5 mol 1-! of Ru (III) for catalyst B and 0.776 x 10.6 to 3.100 x 1 0 .6 mol 11 of RU(III) for catalyst F at constant substrate concentration of 11.80 x 10-3 mol 1!, 35~ and 1 atm pressure (Tables 5 and 6). The order of reaction calculated from the plots of log (initial rate) vs. log [Catalyst] was found to be fractional with respect to catalyst concentration for both the catalysts. This may be due to non accessibility of catalytic sites, as well as steric hinderance because of complex nature of the catalyst [5]. 3.2.3 Influence of temperature Catalytic oxidation was studied over a range of 30-45~ at a fixed catalyst concentration of 3.22 x 10.5 mol 1-~ ofRu (III) for catalyst B and 1.55 x 10.6 mol 1! ofRu (III) for catalyst F, at 35~ 1 atm pressure and a substrate concentration of 11.80 x 10.3 mol 1-!. An increase in the rate with temperature was observed. The values for energy of activation calculated from the slope of the plot of log (initial rate) vs. 1/T (Fig. 3) were found to be 5.86 and 9.36 Kcal mol t for catalyst B and F respectively; corresponding entropy of activation was found to be-59.91 and -45.44 eu. 3.3 Oxidation of Cyciohexane using unbound complex [Ru DAP CI2]CI Oxidation of cyclohexane was also studied using homogeneous complex of Ru(IIl) with 1,2 DAP under similar condition (Table 8). However for convenience same quantity of catalyst could not be used as the same quantity of unbound catalyst gave immeasurable.oxygen uptake. Inspite of using larger amount of Ru (III), a lower reaction rate was observed as compared tO polymer supported catalyst. Effect of various parameters such as concentration of substrate and catalyst, temperature, amount and nature of solvent is seen and the results are summerised in Table 8. The energy of activation was found to be 7.04 Kcal moP. 4. RATE EQUATION The reaction mechanism for the oxidation of olefins by metal ions / complexes in homogeneous medium is studied widely and the formation of peroxo and oxo complexes was suggested to be responsible for the transfer of oxygen to the substrate. Vaska et al., have reported formation of peroxo complexes when dioxygen is bound covalently to the metal centre [ 16]. The formation of oxo complex and the transfer of oxygen via this route has been suggested by Taqui Khan et al.; in the oxidation of olefins catalysed by Ru(III) complex in homogeneous medium [ 17]. On the basis

1171 of experimental results as well as evidence from the literature, the following mechanism and rate equation are proposed. /O--O Ru (III) complex +02 ~ Ru(IV) R~uu(IV) m > Ru(V) = O Ru (V) = O + substrate

> product + Ru(III) complex

Keeping the amount of 02 constant, the rate law may be written as 9 R' = k [Catalyst] [Cyclohexane] Thus on increasing the amount of the catalyst as well as the concentration of cyclohexane, an increase in the rate is observed (Tables 5 and 6).

Table 5 Summary of the kinetics of oxidation ofcyclohexane by polymer bound catalyst B in 20ml methanol at 1 atm pressure. [Ru(III)] [Cyclohexane] Temp Rate of Reaction (mol 1l) 105 (mol 1~) 103 (~ (ml min-I) 102 3.22

5.94 11.80 14.80 17.80 23.00

35

2.58 2.66 3.05 3.24 3.65

3.22 4.01 4.83 6.43

11.80

.35

2.66 3.85 4.25 4.64

3.22

11.80

30 35 40 45

2.52 2.66 3.50 3.93

5. CONCLUSION Polymer bound Ru(III)-DAP complexes were found to be stable upto 100~ These catalysts were found to be effective for oxidation of cyclohexane under mild operating conditions. The rate of reaction was studied by varying different parameters and the order of reaction with respect to [catalyst] as well as [substrate] was found to be fractional for both the catalysts. This

1172 Table 6 Summary of the kinetics of oxidation of cyclohexane by polymer bound catalyst F in 20ml methanol at 1 atm pressure [Ru(III)] (mol I-') 10 6

[Cyclohexane] (mol 1") 103

Temp (~

Rate of Reaction (ml min-') 102

1.55

5.94 11.80 14.80 17.80 23.00

35

2.50 2.98 3.72 4.10 4.50

0.776

11.80

35

2.77 2.98 3.78 4.40 4.60

1.55

11.80

30 35 40 45

2.79 2.98 3.56 4.61

might be due to non-availability of active sites as well as steric hinderance. The entropy of activation calculated from the kinetic data was found to be -59.91 and -45.44 eu for catalysts B and F respectively which indicates loss of freedom due to fixation of catalyst molecules on the polymer matrix. The activity of these catalysts for the oxidation.of cyclohexane was observed to be higher than their homogeneous counterparts. A probable reaction mechanism is also proposed.

C

.'xt/1

7I-

__o !

-1.4

v

o

2O 200

Temp.(C)

400

Fig. 2 DTA-TG curves of (P) P(S-DVB) with 15% cross-linking (C) catalyst B

15

3.20

I/Txl6

Fig. 3 Arrhenius plots for catalysts B and F

3.30

1173 Table 7 Summary of the kinetics of oxidation of cyclohexane by homogeneous complex [Ru DAP C12]C1 [Ru] (mol 1!) 103

[Cyclohexane] (mol 1!) 103

Temp. (~

Rate of reaction (ml min-~) 102

1.60

5.94 11.80 14.80 17.80 23.00

35

2.01 3.16 3.34 3.50 3.98

0.643 0.803 0.960 1.280 1.600

11.80

35

1.90 1.99 2.08 2.18 2.34

1.60

11.80

30 35 40 45

2.81 3.38 3.98 4.84

Acknowledgement Authors would like to thank Prof. A.C. Shah, head, Chemistry department and R & D, IPCL for facilities as also to UGC, New Delhi for financial support to one of us (JJ).

REFERENCES 1. 2. 3. 4. 5.

F.R. Hartley, Supported Metal Complexes, Reidel, Dordrecht, 1985. M. M. Miller and D. C. Sherrington, J. Catal., 152 (1995) 368. D. C. Sherrington and H. G. Tang, J. Catal, .142 (1993) 540. W. Derong and S.Licai, Chem. Eng. Sci. (1992) 3673. Jacob John, M. K. Dalai, D. R. Patel and R. N. Ram, J. Macromol. Sc. Pure Appl. Chem. A.34 (3) (1997) 409. 6. M. K. Dalai and R. N. Ram, European Polym. J. 1997, (In Press). 7. D. R. Patel, M. K. Dalai and R. N. Ram, J. Mol. Catal. A : Chemical, 109 (1996) 141. 8. J. John and R. N. Ram, Polym. Intl., 34 (1994) 369. 9. J.N. Shah, D. T. Gokak and R. N. Ram, J. Mol. Catal., 60(1.990) 141, 10.D.T. Gokak, B.V. Kamath and R.N. Ram, React. Polym. 10 (1989) 37. ll.R.S. Drago and J. Gaul, J. Am. Chem. Sot., 102 (3) (1980) 1036. 12.B.S. Furniss, A. J. Hannaford, V. Rogers, P. W. G. Smith and A. R. Tatchell (Eds) Vogel's Textbook of Practical Organic Chemistry, 4th Ed., ELBS and Longmann, London, 1978.

1174 13. D. T. Gokak, B. V. Kamath and R. N. Ram, J. Appl. Polym. Sci., 35 (1988) 1523. 14. J. D. Spivack and S. Vauey, U. S. Pat. 3 281 505 (1966). 15. A. K. Singh, M. Katyal and R. P. Singh, J. Ind. Chem. Soc., 53 (1970) 691. 16. L. Vaska, Acct. Chem. Res., 9 (.1976) 275. 17. M.M. Taqui Khan, Ch. Sreelatha, S.A. Mirza, G. Ramchandraiah and S.H.R. Abdi, Inorganica Chimica Acta, 154 (1988) 103.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1175

Photoimmobilized catalysis for low-temperature oxidation of olefins L.V. Lyashenko, V.M. Belousov, E.V. Kashuba L.V.Pisarzhevskij Institute of Physical Chemistry, National Academy of Sciences, Ukraine, 252028, Kyiv, Prospekt Nauki, 31

Heterogeneous metal complex catalysts, synthesized by a new photoimmobilization method, for low-temperature oxidation of olefins C3-C4 to partial oxidation products are proposed. The process selectivity is about 100 %.

1. INTRODUCTION The catalytic oxidation of light olefins on complex oxide catalysts is known to proceed at temperatures above 630 K. Promising catalysts for the low temperature oxidation of hydrocarbons can be heterogeneous metal complexes whose efficiency is largely determined by the method of their preparation. We have developed a new nontraditional method to synthesize heterogeneous metal complex catalysts, which distinguishes by using UV-irradiation directly in the process of complex immobilization on the support[ 1-3]. Here we present the results of studying the oxidation of propene to acrolein and isobutene to methacrolein on metal complex catalysts synthesized through photochemical reactions. We have established that the application of photoimmobilized metal complex catalysts decreases the reaction temperature from 630 K to 300-320 K at selectivity about 100 %.

2. EXPERIMENTAL

2.1 Preparation of catalysts. The catalysts were prepared by the immobilization of metal ions on the support surface during metal photoreduction under UV irradiation from the salt solutions in proton solvents in the presence of benzophenone. The reaction was carried out in a molybdenum glass vessel using mercury lamp under stirring for 2 h. Large pore silica gel with a surface area of 380 mE/g and particle diameter of 0.05 mm was used as a support. After settling the precipitate was washed out by decantation with isopropanol and dried at 365-370 K in flowing argon. The catalysts thus prepared will be referred to as photoimmobilized ones (Me-Ph). For comparison, non-irradiated immobilized catalysts were synthesized under similar conditions (Me-Iml) as well as ones from solution in CC14 (Me-Im2).

1176 We also prepared supported catalysts by the impregnation of silica gel in isopropanol with metal salt, followed by its decomposition under heating (Me-Sup). The heterogeneous metal complex catalysts containing Cu, Mo, W, V, Ti, Ag were examined. 2.2 Experimental technique. Catalytic activity was measured in a flow reactor. Reaction mixture to test catalysts was 1,22 vol.% of hydrocarbon in air. The oxidation products were analysed chromatographically. To gain information on metal surface compounds of both photoimmobilized and immobilized samples their diffuse reflectance spectra in IR, visible and UV region were measured using spectrophotometer "Hitachi-340". ESR spectra of catalysts also were examined using radiospectrometer "Varian-E-9".

3. EXPERIMENTAL RESULTS AND DISCUSSION 3.1. Isobutene oxidation. The data obtained on isobutene oxidation are represented in the Table 1, Figures 1 and 2. To compare the catalytic activities of the samples synthesized by various methods and containing different amounts of metal, we calculated the atomic catalytic activity (ACA), i.e. the rate of methacrolein formation (mol/s) per one g of metal. Table 1. Oxidation of isobutene to methacrolein on metal complex catalysts prepared by various methods Catalyst Temper.of Selectivity and ACA reaction yield of meth106 start,K acrolein, % mol/s g Me Cu-Ph 290 100 25 4.2 Cu-Iml 438 34 21 0.7 Cu-Im2 453 32 11 0.3 Cu-Sup 433 40 12 0.6 W-Ph W-Iml W-Im2 W-Sup

305 310 400 490

100 90 35 0

40 13 1 0

5.2 3.6 0.4

Mo-Ph Mo-Iml Mo-Im2

293 293 330

100 100 80

28 11 4

5.5 1.6 0.1

V-Ph V-Im2 V-Sup

295 346

100 12 74 2 inactive till up 510 K

1.4 0.1

1177

100

20

Q-

2

o

50 ~O r-~ G) ,H .H rH o 0 cO 4-~ r

5

'~ r-I g)

323

373

/.23

2O

/.73 T(K)

523

573

100

0

0,1

!

15 10

50

323

373

/.23

/.73 T(K)

523

573

Figure 1. Temperature dependence of methacrolein (1) and C02 (2) yield on Cu-Ph (a) and Cu-Iml and Cu-Sup (b) catalysts. It can be seen that photoimmobilized catalysts are quiteadvantageous as compared to other ones. A specific feature of isobutene oxidation on photoimmobilized catalysts obtained consists in the fact that this process takes place almost at room temperature. For example, on Cu-Ph reaction starts at 290 K, with increasing temperature the amount of methacrolein formed rises, attains its maximum value at 355-365 K and then at 430 K it decreases. The maximum degree of isobutene conversion is 25-30 %. On Cu-Iml and Cu-Sup isobutene oxidation starts at 420-450 K (Figure 1). Partial oxidation products (methacrolein, acetaldehyde, acetone - -~ 0.05 vol. %) are formed with maximal rate at 520-550 K. Above 570 K isobutene is oxidized to CO2 and H20. The similar results are obtained over Ti-containing catalysts. In the presence of Ti-Ph isobutene is oxidized to methacrolein in the range of 330-420 K and to COz and H20 at 430-520 K. The Ti-Iml, Ti-Im2 and Ti-Sup are inactive at low temperatures, producing CO2 and H20 at the temperatures above 520 K.

1178

A/ t0 3o

mol/s

gMe

t 20

2

10

,3 z/oo T,/4 Figure 2 Temperature dependence of methacrolein formation rate on Mo-Ph (1), Mo-Iml (2) and Mo-Im2 (3) catalysts. The ACA value for Mo-containing photoimmobilized samples (Mo-Ph) is the same order of magnitude as that for the catalysts immobilized from isopropanol (Mo-Iml). However the methacrolein yield relative to the initial isobutene on Mo-Ph is much higher than on Mo-Iml. ACA of the catalysts, synthesized in isopropanol (Mo-Iml) is higher than that for the catalysts prepared from CC14 (Mo-Im2). On Mo-Sup samples isobutene undergoes no oxidation at temperatures of up to 490 K, starting from 510 K a cracking process takes place. The oxidation of isobutene on W- and V- containing catalysts proceeds in the same way as on Mo-containing samples. Another peculiarity of isobutene oxidation on Cu, Mo, W, V, Ti catalysts is the high selectivity towards methacrolein formation (100 % selectivity is observed at 300-390 K). In the region of maximal yield ofmethacrolein there is no even trace amounts of CO2. In some cases (Mo-Ph, W-Ph) acetone and the cracking products are observed at temperatures above 390 K.

3.2. Propene oxidation to acrolein The oxidation of propene on photoimmobilized Cu-Ph, Mo-Ph and W-Ph catalysts occurs at a measurable rate already at room temperatures. The selectivity to the partial oxidation product, acrolein, proved to be 100 %. The rate of formation of the reaction products on the above catalysts against temperature is represented in Figure 3. The maximal rate of the acrolein formation is observed on Cu-Ph, Mo-Ph and W-Ph at 345, 350 and 340 K respectively. Above these temperatures the photoimmobilized catalysts lose their activity.

1179

[ -w-.to 400

SO

~ ~, f,

300

,,

,|,,

35O

Figure 3. Temperature dependence of acrolein formation rate on Cu-Ph (1), Mo-Ph (2) and WPh (3) catalysts. Within 310-430 K the rate of propene oxidation on Me-Im2 and Me-Sup is less by order of magnitude as compared to that on Me-Ph. At constant temperature the activity of the photoimmobilized catalysts decreases gradually. Temperature rise from 300 to 360 K leads to a new increase in the rate of partial oxidation production. Time decrease in the activity of these catalysts at constant temperature is described by the equation W=A lnt, which is typical for processes with a variable number of active sites. Decrease in the activity of Me-Ph catalysts after attaining the low-temperature maximum can be ascribed to reoxidation of photoreducted metal ions. 3.3. Other oxidation reactions The supported Ag-containing compounds are known to be typical catalysts for production of the ethene oxide from ethene. It was interesting to examine Ag-Ph catalysts in this reaction and to compare them with Ag-Im2 and Ag-Sup. It was established that ethene oxide, CO and CO2 were formed on Ag-Sup at temperature range of 370-410 K. Such the products were formed at the presence of Ag-Im2, although the reaction took place at more high temperatures (450-500 K). The major product of partial oxidation on Ag-Ph was butene-2,3 oxide. It has been formed begining from 340 K with 100 % selectivity and ethene conversion being of 30-35 % at 408-415 K. The ethene oxide was detected in trace amounts.

1180

/

/ I

/

doo gOo

..

' mJ

r /,7/,5 r / L

9

|

i

Figure 4. The diffuse reflectance spectra of W-containing catalysts 91,1"- W-Ph, 2,2"- W-Im2, 3 -WO3. Mo-and W-containing systems were examined in the reaction of n-butane oxidation. Mo-Ph and W-Ph catalysts were found to be active in this process. There were two temperature ranges for n-butane oxidation - the low temperature one at 310-410 K and the other at the temperature higher than 520 K. In the high temperature range the reaction products were acetaldehyde, CO2 and HEO. At low temperatures only one product of oxidation was formed. By the methods of chromatography, field mass-spectrometry, IR and NMR-13C spectroscopy it was determined this product to be an ethene oxide. The Mo-Iml, Mo-Im2, Mo-Sup were active only at high temperature range, to produce the products of deep oxidation only.

3.4. The nature of the active centers of heterogeneous metal complex catalysts. The diffuse reflectance spectra in visible and UV regions indicate, that support surface contains metal as separate complexes or their associates in all systems studied: the absorption edge for the immobilized samples is shii~ed towards the short-wave region in the spectrum as compared to that of solid oxides (for example, Figure 4). The near IR spectrum of the samples of Me-Im2 exhibits two bands at 1.37 and 1.90 I~ that are ascribed to vibrations of the OH-group in H20 molecule. For the Mo, V, W catalysts immobilized from isopropanol, besides these bands, the spectrum also exhibits 1.42, 1.681.801~ bands attributed to C-H vibrations of the -CH3 and _CH groups and a 1.32 !~ band

1181 ascribed to vibrations of the OH-group in the alcohol molecule. Hence it can be suggested that the active compound in these samples is present in the form of separate complexes, whose coordination sphere involves water and alcohol molecules. Since these complexes are not washed out by solvent, one may suggest that they are strongly bonded to the support surface. We have also examined ESR spectra of the catalysts. These spectra indicate the presence of W 5+ and Mo 5+ions in the systems studied. ESR signal for the samples synthesized from CC14is negligible. The data obtained permit to state that the site of the low-temperature active center is W 5+ and Mo 5+, involved in the WO 3+ and MoO 3+ oxo ions. The intensity of ESR signal correlates qualitatively with the activity of W- and Mo-containing catalysts in the sequence: Me-Ph > Me-Iml > Me-Im2. ESR spectra of V-containing systems show a SFS structure, characteristic ofvanadyl ions in an octahedral environment. Chemical analysis indicated the presence of Cu + ions on Cu-Ph catalysts and nearly complete absence in immobilized and supported samples. Thus we may conclude that the low-temperature catalyst surface contains metal ions (electronic and ESR spectral data) in the form of complexes whose coordination sphere involves water and solvent molecules (IR spectral and elemental quantitative analysis). On the base of spectroscopy studying the composition of surface compounds may be offered as

CH3 Si ~

O ~

(OU)x

C ~

O...Me (n-1)+/

CH3

~

(UO)y

(ROH)z

The correlation between spectra of the solution from which the immobilization was carried out and diffuse reflectance spectra of the catalysts shows first forming of the reduced metal complex under UV-irradiation followed by its immobilization on the silica surface.

4. CONCLUSION By the photoimmobilization of Cu, V, Mo, W, Ti and Ag complex ions on silica gel the systems including the metal ions in a low oxidation degree were synthesized. Their coordination sphere was consisted of hydroxy ions, water and isopropanol molecules. Such systems possess unique catalytic properties. They catalyse the hydrocarbon selective oxidation at sufficiently low temperatures: isobutene at 300-320 K, propene at 310-350 K, n-butane at 330-350 K. The significant selectivity towards partial oxidation products (about 100 %) is observed at the quite high hydrocarbon conversion catalytic ( 30 - 40%). A special feature of photoimmobilized systems is their ability to catalyse the reaction in unexpected direction as compared to immobilized and supported catalysts of the same chemical nature. Two new catalytic reactions, proceeding over the photoimmobilized catalysts at a low temperature, have been discovered. One of them is the n-butane oxidation to ethene oxide over

1182 Mo-Ph and W-Ph catalysts and another one is the formation of butene-2,3 from ethene over Ag-Ph systems. To our mind, such properties of photoimmobilized systems can to be explained by two factors: the first one is the presence of metal ions in a partially reduced state and the second deals with the presence of organic ligands in the surface complexes. Due to the first one the reaction can be performed at a quite low temperature while the second favors the high selectivity of the process. Photoimmobilized catalysts have been recently shown to be active and of a steady state action in hydrogenation and ammonolysis reactions.

REFERENCES

1. E.V.Kashuba, L.V.Lyashenko and V.M.Belousov, Kinetikai kataliz. 30 (1989) 474. 2. V.MBelousov, E.V.Kashuba and L.V.Lyashenko, Ukr.khim.zur. 57 (1991) 287 3. L.V.Lyashenko, V.M.Belousov and E.V.Kashuba, Teor. exper.khim. 33 (1997) in press.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1183

Selective oxidation catalysis over heteropoly acid supported on polymer In Kyu Song a, Jong Koog Lee, Gyo Ik Park and Wha Young Lee* aDepartment of Industrial Chemistry, Kangnung National University, Kangnung, Kangwondo 210-702, Korea Department of Chemical Engineering, Seoul National University, Shinlim-dong, Kwanak-ku, Seoul 151-742, Korea* Membrane-like heteropoly acid-blended polymer film catalysts were prepared using a common solvent (or mixed solvents) and they were tested as fixed-bed catalysts for the ethanol conversion reaction in a continuous flow reactor. It was found that heteropoly acid catalyst was finely and uniformly distributed through the polymer matrix. All the film catalysts showed the higher selectivity to acetaldehyde than the bulk solid catalyst. Conversion and selectivity over the film catalysts were also affected by the nature of solvent and polymer. Microporosity of the film catalyst was controlled by the phase separation method. The microporous film catalyst could be regarded as a highly dispersed heteropoly acid catalyst supported on polymer matrix. The film catalysts were characterized by IR, TPD, SEM, EDX, DSC, and ESCA. 1. INTRODUCTION

Heteropoly acids (HPAs) are inorganic acids as well as oxidizing agents [1-3]. They are highly soluble in polar solvents such as water, alcohols, and amines, but some HPAs are insoluble in non-polar chemicals such as benzene and olefins [4-6]. The solubility of HPAs in turn is closely related to their ability to adsorb reactants. Polar substances readily penetrate into the bulk of HPAs to form a pseudo-liquid phase [7], whereas non-polar chemicals are mostly adsorbed on the surface of HPAs [8]. Owing to these characteri.stics, HPAs have been widely investigated and have been used in a commercial process producing methacrylic acid [9-11]. It is well known [12-17] that the acid and redox properties of HPA can be controlled in a systematic way by replacing the protons with metal cations and/or by changing the heteroatom or the framework transition-metal atoms. Novel catalysis of HPA has been also modified by combining HPA with ion exchange resins [18] or conjugated polymers [19]. Another method for the modification of novel catalysis of HPA is to blend HPA with polymer using solvents to form a membrane-like film catalyst [2022]. Dispersity and redox property of HPA can be easily controlled by this method to meet the need for low temperature oxidation reactions. Another advantage of the film catalyst is that porosity can be also controlled by the membrane preparation technique [23] as described in this work.

1184

Membrane-like HPA-blended polymer film catalysts were prepared using organic solvents and they were tested as fixed-bed catalysts for the ethanol conversion reaction in a continuous flow reactor. The effect of solvent and polymer on the catalytic activity of the film catalyst was examined. The porosity of the film catalyst was also controlled by the phase separation method.

2. EXPERIMENTAL 2.1. Preparation of the film catalyst H3PMo12040 (PMo, Aldrich) was purified and calcined at 300oc for the precise quantification. Polysulfone (PSF, Udel-1700 from Union Carbide), polyethersulfone (PES, Victrex 5200P from ICI), and polyphenylene oxide (PPO, poly-2,6-dimethyl1,4-phenylene oxide from Aldrich) were used as blending polymers. Dimethylformamide (DMF), and mixture of methanol (M) and chloroform (C) were used as blending solvents. A homogeneous PMo-polymer-solvent solution was prepared at room temperature. This solution was casted on a glass plate with a constant thickness in ambient condition (56% relative humidity) to form a membrane-like film catalyst. The homogeneous solution was also casted under different relative humidity (RH) in order to control the porosity of the film catalyst by the phase separation method [23]. All the dried film catalysts were thermally treated at 150oc before reaction and characterization. Composition and preparion method of the film catalysts will be described in Section 3 in detail. 2.2. Reaction and characterization Vapor-phase ethanol conversion reaction was carried out in a continuous flow reactor. The film catalyst was cut into small pieces (2 mm x 2 mm) and used as a fixed-bed catalyst. All the film catalysts were treated at 150oc for 1 hr by passing air (5 cc/min) before the reaction. Ethanol was preheated for vaporization and fed to the reactor together with air as both a carrier gas and an oxygen source. Net amount of PMo was 50 mg in each run. The products under steady state condition were analyzed with an on-line GC (HP 5890 II). Conversions and selectivities were calculated on the basis of carbon balance. The film catalysts were characterized by IR (Midac Co. M2000), TPD, SEM (Jeol JMS-35), EDX (Philips PV-9900), DSC (TA Instruments TA200), and ESCA (Perkin-Elmer PHI 581). All the film catalysts were thermally stable during the reaction because the reaction was carried out at temperatures below the glass transition temperature (Tg) of corresponding polymers. Reaction conditions will be described in detail for each run. 3. RESULTS AND DISCUSSION 3.1. Preparation and characterization of PMo-PSF-DMF A homogeneous PMo-PSF-DMF solution was successfully obtained by dissolving both PMo and PSF in a common solvent of DMF. The viscous and greenish PMo(4.76 wt%)-PSF(23.81 wt%)-DMF(71.43 wt%) solution was casted on a glass plate with a constant thickness at 56% RH to form a membrane-like film, and subsequently it was dried for 4-5 hrs at the same condition. The thickness of the prepared PMo-PSF-DMF film catalyst was 0.017 mm.

1185

The film catalysts comprising PMo and each polymer material show two different thermal behaviors. Tg of a polymer increases when it forms a physicochemical blending with PMo, and, on the other hand, Tg of some polymers decreases when it forms a physical blending. Fig. 1 shows the thermal behavior of PMo-PSF-DMF and PSF-DMF (PMo-free film). Tg of PSF-DMF and PMo-PSF-DMF were found to be 187oc and 174oc, respectively. This result means that PMo in PMo-PSF-DMF acts as an impurity and that the blending of PMo with PSF is physical. The physical blending was also confirmed from IR measurements. Typical bands of the Keggin structure of PMo in the PMo-PSF-DMF film was not changed. Fig. 2 shows the crosssectional SEM micrograph of PMo-PSF-DMF film. No visible evidence representing PMo was found in the PMo-PSF-DMF film and there was no distinctive difference between PSF-DMF and PMo-PSF-DMF. This indicates that PMo was not recrystallized into the large particles but was finely distributed as fine particles invisible in the SEM in the PMo-PSF-DMF. The uniform distribution of PMo in the PMo-PSF-DMF film was also confirmed by EDX analysis as shown in Fig. 3.

3o

,T

~ .8,7~

(9

*

0

&

100

I

I

200

Temperature (~

I

,

300

Figure 1. Thermal behavior of (a) PMo-PSF-DMF and (b) PSF-DMF. In order to investigate any interaction between PMo and PSF, the oxidation state of molybdenum in bulk PMo and in PMo-PSF-DMF film was measured by ESCA. The spectrum could be fitted with only one doublet corresponding to Mo 3d3/2 and Mo 3d5/2. The binding energies of Mo 3d3/2 and Mo 3d5/2 in both catalysts were 235.3 eV and 232.1 eV, respectively. It was confirmed that there was only one type of molybdenum (VI)in both PMo and PMo-PSF-DMF. A DMF-TPD experiment on bulk PMo revealed that DMF desorption started at 150oc and reached the maximum at 270oc and 337oc. These temperatures are higher than Tg of PMo-PSF-DMF film. This fact suggests that DMF (organic base) is chemically adsorbed on the acid sites of PMo in the PMo-PSF-DMF film and affects the acidic function of the film catalyst. It is summarized that highly dispersed PMo catalyst supported on PSF was obtained by blending these two materials. Although the oxidation state of molybdenum of PMo in the PMo-PSF-DMF was not changed, the acidic catalytic activity of the film catalyst might be affected by DMF that was adsorbed on the acid sites of PMo.

1186

Figure 2. Cross-sectional SEM micrograph of PMo-PSF-DMF (xl,000).

Figure 3. EDX image of PMo-PSF-DMF film by mapping on only molybdenum.

3.2. Catalytic activity of PMo-PSF-DMF

Table 1 shows the catalytic activity of bulk PMo and PMo-PSF-DMF film catalyst at 170oc. Acetaldehyde is formed by oxidation reaction while ethylene and diethylether are formed by acid-catalyzed reaction over PMo catalyst [24]. As shown in Table 1, PMo-PSF-DMF film catalyst shows the higher ethanol conversion than bulk PMo. The PMo-PSF-DMF shows remarkably enhanced yield and selectivity for acetaldehyde, but it shows obviously decreased yield and selectivity for ethylene and diethylether compared to the bulk PMo. The oxidation activity of the film catalyst was about 10 times higher than that of bulk PMo. It was believed that the enhanced oxidation activity of the film catalyst was due to the fine distribution of PMo through PSF matrix whereas the reduction of an acidic activity was due to DMF which was strongly adsorbed on the acid sites of PMo. The DMF effect was also confirmed by the catalytic activity of PMo-DMF in Table 1. PMo-DMF showed the suppressed acidic catalytic activity compared to the bulk PMo. Above results imply that the film

1187

catalyst can be applied to the low temperature oxidation reactions to obtain high yield and selectivity for oxidation product by enhancing PMo dispersion and by suppressing the acid-catalyzed reaction. Table 1 Catalvtic activitv of PMo-PSF-DMF film catalvst at 170oc Catalyst Amount of EtOH converted to each product Ethanol (xl05 mole/g-PMo-hr) & (% selectivity) conversion(%) CH~CHO C;~H4 C2H5OC2H5 Bulk PMo 2.60(16.4) 2.41 (15.2) 12.00(68.4) 2.8 PMo-DMF a) 4.16(64.7) 0.82(12.7) 1.50(22.6) 1.0 pMo-PSF-DMF b) 26.70(80.9) 9.87(3.0) 5.30(16.1) 5.3 W/F=32.43 g-PMo-hr/EtOH-mole, air=5 cc/min, film thickness=0.017 mm, a)PMo recrystallized from DMF, b)PMo(4.76 wt%)-PSF(23.81 wt%)-DMF(71.43 wt%) solution was casted and dried at 56% RH

3.3. Preparation and characterization of PMo-polymer-MC A new method for the preparation of PMo-imbedded polymer film catalyst using mixed solvents was successfully developed toward the modification of novel catalysis of HPA. HPA and polymer can be easily blended if both materials are soluble in a common solvent as in the case of PMo-PSF-DMF. Though HPA and polymer are not soluble in the same solvent, if a solvent dissolving HPA and another solvent dissolving polymer are miscible, HPA and polymer can be blended using the solvent mixture. Methanol was used for PMo and chloroform was used for polymer. A homogeneous PMo(1.22 wt%)-polymer (6.90 wt%)-methanol(4.41 wt%)chloroform(87.47 wt%)solution was casted on a glass plate at 56% RH, and subsequently it was dried for 4-5 hrs at the same condition. The thickness of the prepared film catalyst was 0.017 mm. PSF, PES and PPO were used as blending polymers. PMo-free polymer films were also prepared at the same condition for comparison. The PMo-blended polymer film catalyst was denoted as follow ; for example, PMo-blended PSF film catalyst prepared using methanol (M)-chloroform (C) mixture was denoted as PMo-PSF-MC. Fig. 4 shows the SEM micrograph of bulk PMo and PMo-PPO-MC film catalyst. The bulk PMo was large cluster with diameters of 10-100 llm. All the film catalysts retained greenish color implying the fine distribution of PMo through each polymer matrix. No visible evidence for PMo in PMo-PSF-MC and in PMo-PES-MC was found in the SEM images as in the case of PMo-PSF-DMF. This indicates that PMo is uniformly and finely distributed through PSF and PES matrix. On the other hand, PMo in PMo-PPO-MC exists as agglomerates having diameters of 1 I~m or less, although PMo dispersion is much improved compared to the bulk PMo. Thermal behavior of the films and film catalysts are shown in Fig. 5. Tg of PESMC was 236oc whereas that of PMo-PES-MC was 219oc. Tg of PMo-PSF-MC was not detected from a room temperature to 350oc. However, considering that the physical state of PMo-PSF-MC was changed after the reaction over 170oc and became fragile, Tg of PSF was supposed to be lowered after blending with PMo. The decreased Tg of PSF and PES after blending with PMo means that there is n o

1188

interaction or no chemical bonding between PMo and polymer in the film catalyst, and PMo acts as an impurity for each polymer as in the case of PMo-PSF-DMF. The increased Tg of PPO after blending with PMo suggests that there is a certain interaction between PMo and PPO and that PMo is not an impurity for PPO. Chemical state of PMo-PPO-MC is not clear, but it can be presumed that PMo-PPOMC may show different catalytic activity from PMo-PSF-MC and PMo-PES-MC.

Figure 4. SEM micrograph of (a) bulk PMo (x480) and (b) PMo-PPO-MC (x3,000).

A

,~

(d)

o

m

14.

"

(b).

~

~

236~

185oC

m

r ..r

221~

0

!

I

100

I

I

200

I

,Y'-"-

300

Temperature (~ Figure 5. Thermal behavior of (a) PPO-MC, (b) PMo-PPO-MC, (c) PSF-MC, (d) PMo-PSF-MC, (e) PES-MC, and (f) PMo-PES-MC.

1189

3.4. Catalytic activity of PMo-polymer-MC

Ethanol conversion and product selectivity over the film catalysts are listed in Table 2. All the film catalysts show the higher ethanol conversion than bulk PMo. The enhanced conversion over the film catalyst is believed to be due to enhanced surface area of PMo. The conversion is in the follwoing order; PMo-PSF-MC > PMoPES-MC > PMo-PPO-MC > PMo. PMo-PPO-MC shows the smallest conversion among three film catalysts as expected in SEM images of Fig. 4. This may be partly resulted from partial agglomeration of PMo throughout PPO. Bulk PMo and PMo-MC show similar conversion and selectivity. This means that the mixed solvent has no influence on the catalytic activity of PMo unlike DMF of the PMo-PSF-DMF. It also suggests that the main reason for the enhanced activity of the film catalyst is not the effect of mixed solvent but the enhanced PMo dispersion upon blending. The enhancement of ethanol conversion over the film catalyst contributes to the increase of acetaldehyde yield via oxidation reaction over highly dispersed PMo throughout the polymer support. The selectivity to the oxidation reaction over PMo-PPO-MC is three times or more to other two film catalysts and that to dehydration reaction is 50% or less. Lower surface area of PMo-PPO-MC may be responsible for the lower activity but different selecivity can not be explained by the different surface area. The increased Tg of PPO after blending with PMo suggests that there is some interaction (like chemical bonding) between PMo and PPO. The interaction between two materials can contribute to the inhibition of acidic activity of PMo-PPO-MC. Table 2 Catalytic activity over PMo-polymer-MC at 170oc Selectivity(%) Catalyst Conv e rsion (%) CH3CHO C2H4 C2H~OC~H 5 Bulk PMo 6.9 12.8 8.4 78.8 PMo.MC a) 7.4 10.5 8.6 80.9 PMo-PSF-MCb) 39.5 20.0 16.1 63.9 PMo_PES_MC b) 33.7 9.0 22.4 58.6 p Mo_PPO_MC b) 13.4 59.4 9.8 30..8 W/F=169.1 g-PMo-hr/EtOH-mole, air=5 cc/min, film thickness=0.017 mm, a)PMo recrystallized from methanol-chloroform mixture, b)PMo(1.22 wt%)-polymer (6.90 wt%)-methanol(4.41 wt%)-chloroform(87.47 wt%) solution was casted and dried at 56% RH In order to confirm the individual effect of polymer materials on the catalytic activity, the perm-selectivities of O2/ethanol through the film catalysts were measured at 80oc where the extent of reaction was negligible. As shown in Table 3, the ratio of O2/ethanol in permeation side is smaller than that in feed side. This means that the 02 permeability is smaller than the ethanol permeability and that the permeation rate of 02 is the rate-determining step. The lowest ratio of O2/ethanol through PMo-PES-MC may be responsible for the lowest acetaldehyde selectivity over PMo-PES-MC. O2/ethanol ratio in permeation side and acetaldehyde selcetivity showed the same trend in the following order as shown in Table 2 and Table 3 ; PMo-PPO-MC > PMo-PSF-MC > PMo-PES-MC.

1190

Table 3 Permeability ratio of 02/ethanol through the film catalyst at 80oc Catalyst Pressure (atm) , Permeability ratio of O2/ethanol Feed side Permeation side PMo-PSF-MC a) 0.9 1.04 0.57 PMo-PES-MC a) 0.9 1.04 0.11 PMo-PPO-MC a) 1.1 1.04 0.85 film thickness=0.017 mm, permeation area=17.65 cm 2, a)PMo(1.22 wt%)-polymer (6.90 wt%)-methanol(4.41 wt%)-chloroform(87.47 wt%)solution was casted and dried at 56% RH

3.5. Porosity control of PMo-PSF-DMF Another advantage of PMo-PSF-DMF is that porosity of the film catalyst can be controlled by the membrane preparation technique. The homogeneous PMo(4.76 wt %)-PSF(23.81 wt %)-DMF(71.43 wt%) solution was used for the preparation of microporous film catalyst by the phase separation method. Water vapor was used as a non-solvent for PSF. Phase separation rate was controlled by modulating water vapor concentration (RH). RH might affect DMF evaporation rate and phase separation rate. Fig. 6 shows the cross-sectional SEM micrograph of the film catalysts which were prepared at different condition. PMo-PSF-DMF(56V) was prepared by casting the solution at 56 % RH and by drying it in vacuum. PMo-PSF-DMF(85) was prepared by casting and drying the solution at 85 % RH. PMo-PSF-DMF(56V) had no micropores because all water vapor and DMF were evacuated as shown in Fig. 6. PMo-PSF-DMF(56V) was a non-porous film and it showed no measurable microporous properties. On the other hand, PMo-PSF-DMF(85) had well-developed micropores and honey comb type cells with an average pore diameter of 0.25 l~m.

Figure 6. Cross-sectional SEM micrograph of (a) PMo-PSF-DMF(56V) (x4,000) and (b) PMo-PSF-DMF(85)(x2,000).

1191

Its total pore area was about 25 m2/g. The blending of PMo with PSF is a physical one and PMo-PSF-DMF film can be regarded as a finely distributed PMo catalyst supported on PSF as described in Section 3.1. Most of PMo in the PMo-PSFDMF(85) film catalyst is presumed to exist on/near the surface of pore wall as an encapsulated and physisorbed state after the phase separation process. It is proposed that each of micropore and honey comb type cell acts as a micro-reactor having well distributed PMo on/near the wall. 4.

CONCLUSIONS

Membrane-like HPA-blended polymer film catalysts were prepared using organic solvents (a common solvent or mixed solvents) and they were tested as fixed-bed catalysts for the ethanol conversion reaction in a continuous flow reactor. It was found that PMo was finely and uniformly distributed through the polymer matrix. All the film catalysts showed the higher acetaldehyde yield and selectivity than the bulk solid PMo. Conversion and selectivity over the film catalyst were also affected by the nature of solvent and polymer. A microporous film catalyst was also successfully prepared by the phase separation method with the modulation of RH. The film catalyst could be regarded as a highly dispersed heteropoly acid catalyst supported on polymer. It was concluded that the film catalyst could be applied to the low temperature oxidation reactions to obtain high yield and selectivity for oxidation product by enhancing catalyst dispersion and by suppressing the acid-catalyzed reaction. REFERENCES

1. N. Mizuno and M. Misono, Chem. Lett., (1984) 669. 2. M. Ai, Appi. Catal., 71 (1981) 88. 3. I. V. Kozhevnikov, Catal. Rev.-Sci. Eng., 37 (1995) 311. 4. M. Misono, Mater. Chem. Phys., 17 (1987) 103. 5. J. B. Moffat, J. Mol. Catal., 52 (1989) 169. 6. T. Okuhara, T. Nishimura and M. Misono, Chem. Lett., (1995) 155. 7. M. Misono, K. Sakata, Y. Yoneda and W. Y. Lee, in T. Seiyama and K. Tanabe (Eds.), New Horizons in Catalysis, Proc. 7th Int. Cong. Catal., Tokyo, 30 June4 July 1980 (Stud. Surf. Sci. Catal., Vol 7B), Kodansha-Elsevier, 1980, p. 1047. 8. M. Misono, Catal. Rev. -Sci. Eng., 29 (1987) 269. 9. M. Ai, J. Catal., 116 (1989) 23. 10. N. Mizuno, T. Watanabe and M. Misono, Bull. Chem. Soc. Jpn., 64 (1991) 243. 11. H. Mori, N. Mizuno and M. Misono, J. Catal., 131 (1990) 133. 12. H. C. Kim, S. H. Moon and W. Y. Lee, Chem. Lett., (1992) 1987. 13. M. Ai, Appi. Catal., 4 (1982) 245. 14. N. Mizuno and M. Misono, J. Mol. Catal., 86 (1994) 319. 15. S. S. Hong and J. B. Moffat, Appl. Catal., 109 (1994) 117. 16. C. L. Hill and C. M. McCartha, Coor. Chem. Rev., 143 (1995)407. 17. T. Okuhara, N. Mizuno and M. Misono, Adv. Catal., 41 (1996) 113. 18. K. Nomiya, H. Murasaki and M. Miwa, Polyhedrons, 5 (1986) 1031. 19. A. Pron, Synth. Met., 46 (1992) 277. 20. J. K. Lee, I. K. Song, W. Y. Lee and J. J. Kim, J. Mol. Catal., 104 (1996) 311.

1192

21. 22. 23. 24.

I. K. Song, S. K. Shin and W. Y. Lee, J. Catal., 144 (1993) 348. I. K. Song, J. K. Lee and W. Y. Lee, Appl. Catal., 119 (1994) 107. Y. B. Kim, MS Thesis, Seoul National University, Seoul, Korea (1996). T. Okuhara, A. Kasai, N. Hayakawa, Y. Yoneda and M. Misono, J. Catal., 83 (1983) 121.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

A s t u d y of V2Os-KaSO4-SiO~. c a t a l y s t s o x i d a t i o n of t o l u e n e to b e n z a l d e h y d e

1193

for

catalytic

vapor-phase

A. O. Rocha Jr a, A. L. Chagas a, L. S. V. S. Sufi~ ", M. F. S. Lopes a and J. A. F. R. Pereira b aEscola Polit~cnica - Universidade Federal da Bahia - Rua Aristides Novis, 2, 20 andar, Federaqio, 40210-630 - Salvador, Bahia, Brasil bFaculdade de Engenharia Quimica - U N I C A M P - C.P. 6066 - 13081-970 Campinas, Silo Paulo, Brasil The vapor phase catalytic oxidation of toluene to benzaldehyde has been studied over V90~-I~SO4-SiO9 catalysts in an isothermal differential reactor. The experiments were carried out at atmospheric pressure, temperatures from 410oC to 470~ and the modified spatial time (W/FTo) ranging from 0 to 180 g cat/mol toluene/h. The experimental tests showed the best performance for the catalyst obtained by co-precipitation. These results may be due to a crystalline phase identified in the process of catalyst characterization. Reaction kinetics was determined using the Mars and van Krevelen model. 1. INTRODUCTION The gas phase partial oxidation of toluene to benzaldehyde is an industrially important reaction due to the fact that benzaldehyde is a common intermediate in a wide variety of chemical reaction processes. Almost haft of benzaldehyde world production is employed in the synthesis of food additives (mainly flavoring). Although the partial oxidation of aromatic hydrocarbons is widely treated in the literature, the oxidation of toluene to benzaldehyde presents few information, thus requiring systematic studies (1). The catalysts which have presented the most suitable characteristics for this oxidation are the metal oxides and metal oxides mixtures of transition elements of the V and VI groups, and the literature reports information related to the formulation, preparation and evaluation of the catalysts (2 - 6), although very few data have been published related to the reaction kinetics. Gunduz and Akpolat (5) present experimental kinetic data of gas phase oxidation of toluene to benzaldehyde over V205 catalysts. Their results are based on the redox model and are restricted to the temperature of 430oC. Also, it is not found in the literature enough data which allow to analyze the activity and behavior of V20~ catalysts based only on their physical characteristics.

1194 2. E X P E R I M E N T A L

2.1. Preparation of catalysts In order to obtain a better understanding of V205 - K2SO4 - Si02 catalysts behavior, in a wide range of operational conditions, a systematic study of two catalysts was carried out. One of them was prepared using the co-precipitation technique and the other obtained through impregnation on a commercial support. They are denominated respectively VK-COPREC catalyst (coprecipitation) and VK-SUPPORT catalyst (impregnation). The VK-SUPPORT was prepared through wet impregnation of the active phase on 20 g of commercial silica (SHELL $980) using an oxalic acid solution of NHaVO3 (7.8940 g) and K2SO4 (6.5919 g). The process of impregnation was carried out in a rotoevaporator at 80oC for 8 hours. Afterwards the impregnated support was dried in an oven at 95 - 100oC during 4 hours. The VK-COPREC catalyst was obtained through co-precipitation using Ray and Mukherjee (6) methodology. Silica gel was initially prepared from the addition of sulfuric acid to a solution of sodium silicate (22.35 g). The prepared gel was washed and dissolved in a solution of potassium hydroxide (20.55 g) to originate potassium silicate, and to this solution diluted H2SO4 was added. The precipitate, after being dried for 10 hours in an oven, was impregnated with a NH4VO~ solution in a roto-evaporator for 30 hours, and then dried out at 95 - 100oC in an oven.

2.2. Experimental Apparatus All the catalysts were tested in an experimental set-up operated at atmospheric pressure, as shown on the flow diagram (figure 1). A stream of air, after being filtered and dried, was divided in two. One of the streams was saturated by bubbling into toluene, and the other, consisting of pure air, was adjusted and mixed to the first one to obtain a stream with the required reactor inlet concentration. The reaction takes place in a stainless steel differential reactor maintained in an oven with controlled temperature. Samples of the exit gas of the reactor were fed to a gas chromatograph with a 10'x 1/8" packed chromatographic column of 10% OV-101 in Chromosorb W type AW-DMCS 80/100, which permitted an excellent separation between toluene and benzaldehyde. The calibration methodology was similar to the experimental procedures using streams of known composition of toluene or benzaldehyde bypassing the reactor. For the benzaldehyde calibration, nitrogen was used instead of air to avoid any possible oxidation of the benzaldehyde to benzoic acid. Catalyst particles size -35 +48 Tyler mesh were used in all tests. Porosity was measured using a mercury porosimeter. A 0.1356 ~m pore mean diameter was determined. The Satterfield and Sherwood (7) methodology was used to verify that reaction occurs without any diffusional limitation (internal or external). The effective diffusivity was estimated from the porosity measurements and binary ~sion coefficient and pore tortuosity published in the literature, leading to an estimated value of 10 1 for the generalized Thiele Modulus based on the reaction rate. The effectiveness factor was then considered as 1.0.

1195 The external diffusional effects were evaluated from the generalized JD factor, being (Pwb - Pws) _= 10-Satin, where Pwb and Pws are the partial pressure of toluene in the gas phase and at the catalyst surface, respectively. Thus, Pws iS about 0.4% less than Pwb, SO it was assumed that PT~ _--Pws. In all the experiments 0.3 g of catalyst were used, and the spatial time varied in a range of 0 to 180 gcat.h/mol. Toluene concentration in the reactor feed stream was 0.5% molar, and the reactor was operated at temperatures ranging from 410~ to 470~ The catalyst was oxidized and activated within the reactor.

T B - Termostatic Bath C - Compressor G C - Gas Cromatograph C S - Silicagel Column C F - Filter T C - Temperature Controller F O - Oven I R - Integrator/Registrator T I - Temperature Indicator M - Manometer R E - Reactor R1, R2 - Rotameter SB - Saturator filled with benzaidehyde ST - Saturator filled with toluene T1 to T 5 - Thermocouple CS

F

Figure 1 - Experimental apparatus

3. R E S U L T S AND D I S C U S S I O N The experimental results for catalyst evaluation show that the VK-SUPPORT (V205 and I~SO4 in a silica support) presents low activity and is extremely unstable, the activity decaying very rapidly with time. On the other hand, the VK-COPREC (V20~ - K2SO4 - SiO2 obtained through co-precipitation) shows to be active, presenting a selectivity of almost 100% at 450~ Also, this catalyst is very stable for long periods of reactor operation. Figure 2 presents the selectivity of toluene to benzaldehyde as a function of the reactor temperature for W / Fwo = 110 g cat.h/tool, for the VK-COPREC catalyst.

1196 100 90 A

o~

80 70 60

o~

50 40

rl)

30 20

W/FTo= 111 gcat h/mol

10 0 410

I

I

I

I

I

420

430

440

450

460

470

Reaction Temperature (~ Figure 2 - Effect of t e m p e r a t u r e on the selectivity

During the experimental runs it was shown that the above selectivity is dependent on the reduction level of the catalyst. These data agree with the published results of Trimm and Irshad (2). It was found that the selectivity of toluene to benzaldehyde increases with reaction time, while the toluene conversion shows to be practically constant. These results indicate that a given ratio of V +~ / V +4, after the process of oxidation must be attained, so that satisfactory values of toluene to benzaldehyde conversion are obtained.

3.1. M o d e l i n g of the r e a c t i o n The reaction kinetics was analyzed using the Mars and van Krevelen (8) redox model. C 7 H 8 + 0 2 -~ C 7 H 6 0 + H 2 0 (1) The model equation which shows a better fitting to the experimental data presents two parameters kl (the constant rate of toluene oxidation) and k2 (the constant rate of catalyst oxidation). The reaction rate is second order dependent on toluene partial pressure (PT) and first order dependent on oxygen partial pressure (Po), leading to the following equation

rT =

kl k2 P~ PO

(2)

kl P~ + k2 PO The model parameters were adjusted through the Marquardt method (9), for each temperature, and from the values of kl and k2 the Arrhenius equation constants were determined. Their values are A1 - 7,29 x l0 s mol/g h atm 2 A2 - 7,65 x 105 mol/g h atm

E.1 - 22950 cal/mol Eal - 16021 cal/mol

1197

where A, and A2 are the pre-exponential factors and Eal and Ea2 the values of the activation energy, for the Arrhenius equation. On figure 3 a comparison between the experimental data of the reaction rate and the calculated values from equation (2) is presented.

2,5oE-o3

2,00E-03

r

r

1,5OE-O3

1,OOE-03

r

5,00E-04

9 S...

i

0,00E+00 0,00E+00 5,00E-04 l OOE-03 1,50E-03 2,00E-03 2,50E-03

r e x p (mol/g cat.h)

Figure 3 - Comparison between the experimental data and the calculated values.

3.2. Physical characteristics of catalyst The Electron Scanning Microscope analysis of the VK-SUPPORT and VKCOPREC catalysts show a significant difference between the two used catalysts, although both catalysts present similar textures when they are new. For the VK-SUPPORT catalyst it was found that a substantial part of the impregnated phase had been lost in the used catalyst compared to the freshly impregnated catalyst, as may be clearly seen on figures 4 and 5. The VKCOPREC catalyst shows a very different behavior (figures 6 and 7). While the fresh catalyst shows a similar texture to the VK-SUPPORT, the used VKCOPREC catalyst shows the formation of a crystalline phase, with needle shape crystals, suggesting that this new phase composed of a mixed oxide, is responsible for the high selectivity and stability of this catalyst. The X-ray diffraction tests permit to identify various phases that are present in the two catalysts, both fresh and after use. The X-ray diffraction results on figures 8 to 11 show that the VK-COPREC catalyst presents a better definition of a crystalline structure than the VK-SUPPORT catalyst. The diffraction results show also that the silica, in both catalysts, presents an amorphous structure.

1198

9~ ' : ' : ~ ' : - : : ' ~ : - : : ~ : ' ' . . : + ' . w '.',.," '^ .~.~ -' ', ' ~ .~.[~.~f~. ~.".~'~. "..-,~.'4:/.::":-:.:.74.:-:.:.~:.:~:..-:.:~.'.:.:+:.~.u~,.:.r " .: ~ ,. .~:..~.~.+: +:.:.: :. :.:.:+.::...:.: : :~4:+:~+. :...::.:.~.~.~....q.:.:~;,.-~::

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:-: :. <:..-.:-:..:~.

r ~ .:

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Figure 4- SEI photograph of a fresh impregnated catalyst.

.

. ..: 9

::i ..... :;: ::::ii: :::: ::!:i~::~.>7:.7~: . :;,~:~i:;:!:~:>::i.i:i:~:~!.:. :?~.,.:;:;~.~.::::~-:::~:: ~'~:~

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.

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:

. . . ::: .~...: ::: :.:. ::::.:::::::::::::::::::::::::::::.:.}.::~ "::-.::.

,+

:::::•

..::::::~:::::: :::::::::::::::::::::::..~;::::~

:

Figure 5 - SEI photograph of a used impregnated catalyst.

~i!~i.::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::

i:?:::?

"

:

:: :::::::::::::::::::::::::::::::.:~.

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:

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~iii:i:i!i~!ii!i!!!::~:::~::~~iii:!:?::: ! :::::::ii!:~:

Figure 6 - SEI photograph of a fresh coprecipitate catalyst.

:::~ ::......~i~i!~:~:~!.~

Figure 7- SEI photograph of a used coprecipitate catalyst.

The K2SO~ phase is identified in the VK-SUPPORT catalyst, while in the VKCOPREC catalyst peaks relative to the I~Na(SO4)2 phase are present. This result is coherent with the fact that the silica gel is obtained through co-precipitation from sodium silicate, and even with an efficient washing it is practically impossible to eliminate all the sodium (10). The X-ray diffraction tests also show the presence of the K4V207 phase in the impregnated catalyst. In the coprecipitate catalyst the phases where vanadium is present do not show good definition. It seems that various phases are present both in the form of vanadium oxides as in the form of potassium vanadates. In the used coprecipitate catalyst, as shown on figure 11, a well defined peak, corresponding to a value of 9.26 in 20, is probably due to the sodium vanadate lines. The results show that the high selectivity and stability of the VK-COPREC catalyst may be due to the presence of phases containing sodium vanadates and sulfates.

1199

VK Fresh Impregnated 600

200 f---,~.., '~v,r ~ i

1

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Figure 8 - X-ray diffraction p a t t e r n s of the fresh i m p r e g n a t e d catalyst. VK Used Impregnated

~.~

o,-i

1000 800 t

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400 200

~oo

"~'t":

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Figure 9 - X-ray diffraction p a t t e r n s of the used i m p r e g n a t e d catalyst. VK Fresh Coprecipitate 2000 t~

1500

~ lOOO '~

500 0

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1500

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............................................................................................................................................................................. i

1000

500

2o Figure 11 - X-ray diffraction p a t t e r n s of the used coprecipitate catalyst.

1200 The thermo-gravimetry analysis (TGA), figures 12 and 13, indicate that the VK-SUPPORT catalyst presents a marked peak which corresponds to the loss of water through hydroxyls at temperatures ranging from 160~ to 300~ while for the VK-COPREC a uniform distribution of terminal hydroxyls is detected, indicating a better dispersion of the active phase in the same range of temperature.

105-

-- 2.5

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QO00

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1201 4. CONCLUSIONS The results shown in the present work permit to conclude that the V~OsK2SO4-SiO~ catalyst obtained through coprecipitation shows a good activity and selectivity for the catalytic oxidation of toluene to benzaldehyde, and that the conversion to benzaldehyde depends on the V§ § ratio. The results also show that the proposed redox model fits with accuracy to the experimental data. The results of scanning electron microscopy, the X-ray diffraction and the thermogravimetry (TGA) analysis suggest that the high selectivity and stability of the VK-COPREC catalyst are due to the needle shape crystalline phase formed after the catalyst activation and experimental runs are carried out. This catalyst also shows a better dispersion of the active phase than the supported catalyst. REFERENCES

1. Haber, New Developments in Selective Oxidation by Heterogeneous Catalysis, Amsterdam, Elsevier, 1992. 2. D.L. Trimm and M. Irshad, Journal of Catalysis, 18 (1970) 142. 3. K. Papadatos and K.A. Shelstad, Journal of Catalysis, 28 (1973) 116. 4. A J., Van Hengstum, J.G. Van Ommen, H. Bosch and P.J. Gellings, Applied Catalysis, 8 (1983) 369. 5. G. Giinduz and O. Akpolat, Ind. Eng. Chem. Res., 29 (1990) 45. 6. S.K. Ray and P.N. Mukherjee, Indian Journal of Technology, 21 (1983) 137. 7. Satterfield and T.K. Sherwood, The Role of Diffusion in Catalysis, AddisonWesley Publishing Company, Inc., 1963. Mars and D.W. van Krevelen, Chem. Eng. Sci. Suppl.,3 (1954) 41. 9. Marquardt, Soc. Ind. Appl. Math. J., 11 (1963). 10.A.B. Stiles, Catalyst Supports and Supported Catalysts, Butterworth Publishers, United States of America, 1987. o

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1203

O2/H2 oxidation of hydrocarbons on the catalysts prepared, from Pd(ll) complexes with heteropolytungstates N.I.Kuznetsova, L.I.Kuznetsova, S.V.Koscheev, E.B.Burgina

L.G.Detusheva,

V.A.Likholobov,

M.A.Fedotov,

Boreskov Institute of Catalysis, Novosibirsk 630090, Russia Palladium(II) complexes with heteropolyanions PWllO397"and PW90349-, and originally synthesized bimetallic Pd(II)-Fe(III) complexes with PW90349"were used for the preparation of SiO2 supported catalysts of hydrocarbons oxidation. The composition of the complexes in aqueous solution was characterized by 31p NMR and IR spectroscopy. The supported samples prepared from Pd(II) complexes with the PWllO397"anion exhibited a considerable activity in a liquid-phase oxidation of benzene and cyclohexane with a gas mixture of O2/H2. H2 pretreatment of the catalysts gave rise to increasing the yield of oxigen containing organic products. It was witnessed by XPS and IR studies that heteropolytungstate principally retained its structure and a part of Pd(II) ions became reduced to Pd(0) after supporting and H2 treating the samples at a temperature of 100~ Decomposition of the heteropolytungstate proceeded at a temperature of 450~ resulting in a loss of catalytic properties of the sample. The samples prepared from Pd(lI) complexes with another heteropolytungstate PW90349" anion showed poor catalytic activity in oxidation of hydrocarbons. By contrast, bimetal Pd(II)Fe(III) complexes with the same anion gave active catalysts after supporting and H2 treatment. The specific interaction of palladium species with PWllO397"or Fe(III) in the complexes with heteropolytungstates determine the catalytic properties of the supported samples.

1. Introduction

Hydrocarbon oxidations on platinum metal catalysts readily proceed in a liquid phase in the presence of co-reductant, most simply, dihydrogen. Under the action of co-reductant the active oxygen species are generated from dioxygen. The properties of active oxygen are close to the properties of hydroperoxides, therefore, the two type systems operate in a similar way. With O2/H2 mixture the oxidative transformation of aromatics [1-4], cyclohexene [5-6] and saturated hydrocarbons [7-8] have been studied. It should be noted that some of the systems are effective enough to generate interest of industrial chemistry. The complex composition of the systems using O2/H2 oxidant gives evidence that pure metal catalysts do not operate effectively without promoter, which can be halide ions and/or transition metal oxides. Our idea was to bring into contact platinum metal ions with probable promoter by means of synthesis of metal ion containing heteropolytungstates. As revealed in our previous study, palladium complexes with heteropolytungstate anion PWllO397 do catalyze oxidation of benzene with O2/H2 mixture. Water soluble palladium complexes operate in a two-phase liquid medium. This paper describes preparation, characterization and catalytic performance of SiOa supported palladium complexes with heteropolytungstates.

1204

2. Experimental 2.1. Preparation of Pd(II) complexes with heteropolytungstates in aqueous solutions.. Palladium(II) complexes with the PWI10397" anion, PWllPd and PWllPd2, were prepared as previously described [9] through the interaction of Pd(H20)42§ and NaTPWI iO39 in aqueous solution. Resulting solution contained 0.03 M PWllO397" at a Pd(II) to PWIIO397" ratio of 1 and 2. An alternative procedure was started from Pd(OAc)2 (0.15 mmol in acetone solution) that was mixed with equimolar amount of NaTPWIIO39 in a water. Acetone was evaporated with air flow at room temperature and an aqueous solution was added by dilute HC104 to pH 2 and PWllO397" concentration of 0.03 M. The resulting compound was denoted as PWlIPd(OAe). Palladium(II) complexes with the PW90349 anion, (PWg)2Pd3 and PWgPd3, were prepared from Pd(H20)42+ in sulfuric acid solution, that was added by Na2CO3 concentrated solution to pH 2, and a solid salt NasHPW903424H20 [ 10] was immediately introduced under stirring. A ratio of reagents Pd(II) to PW90349" was 1.5 or 3. After 1-2 minutes dissolving was completed and pH was readjusted to 3.7. Resulting solution contained 0.02 M PW90349. Palladium(II)-iron(III) complexes with the PW90349- anion, (PWg)2PdFe2, (PWg)2Pd2Fe and PWgPd2Fe, were prepared through mixing stoichiometric amounts of Fe2(SO4)3 in a water with Pd(H20)42+ in sulfuric acid. The solution was added by Na2CO3 to pH 2 and a solid salt NasHPW903424H20 was immediately introduced. A ratio of reagents Pd(II) to PW90349" was 0.5, 1 or 2. After 1-2 minutes dissolving was completed and pH was readjusted to 3.7. Resulting solution contained 0.02 M PW90349". 3~p NMR characteristics of the aqueous solutions prepared are given in Table 1. Complexes with heteropolytungstate anions were precipitated through adding 10-fold excess of (C4H9)4NNO3 or CsNO3. From elemental analysis of appropriate tetrabutylammonium salts, the atomic ratio of P:W:Pd:Fe were calculated. Table 1.31p NMR characteristics of complexes in aqueous solution (chemical shift 8 and integral intensity of the peak referred to the total 31p (in brackets)) and composition of the corresponding tetrabutylammonium salts. Complex

8, p.p.m. (% referred to E 31]~))

P:W:Pd:Fe

PW1 iPd PWIIPd2 PWllPd (OAc) (PW9)zPd3 PW9Pd3 (P W9)2Pd2Fe (PW9)2PdFe2 PW9PdzFe

-12.8(55); -12.8(60); -12.8(55); - 12.5(90) - 12.5(75) -45.8(30); -37.8(82); -37.0(12);

1.0:11.3:1.1 1.0:10.6:1.9 no data 2.0:18.4:2.3 1.0:10:2.7 1.0:10.1:0.8:0.5 1.0:8.1:0.4:0.9 1.0:8.0:1.3:0.9

-13.2(45) -13.2(40) -13.2(45)

-66.2(55) -66.0(12) -45.5(18); -63.9(38); -78.8(22)

1205

2.2. Synthesis of Si02 supported samples. Solid samples were prepared through the impregnation of SiO2 (powder with surface area of 200 m 2 g-l) with aqueous solutions of Pd(II) or Pd(II)-Fe(III) complexes with heteropolytungstate anions. Moisture was evaporated to dryness with air flow at room temperature; whereupon the samples were dried at 100~ for 1 hour. Reference samples Pd(OAc)2/SiO2 and Pd(OAc)z/Na7PWllO39/SiO2 were prepared by means of standard impregnation procedure. Acetone solution of palladium acetate and aqueous solution of heteropolytungstate salt were used. After evaporation of the solvent at room temperature, the samples were dried at 100~ The second sample was prepared by means of consequent impregnating and drying procedures. When used, the calcination was performed in air at 400 ~ or 450~ -. Hydrogen treatment of the samples was made in a gas flow (Hz:N2=l:l) at a temperature from room to 200~ for 1 hour.

2.3. Characterization. 31p NMR spectra were registered on a Bruker MSL-400 spectrometer at room temperature. The resonance frequency of 161.98 MHz and accumulation frequency of 0.05 Hz were used, chemical shifts (8) were measured with respect to an external aqueous solution of 85% HaPO4. 0.1 M solution of Na7PWllOa9 has been used to estimate the concentration of heteropolytungstate. IR spectra were recorded in KBr on Specord IR-75 (TBA-salts) or on BOMEM spectrometer (supported samples). CO adsorption measurements were performed at ambient temperature in a pulls mode with hydrogen as a carrier gas. The samples were pretreated in hydrogen flow at 100~ X-ray photoelectron spectroscopy (XPS) data were registered on ES-2401 instrument using low power (100 wt) Mg I ~ radiation source. The spectra were obtained at room temperature, the binding energy was corrected against one of Si 2p (103.5 eV) as a standard.

2. 4. Catalysts testing. Catalytic reactions were carried out in a glass thermostated flask. To observe gas consumption, the flask was connected with a gas burette. The flask was loaded with a portion of catalyst, solvent (CH3CN) and substrate (C6H6 or C6H12), the system was blown with a gas mixture of O2/H2 and sealed. Suspension was subjected to intensive stirring. After 1 hour the reaction was stopped and the organic products were analyzed by GC (Tswet-500 instrument with FID and a column of 0.4% 1nitroanilineantraquinone on carbon black) and GC-MS (LKB model 2091). 3. Results.

3.1. Characteristics of Pd(II) and Pd(II)-Fe(III) complexes with heteropolytungstates. The composition of aqueous solutions denoted as PWI1Pd and PWllPd2 (Table 1) was recently studied [9]. When the molar amounts of Pd(II) and PWI10397" being equal, the solution contained two type complexes of identical composition: PWllPd and PWllPd-OPdWzlP. In the case of PWI~Pd2 solution, overstoichiometric palladium formed bi- and

1206 polynuclear fragments bonded to heteropolytungstate. Both solutions included appreciable amounts of Na2SO4. The compound PWllPd(OAc) was obtained in a solution free of sulfate ions. From alp NMR data the complex was very similar to PWI1Pd, and heteropolytungstate anions were involved into binding with Pd(II). Heterpolytungstate PW90349" anion possesses three vacancies in the structure, being capable to form compounds of the type (~-1,2,3-PW9V30406 [ 11]. Two-charged M ions tended to give complexes of a composition (PW9Oa4)2M3, where M may be Pd(II) ions [12]. To follow this direction, in this study two complexes of the different average composition were prepared under variation of a ratio of reagents: Pd(H20)42+ and NasHPW9034. Both solutions exhibited one of the same principal signal in alp NMR 6=-12.5 ppm, that was different from 6--12.8 ppm for PWIIPd. IR spectra of both PW9Pd3 and (PW9)EPd3 TBA-salts comprised the typical frequency bands of Keggin-type metal-substituted heterpolytungstates (Fig.l): Vas P04 about 1070 cm "1, Vas W--O 950 cm "1, Vas W-O-W(M) about 880-700 cm -] [13]. The distinction of the composition of PW9Pd3 (or 1 (PW9)2Pd3) from PWI]Pd TBA-salts suggested stability of the parent PW9 fragment. When Fe(III) being added concurrently with Pd(II), the formation of bimetal Pd(II)Fe(III) complexes with PW90349 was observed. The bimetal complexes provided several broad 31p NMR peaks (Table 1) distinguished from the signals of both Pd(II) and Fe(III) complexes with PW90349.(Fe(II) complex with PW90349" prepared according to the known procedure [13] and oxidized with dioxygen to Fe(III), gave alp NMR signal at 6=286 ppm) Broadening and considerable negative chemical shift of 31p signals reflected cp~ the specific interaction of heteropolyanion with paramagnetic Fe(III) ions. No changes of 1200 1000 800 600 principal IR bands of TBA-salt were produced Wavenumber, cm-I by Fe(III) ions (Fig.l). Three solutions were m

i

i

i

i

- - - i

....

~ -

i

i

Fig.1. IR spectra of TBA-salts of the obtained which contained Pd(II) and Fe(III) PWqPd3 (1) and PWqPdEFe (2) complexes, ions totally involved into bimetal complex with PW90349 anions. The average composition of the complexes is given in Table 1.

3.2. Characteristics of supported samples. Because of a strong absorption of SiO2 and Na2SO4 in the IR region around 1100 cm "1, the characteristic frequencies of SiO2 supported heteropolytungstates in a region of 1000-700 cm "1 were considered. The principal IR bands of the supported and dried at 100~ samples were identical to those of the solid Cs-salt of PWI]Pd (compare spectra 1 and 2 in Fig.2). Negligible transformation of the spectrum occurred after the low temperature hydrogen

1207 treatment of P W I I P d / S i O 2 that indicated the parent heteropolytungstate anion to keep structure after low temperature reduction of the sample. The hydrogen treatment at the temperature of 100~ gave rise to appearing a small shoulder with v 978 c m "1 (spectruna 3) that could be explained by the reversible transformation of the heteropolytungstate anion [ 14], and was developed with increasing the temperature of treatment. Calcination of the samples in air brought about the transformation of the parent PWllO397 anion to PW120403 at 400~ (v 984 cm -1 on spectrum 4) and the complete destruction of heteropolytungstate to WO3 at 450~ (broad band around 850 c m -1 o n spectrum 5). Issuing from IR spectra, the conditions of stability of the supported heteropolytungstate were consistent with the data of XPS. As shown in Fig.3, spectra 1-2 for W 4f, the electronic state of tungsten corresponding to W 6+ [ 15], was kept after the hydrogen treatment. However, sintering of the sample prior to reduction resulted in some broadening W 6+ signal without distinct splitting the 4f7/2and 4t"5/2 lines (compare spectra 2 and 3), that indicated the destruction of the regular structure of anion. XPS spectra gave also interesting information on the palladium species on the surface. The peak at E=336.5 eV in the spectrum of the sample before reduction (1 for Pd 3d in Fig.3) was attributed to Pd 2+ [16]. The appearance of the Pd ~ species after reduction was evidenced by shifting this peak to 334.8 eV and by admixing the signal of reduced Pd species at about 340 eV (spectrum 2). The similar spectrum belonged to the F., ~ -7. sample passed sintering prior to the hydrogen treatment (spectrum 3). The clearly appeared in spectrum 3 band at 331 eV was ascribed to palladium hydride. Distinctive features of spectrum 3 was more symmetrical shape of the peaks at 334.8 eV and around 340 eV that is explained by lower contribution of oxidized state. Thus, the hydrogen treatment of the supported samples at 100~ resulted in reduction of a part of Pd 2+ to Pd ~which proceeded more completely after sintering. A part of accessible palladium species on the surface of the Pd containing supported samples was estimated by means of CO adsorption measurements; the data are depicted in Fig.4. ~ooo 9b0 a6o 700 Wavenumber, cm-1 So high dispersity of palladium (if the conventional stoichiometry of absorption of CO:Pd=I is accepted) is not surprising for the reference samples Fig.2. IR spectra of Cs+-salt of prepared from acetone solution, but is unusual for P W l l P d (1); untreated PW11Pd(OAc)/SiO2, because a water solution was used in PWllPd/SiO2 (2); PWllPd/SiO2 this case. This is, probably, resulted from the interaction after H2 treatment at 100~ (3) of Pd(II) with the heteropolytungstate in a parent and after sintering at 400~ (4), solution. All the other samples containing NaESO4 w e r e 450~ (5). t... o rs~

1208 characterized by smaller fraction of accessible P d , probably, because of blocking effect of inert salt. W 4f 60

|

,

,

,

,

i

,

,

,

,

i

,

,

Pd 3d

,

,

|

,

,

,

36.!

334.8

/'

3"11

/

50

40

,

,

i

,

,

,

,

336.5

67.5

65.0

t

3 ~/f~r. 30

i

3,,.3

\

\

3 !

20

62.5

2 10

.

.

.

.

,

.

.

35

3Heprx~

.

.

!

~

.

.

.

.

|

9

45

cBa3H, 3B

,

320

330 3Hepm~[

340 CBII3H,

350 3B

Fig.3. XPS spectra for W 4fand Pd 3d of PWllPd/SiO2" untreated (1), H2 treated at 100~ (2), sintered at 450~ and H2 treated at 100~ (3).

3.3. Catalysis. The catalytic properties of the supported samples were tested in oxidation of cyclohexane and benzene with a mixture of O2/H2 gases at a temperature of 20-40~ . Cyclohexanol and cyclohexanone were obtained from cyclohexane and phenol with admixtures of cyclohexanol and hydroquinone (no more than 2% mol. of each) was obtained from benzene. A suspension of unreduced samples in CH3CN consumed the O2/H2 mixture. Upon reaction with O2/H2 gases and hydrocarbon substrates the samples were subjected to reduction reflected by a change of a color from yellow to gray. The yield of organic products is depicted in Fig.4 (upper diagram). The samples containing Pd(II) complexes with PWllO397" showed substantial catalytic activity, whereupon, PWIIPd/SiO2 producing higher yield of organic products than PW11Pd(OAc)/SiO2. The PWIIPd/SiO2 catalyst became inactive after sintering at 450~ All the other samples were inert in unreduced state, except for small activity of one of bimetal Pd(II)-Fe(III) sample.

1209 A

O

E

ElO

II

0

phenol

2 2

3

T-4

5

6

7

8

9

10

11

3

4

5

6

7

8

9

10

11

a_

~

a_

~

~

..~

rt

~

20

O

E E

15

o

tz.~~ l O

o Ix. 5

% Pd COads/Pd

1

2

ta

o_

0

~

EL

131_

n

I::L

EL

a_

0.54 0.31

~

o

=

"~ EL

1.08 0.54 0.54 0.60 0.54 0.57 0.60 0.38 0.19 0.8 0.23 0.80 0.14 0.78 0.75 0.25 0.27

Fig.4. The yield of the products of hydrocarbons oxidation in the presence of SiO2 supported Pd(II) complexes with heteropolytungstates: (A) before H2 treatment, (B) after H2 treatment at 100~ (CO adsorption was measured after the treatment). Conditions: 0.03g of catalyst suspended in lml of CH3CN, 0.1ml of benzene or cyclohexane, O2/H2=1/2,1 hour. Preliminary gas-phase hydrogen treatment of the samples gave rise to increasing the consumption of gases and, in most cases, the yield of the organic products (Fig.4, down diagram). Reduction of the PWIlPd/SiO2 and PWI1Pd2/SiO2 samples improved their activity in oxidation. Increasing the temperature of reduction from room to 100~ resulted in increasing the yield of organic products, but further increase of the temperature had an adverse effect on the yield (Fig.5). On the contrary, activity of the PW11Pd(OAc)/SiO2 catalyst was lower after than before gas-phase reduction, and lower than activity of the analogous PWllPd/SiO3 sample, even though the former one was characterized by more accessible Pd surface. Reference samples, which did not contained palladium complexes with heteropolytungstate, remained inactive in oxidation even after reduction. Analogous behavior was exhibited by the PWllPd/SiO2 sample sintered at 450~ it has only negligible activity in reduced state. These date suggest the promotion effect of the PWllO397anion, the presence of which is necessary to create the activity of palladium species in oxidation. The specific structure of the heteropolytungstate appeared to be a crucial factor since destruction of the anion resulted

1210 in a considerable loss of activity in oxidation. On the other hand, the reference Pd(OAc)2sample did not catalyze oxidation even though it contained the necessary promoter. The reason is the method of preparation providing no chemical bonding between Pd(II) and heteropolytungstate.

NaPWllO39/SiO2

9 phenol 9 cyclohexanol

,.. 16 14 12

~

" ~

r,

1

_

_ _

anone

22 20

.,

10 8

0

-1

50

1~

150

2~

~c

16 14

20

25

30

35

40

T~

Fig.5. The yield of the products v s temperature of H2 pretreatment of the PWllPd/SiO2 catalyst. Fig.6. The yield of organic products v s temperature of reaction in the presence of P W 1 1 P d / S i O 2 , HE pretreated at 100~ Conditions: catalyst 0.03g, CH3CN lml, benzene or cyclohexane 0.1ml, O2/H2=1/2, 1 hour. Another complex with the PWI10397" anion containing acetate anion was available in water solution but appeared unstable in supported state. Destruction upon drying and the following reduction at 100~ made the supported complex less active catalyst than analogous PWIIPd/SiO2 sample. Hydrogen treatment of the samples both in a gas phase and in liquid resulted in reduction of a part of Pd(II) to Pd (0), as detected by XPS. After reduction Pd(0) species retained location near by the heteropolytungstate anion, which can participate the catalytic process in this position. As seen from Fig.6, the most favorable temperature of oxidation catalyzed by reduced PWI1Pd/SiO2 sample, was 30~ At the temperature of 30~ the PWllPd/SiO2 catalyst produced 9 mol/g-at Pd h of phenol and 23 mol/g-at Pd h of cyclohexanol plus cyclohexanone. About 20% of a O2/H2 mixture consumed was spent to oxidation of organics, and the rest formed a water. Two solutions (PW9)EPd3 and PW9Pd3 was shown by 31p NMR to contain one of the same complex. In the second solution the overstoichiometric Pd(II) ions can be connected to heteropolyanion in the form of bi- and polynuclear fragments. Both (PW9)EPd3/SiO2 and PW9Pd3/SiO2 samples possessed low catalytic activity in oxidation even in reduced state, that indicated insufficient promoting ability of the PW90349" heteropolyanion. On the contrary, two of three samples containing bimetal Pd(II)-Fe(III) complexes with the PW90349" anion developed appreciable activity after reduction. The Fe(III) ions rather than heteropolytungstate are responsible for the promotion effect in this case. Bimetal complexes with heteropolytungstate included Pd(II) and Fe(III) ions in different combinations, some of them were advantageous to create the catalytic activity of palladium. The detailed structure of

1211 active in catalysis Pd(II)-Fe(III) complexes is supposed to be the subject of the following study.

4. Conclusion. The results of this study have shown that prepared from Pd(II) complexes with the PWllO397 anion, SiO2 supported samples behave in catalysis of hydrocarbons oxidation in a similar way as their water soluble precursors, but exhibited improved performance because of the more advantageous composition of the catalytic system. In the present conditions, the catalytic activity of palladium species is developed only in the contact with a proper promoter. The effective contact of palladium with the second component is provided when complexes with the heteropolytungstates are applied for the synthesis of the catalysts. The method is appeared powerful in both situations: when heteropolytungstate anion is a promoter by itself and when it serves as a matrix to assemble together Pd(II) with promoting Fe(III) ions. The present study obviously demonstrates the significance of the promotion effect in oxidation with a O2/H2. There is deficient in direct information on the mechanism of this phenomenon. The role of promoter can be speculatively proposed as (i) facilitating oxygen transfer from activator (palladium species) to organic substrate and/or (ii) stabilizing the proper palladium oxidation state. In the latter case oxidized palladium is suggested to participate dioxygen activation. Acknowledgment The study was supported by Russian Foundation for Basic Research, Grant No. 95-0308754a. References 1. A.ltoh, Y.Kuroda, T.Kitano, G.Zhi-Hu, A.Kunai, K.Sasaki, J. Mol. Catal. 69 (1991 ) 215. 2. Jap. Appl. No5-4935, 1993. 3. T.Kitano, Y.Kuroda, M.Mori, S.Ito, K.Sasaki, J. Amer. Chem. Soc. Perkin Prans. 2 (1993) 981. 4. T.Miyake, M.Hamada, Y.Sasaki, M.Oguri, Appl. Catal. 131 (1995) 33. 5. N.I.Kuznetsova, A.S.Lisitsyn, V.A.Likholobov, React. Kinet. Catal. Lett. 38 (1989) 205. 6. N.I.Kuznetsova, A.S.Lisitsyn, A.I.Boronin, V.A.Likholobov, in" B.Delmon and J.T.Yates (Ed.), Stud. Surf. Sci. Catal. Vol.55 (1990) p.89. 7. US Pat. No5.409.876, 1995. 8. S.B.Kim, K.W.Jun, S.B.Kim, K.W.Lee, Chem. Lett. 7 (1995) 535. 9. N.I.Kuznetsova, L.G.Detusheva, L.I.Kuznetsova, M.A.Fedotov, V.A.Likholobov, J. Mol. Catal. A" Chemical (1996) in press. 10 R.Contant, J.-M.Fruchart, J.-P.Ciabrini, M.Fournier, Inorg. Chem., 16 (1977) 2916. 11. P.J.Domaille, J. Amer. Chem. Soc. 106 (1984) 7677. 12. W.H.Knoth, P.J.Domaille, R.L.Harlow, Inorg. Chem. 25 (1986) 1577. 13. C.Rocchiccioli-Deltcheff, R.Thouvenot, R.Franck, Spectrochimica Acta, A 32 (1976) 587. 14. R.I.Maksimovskaya, Kinet. i kataliz (Rus.), 36 (1995) 910. 15. Handbook of x-Ray Photoelectron Spectroscopy. Ed: C.D.Wagner, W.M.Riggs, L.E.Davis et.al. Perkin-Elmer Corporation, Physical Electronics Division, Eden Prairie, Minnesota, 1979. 16. Analysis of surface by Auger and X-ray photoelectron spectroscopy. D.Brigs, P.M.Sikh, Ed., Moskow, Mir, 1987, p.598.

This Page Intentionally Left Blank

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1213

O x i d a t i o n of n - p e n t a n e to phthalic or maleic a n h y d r i d e : The role of t h e VPO c a t a l y s t s t r u c t u r a l disorder Z. Sobalikl, S. Gonzalez 2, P. Ruiz3 and. B. Delmon3 1 j. Heyrovsky Institute of Physical Chemistry, Academy of Sciences of the Czech Republic, Dolejskova 3, 182 23 Prague 8, Czech Republic 2 On leave from Instituto de Quimica. Universidad de Salamanca, Salamanca, Espana. 3 Unitd de catalyse et chimie des matdriaux divisds, Universitd catholique de Louvain, Place Croix du Sud, 2 bte 17, B-1348, Louvain-la-Neuve, Belgium.

ABSTRACT Vanadium phosphate catalysts are obtained from precursors prepared by a two-step synthesis. In the first step, a VOPO4-mixed isobutanol-water intercalate was obtained by precipitation from a solution containing vanadyl isobutoxide, H3P04 and carefi~y adjusted water content (precursor A). In the second step, precursor B was formed by reflux of precursor A in (i) an inert (n-octane) or (ii) reductive (isobutanol) medium. By such a procedure, precursors and catalysts (with PN atomic ratio equal to 1.05) displaying widely different structural defects (XRD, IR) were prepared. Catalysts were tested in the oxidation of n-pentane into maleic (AM) and phthalic (PA) anhydrides. Formation of PA demands a highly ordered structure, while AM could be formed on a highly defective VPO catalyst. INTRODUffHON The oxidation of n-pentane to phthalic anhydride over vanadium phosphate catalyst (VPO) is an interesting reaction that triggers question on the mechanism of selective oxidation reactions. In addition to this fundamental aspect, the reaction might perhaps be of practical usefulness. It would indeed be industrially attractive to replace o-xylene by n-pentane as a feedstock for phthalic anhydride (PA). The corresponding reaction mechanism necessarily involves a very complex set of active sites arranged according the a special surface topology. Such a topology could ensure the exact timing of the consecutive steps, including abstraction of 10 H-atoms, insertion of 3 oxygen atoms and cyclisation of 2 molecules of n-pentane to form PA, and at the same time providing an easy desorption of the final anhydride. Surface properties of a VPO catalyst and the dynamics of surface adsorbed species on the selectivity to PA or maleic anhydrides (MA) and the practical usefulness of this reaction was evaluated by Centi [1]. He showed that, for

1214 economic reasons, a more efficient use of n-pentane is necessary, namely an increase of selectivity of the process to PA must be achieved for future industrial applications. At the present stage of research, parameters controlling the PA/MA formation ratio have only been partially identified and much progress in this field is necessary for future developments. The objective of this paper is to study the influence of the conditions during VPO preparation on the properties of the final precursor and to correlate the solid state properties of these precursors with the catalytic activity in npentane oxidation and in particular in the selectivity to PA. Our attention focused mainly on the extent of structural defects observed in precursors and catalysts. In order to prepare VPO catalysts with varying extent of structt~ral disorder but with constant V/P ratio, the catalyst precursors were prepared via a VOPO4 mixed alcohol/water intercalate, containing various amounts of water, and further converted into the final precursor in an inert (n-octane) or reductive (isobutanol) medium. These two parameters, i.e. water content in the preparation mixture and the character of the medium during transformation into the VPO catalyst precursor, is shown to play a decisive role in the control of the structural characteristics of the final catalyst. Precursors were characterised by elementary analysis, XRD, FTIR, SBET, porosity and XPS methods. The role of structural disorder on the catalytic activity was evaluated using the following classification of structural defects: (i) long range defects showing destruction of links between vanadyl phosphate layers both in XRD and IR spectra; (ii) short range defects exhibiting distortion inside the vanadyl phosphate sheets shown by broadening of the respective IR bands. EXPERIMENTAL a) Materials: V205 (Janssen Chimica, purity > 99.9%), isobutanol (Aldrich, by Karl Fisher titration the content of water was found to be 0.1%), H3PO4 (Fluka, purity > 99%), benzene (Aldrich, purity > 99%), and n-octane (Aldrich, purity > 99%) were used as obtained for V P O preparation. Water was purified on Milli-Q plus system, SiC (particle size 0.3 ram) provided by Prolabo was used as obtained.

b) Vanadyl isobutoxide (V-ISOB) was prepared by refluxing 40g of V2Os in 500 ml isobutanol with addition of 50 g benzene with water continuously removed in a form of water-benzene azeotrop and separated in a Dean-Stark apparatus. Typically, about 6 ml water were isolated by dissolving about 30 g V205. ARer refluxing for about 40 hours, the rest of the unreacted V205 was isolated by filtration of the mixture by a membrane filter (Schelicker and Schuele 100) and the solvent was evaporated to 50 cc at lower pressure for about 16 hours. The vanadium content found in the colourless liquid products was over 95% of the theoretical value expected for pure vanadyl isobutoxide. c) Catalysts precursors preparation. The simplified presentation of the preparation procedure explaining the notification of the precursors during the preparation steps is shown in Fig. 1.

1215 Table I. Amount of water used in the preparation of Precursors A and their identification. order (1) 0 1 2 3 4 5 6 7 number mole (2) 0.2 0.65 1.1 1.9 2.4 3.1 4.2 5.3

(H20N)

ratio (1) this order number was used to denote the individual A, B and C samples. (2) mole of water added to vanadium ratio in the preparation solution of precursor A.

Precursors

A were prepared by addition of H s P O 4 to vanadyl isobutoxideAsobutanol/water mixture. The P N ratio of the mixture was 1.0. The water content, expressed as HsO/V mole ratio in the mixture, was adjusted in individual preparations to a value between 0.2 and 5.3. The mixture was vigorous stirred for 48 hours. The solid isolated by filtration was dried under vacuum at 50 ~ for 2 hours. The notification of the individual samples prepared and differing in the water/vanadium ratio is shown in Table I. Precursors Bo were prepared by refluxing 5.5 g of a precursor A in 100 ml of n-octane for 20 hours. The solid was isolated by evaporation of the solvent under low pressure and dried under vacuum at 80 ~ for 18 hours, filtered and dried at 80 ~ for 18 hours. Precursors Bi were prepared by refluxing 5.5 g of precursor A in 40 ml of isobutanol for 16 hours, filtered and dried under vacuum at 80 ~ for 18 hours. Catalysts C were prepared by equilibration of a precursor B at the npentane/air mixture up to 400 ~ for 20 hours. The notification of an individual Precursor Bo or Bi and a sample C is related to a parent precursor A (see also Table I ). Characterisation methods. Vanadium and phosphorus contents were determined by an ICP analysis (inductive coupled plasma atomic absorption) a f a r dissolution in 0.1 M nitric acid, carbon by measuring the omount of COs produced by total oxidation using Coulomat 702 Stroelheim. The BET specific surface area was obtained using ASPAP 2000 (Micromeritics) by nitrogen adsorption at- 196 ~ after degassing samples at 125 ~ Powder X-ray diffraction analysis was obtained in a high resolution X-ray diffractometer Siemens D-500 XRD using CuKa radiation. For XRD spectra interpretation data available for VOPO4 hydrates and alkoholates i.e., for VOPO4.2 H20 data identified by McMurdie [2], taking the most intensive peak at about 20 =11.8 ~ and indexed as (001), for VOP04. H20 data presented by Ladwing [3] taking the prominent band at 20 =12 ~ indexed as (001), an with the second most intensive line (002) at 2 o =28 ~ and for VOPO4-aliphatic alcohol intercalates as presented by Benes et al. [4]. The identification of XRD reflections of VOHPO4.1/2 H20 and (VO)2P207 were taken according to Bordes et al [5]. IR spectra were recorded with a Bruker IFS 88 spectrometer using KBr technique. Observed infrared bands were interpreted using assignment presented by Ladwig [3] for VOP04 hydrates and by Busca [6] for VOHP04.1/2 H20 and (VO)2P207. XPS analysis was performed with an SSX-100 model photoelectron spectrometer (FISONS) using monochromatized aluminium anode. The binding energy (BE) values were calculated with respect to C ls (BE of C-CH

1216

fixed at 284.8 eV). The fitting routine was used supposing Gaussian/Lorentzian ratio 85/15, the P/V, C ~ and C.~ ratio using sensitivity factors provided by manufacturer and with SiO~ as an external reference. The BE of V2p~ of 518.1 eV and 517.4 eV were used for V5+ and V4+, respectively [7].

Q

O

.

mixing RT

48 hours

_

VOPO4-water--isobutanol intercalate

A(0) to A(4)

n'~

.~-\

r~ ] ~

i

A[0] to A[7]

~risobutanol eflux

R1

/R

2

~ i aBi solution~

~ i a Bi solid - ~

-, Bi 3) /

R 1 - solid state reduction in inert media (n-octane) R2 - reduction "via solution" in isobutanol media R3 - reduction "via solid" in isobutanol media Figure 1. Scheme of samples preparation.

Catalytic tests were made in a U-tube reactor using 0.2 g of VPO d) catalyst. The reaction stream (total flow 30 ml/min) was : n-pentane 0.6 %

1217 vol., oxygen 20 %vol., helium 5 %vol, balance nitrogen. The inlet and the outlet gas streams were analysed by an on-line gas chromatograph equipped with a TCD detector on TENAX and Porapak Q columnA (Altech Associates, Inc.), running under temperature program from 70 to 240 ~ The outlet of the reactor prior to analysis was kept at about 150 ~ to prevent condensation of products. Conversion of n-pentane is defined as the number of n-pentane converted by the number of n-pentane feed (in % tool.). Selectivities to maleic and phthalic anhydrides were expressed as the fraction of n-pentane converted into AM or PA (in % tool.). The data were measured under conditions where the molar conversion of n-pentane was always equal or lower than 15%. 3. RESULTS 3,L P r e p a r a t i o n of catalyst precursor~ Precursors A were prepared in a form of yellow (A[0]) to dark green (A[7]) solid and identified as VOP04 - mixed isobutanol/water intercalates with the amount of isobutanol per VOP04 molecule varying from 1.6 to 0.05, and displaying the basal spacing decrease from about 14.6 /k to about 7.4 ~, respectively.

Precursors B were prepared by various preparation routes and thus the reduction step of V5+ to V4+ was realised in different reaction conditions: (i) By reduction in an inert media (Precursorsr Bo) the colour of the suspension turned due to V5+ reduction under reflux from yellow to black after 3 hours of reflux, and so the isobutanol intercalated in the Precursor A served as a reduction agent. The IR spectra of samples display a broad unresolved band between ca 800 to 1400 cm -1, with a m a ~ m ~ m at about 1085 cm -1 and several very diffuse low intensive bands at the 400 - 700 cm -1 region, tentatively assigned to a very defective structure of a vanadyl phosphate. This was supported by XRD results showing amorphous product, displaying only a low intensity XRD reflection at the 20 = 28.5 ~ (ii) Substantial differences were observed during preparation in the reductive medium (Precursors Bi) using precursors A low (samples A[0] - A[3]), or high (A[4] -A[7]) content of water and accordingly the ~ m p l e s prepared in the reductive media are further divided in the following subgroups: (ii-a) Precursors Bi[0] - Bi[3]. These precursors were formed in the form of a blue solid by slow precipitation from a dark green solution obtained during reflux. Accordingly these are referred to as precursors Bi "via solution". (H-b) Precursors Bi[4] -Bi[7]. No apparent dissolution of the solid was observed by starting from precursors A with a high water content. Instead, the colour of the dispersed solid turned during the first two hours of reflux from yellowgreen into pale blue and then does not change during the rest of the time of the reflux and these samples are further referred in the text as precursors Bi prepared "via solid". Comparing both types of Bi samples it could be summarised that the via solution prepared samples differ in the IR and XRD spectra (see Figure 2 and 3, respectively) in the following structural features : i) there is suppression of Ill bands assigned to POH (~p POH at 1133 cm-1 and v P-(OH) at 933 cm-1. ii) there is a very broad and low intensity of the group of IR bands belonging to the O-P-O structure (80PO at 413 482, 534 and 550 cm-1);

1218 iii) A missing or badly developed IR band at 686 cm -1 assigned to the coordinated water molecules, accompanied by a parallel shift in the position of the maximum of the 8 H20 band. (iv) they display only one broad band in the XRD spectra with 20 of about 30.4 ~ assigned to (202) plane of VOI-IP04. 1/2 H20. (v) the missing line at about 20 = 15.5 ~ indicate a very defective structure with a high disorder along the (010) plane.

g l]

l-

v V9

1800

t(100

1400

1200

1000

w w e n u n d w ~ om"t

Figure 2.

FTIR spectra of precursors Bi

800

600

400

1219

w

m

m

It

20

s

40

Figure 3 X R D spectra of precursors Bi 3,2. Charaetet4zation of s ~ u e t t u ~ defecfs of the catalyst precursors An attempt to characterise the level of the structural order of the layered structures of the vanadyl phosphates is given considering that these defects correspond generally to two categories, i.e. long range or short range defects. The idea is to distinguish roughly between the disorder effecfing interlayer connections of the layers of the cryst~ lattice Gong range defects) and the local disorder around the vanadium atoms (short range defects), see also e.g. [8] and [9] using similar concept. The general features of these two categories used were as follows: i) long range defects showing destruction of links between vanadyl phosphate layers both in XRD and IR spectra. In this category we consider defects indicated by broadening and eventually missing of the XED bands and destruction of links between vanadyl phosphate layers, as suggested by the missing of bands of Sip P-OH and v P-OH at 1135 and 927 cm -1, respectively, and of ~ I-I20 at 686 cm -1. ii) short range defects exhibiting distortion inside the sheets of the structure and manifested by broadening of the IR bands of the vanadyl phosphate molecule. In this category are considered broad or not distinguished other IR bands of vanadyl phosphate molecules at the 1400 - 400 cm -1 region, suggesting defects inside the V O P 0 4 sheets, and so a distortion of the structural order on a molecular level.

1220

The general characteristics of the s~mples prepared by various preparation routes were characterised as follows: Precursors B~[1]-[4] display high extent of defects, both of long range and short range character. Precursors B~ prepared "via solution" (Bt[1] -[3]) contain mostly long range defects, but tl~e short range order is mostly retained. Precursors B~_ prepared "via solid" (Bi[4]-[7]), displayed both short and long range order. By comparing these structural characteristics with the amount of the organic species present in precursors B, and expressed as the (C/V) value, the general trend between the efficiency of elimination of fragments of the reducing agent during formation of precursor from precursor A, and 'the extent of their structural defects is also suggested.

Catalysts Samples C were identified as (VO)2PsO7 with various extent of defects. Only samples prepared via solid displayed some extent of both the short and long range order. Most significantly the IR bands at 790 and 743 cm-1 associated with links between layers were reliable detected only in ~mples via solid and were missing in samples prepared from precursors B via solution. 3.4. Catalytic activity Results of catalytic activity tests of catalysts C together with some structural characteristics are presented in Table II. Table II. Characterisation of the preparation procedure of the VPO catalysts and their activity, in n-pentane oxidation. structural PrecursorB preparation SSET n-pentane SMA.% SPA.% defects (I) route m2/g conversion % evaluation cs) ,

13o(2) Bo(3) t3o(4) Bi(2)

octane 20.0 10.5 S+L 17.5 none octane 24.4 9.9 S+L 18.7 none octane 12.9 11.1 S+L 19.0 none isobutanol, 14.9 11.7 S 5.6 none via solution Bi(3) isobutanol, 16.9 12.8 S 5.6 none via solution B~(4) isobutanol, 26.3 15.6 0 24 4.8 via solid B~(6) isobutanol, 23.7 12.7 0 31 4.6 v/a sol/d B~(7) isobutanol, 19.0 12.4 0 26 9.6 via solid (I) precursor B used for catalytic test. (2) structural defects evaluation: (S+L) - both short and long range defective structure; S - short range defects, long range ordered structure; O - both short and long range ordered structure.

1221

Clearly two regions are observed: one in which both short and long range order are present, and the region in which no or only short range order is observed. M A was the only product of selective oxidation over catalysts produced from Bo or Bi by via solution preparation route, i.e. starting from precursor A with low content of water. Catalysts obtained by the via solid route produced by n-pentane oxidation both M A and PA, and the P A / M A ratio was the highest for samples with the m a x i m u m of water used in the synthesis of precursor. 4. DISCUSSION

By comparing the selectivity of the different catalysts for P A and M A formation in n-pentane oxidation over V P O catalysts, we concluded that P A formation displayed a different sensitivity to the nature of the structural defects of the catalyst. Accordingly, products of selective oxidation of n-pentane over V P O catalysts could be controlled by varying the conditions of the precursor preparation, and formation of M A and P A is directly related to the structural defects of the prepared V P O catalysts. In our previous report [10], it was shown that the formation of maleic and phthalic anhydrides could be controlled by varying the experimental conditions of the precursor preparation. In that work such effect was produced by changing the time provided for developing the V O P 0 4 mixed water/isobutanol intercalate (Precursor A of the present notation). The preparation of VOHPO4.1/2 H20 via full development of the VOPO4-mixed water/isobutanol intercalate favoured the formation of a precursor which promotes production of PA. Only M A and no P A was observed over the catalyst prepared directly from solution (i.e.to some extent similar to samples Bill] -[3] of the present work). Selectivity was correlated with the structural order of the samples. Nevertheless, due to some specific features of the Precursor A formation under conditions used in the previous work [10], these samples also displayed differences in both the X P S surface atomic P/V ratio and the bulk P N value, thus making less convincing the arg~_~ments for the role of structural order of the V P O catalysts in the selectivityto P A or MA. On the contrary, in the present work, samples show the same and nearly stoichiometric P/V bulk atomic ratio. The procedure of the development of the VOHPO4.1/2 H20 (Precursor B)was then controlled principally by the amount of water added during preparation of precursors. The prepared somples then differed by the amount of intercalated organic material as shown by the changes of the C N values. This results in a distortion of the layered structure of the precursor and consequently in the V P O catalyst (XRD and FTIR). Our results confirm the fact the selectivity in the formation of M A and P A is directly related to the structural characteristic of the prepared V P O catalysts. CONCLUSIONS

Using the evidence of the changes in the M A and PA selectivity and the characterisation of the nature of the structural disorder, we suggest that M A could be formed on a highly defective V P O catalyst,but that both the short and long range structural orders are necessary for P A formation by oxidation of npentane over V P O catalysts. It could be speculated that a high structural order is necessary to create a tridimensional surface topology around the

1222 complex active sites with an optimal distribution of the individual active functions to provide for the concerted process of the PA formation. ACKNOWLEDGMENTS The stay of Mrs S.R.G. C a r r ~ was supported by the Direcci6n General de Investigaci6n Cientffica DGICYT (Programa FPU), Ministerio de Educaci6n y Ciencia de Espafia, which is gratefully acknowledged. The Service de Programmation de la Politique Scientifique (Belgium) is gratefuly acknowledged for its Concerted Action grant, especially for the support of Ing. Z. Soba!ik and Dr. P. Ruiz. The authors also thank Mr. M. Genet for his help in the XPS analysis. R~'~CI~ 1. G. Centi, J. T. Gleaves, G. Golinelli and F. Trifiro, III European Workshop Meeting "New Developments in Selective Oxidation" (P.Ruiz and B. Delmon, Eds), Stud. Surf. Sci. Catal., 72 (1992) 231. 2. H. McMurdie, Powder Diffraction 1 (1986) 98. 3. G. Ladwig, Z. Anorg. Allg. Chem., 338 (1865) 266. 4. L. Benes, J. Votinsky, J. Kalousova, and J. Klikorka, Inorg. Chim. Acta, 114(1986)47. 5. E. Bordes, P. Courtine, and J.W. Johnson, J. Solid State Chem., 55 (1984) 270. 6. G. Busca, F. Cavani, G. Centi, and F. Trifiro, J. Catal., 99, (1986) 400. 7. S. R. Carraz~in, C. Peres, J.P. Bernard, M. Ruwet, P. Ruiz, and B. Delmon, J. Catal., 158 (1996) 452. 8. S. Albonetti, F. Cavani, F. Trifiro, P. Venturoli, G. Calestani, M. Lopez Granados, and J.L.G. Fierro, J. Catal. 160 (1996) 52. 9. L.M. Cornaglia, C. Caspani, and E.A. Lombardo, Appl. Cats]. 74 (1991) 15. 10. Z. Sobalik, S. Gonzalez and P. Ruiz, Preparation of Catalysts VI. Scientific bases for the Preparation of Heterogeneous Catalysts, Stud. Surf. Sci. Catal. 91(1995)727.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1223

E l e c t r o c h e m i c a l O x i d a t i o n of P r o p e n e U s i n g a M e m b r a n e R e a c t o r w i t h Solid E l e c t r o l y t e S. Hamakawa, T. Hayakawa, K Suzuki, M. Shimizu and K. Takehira National Institute of Materials and Chemical Research, 1-1 Higashi, Tsukuba, Ibaraki 305, J a p a n

An electrochemical reactor, using ceria-based solid electrolyte coated with yttria-stabilized zirconia (YSZ), and gold and silver as the anode and cathode, respectively, has been employed for the selective oxidation of propene at 450~ On applying a direct current to this system, acrylaldehyde was formed at the gold anode, and its formation rate increased with increasing direct current. Selectivities to acrylaldehyde, CO and C02, based on converted propene, were 13.4, 25.6 and 61%, respectively, when the YSZI SDC was used as an electrolyte membrane. It is found t h a t selectivity to acrylaldehyde in this study was higher than that (ScHo=8.5%) obtained by using SDC alone as an solid electrolyte. This result suggests t h a t YSZ coating on the ceria-based solid electrolyte leads to inhibit complete oxidation of propene at the surface of cerium oxide. From the dependence of the selectivity to oxidation products on the thickness of YSZ, it is considered t h a t the selective oxidation of propene occurred at the Au-YSZ-gas triple phase boundary by the oxygen species pumped electrochemically through the ceria-based solid electrolyte and the YSZ.

1. I N T R O D U C T I O N In the oxidation chemistry, activation of molecular oxygen can be achieved by its reduction, i.e., 02 -* 0 2 ~ 0 --* 0 2 . The transient intermediate oxygen species, e.g. 0 2 , 0 , would play an important role as an active species for oxidation of hydrocarbons. We have reported that a fuel cell type reactor with an oxide ionic conductor, e.g. yttria-stabilized zirconia (YSZ), for the selective oxidation of hydrocarbons, as shown in Figure 1, can serve as an 'oxygen pump' [1 ], by which the oxygen flux transferred across the YSZ can be controlled by the electric potential externally applied between the anode and cathode [2-4]. Four electron reduction of molecular oxygen, 02, to oxide ion, 02. , at the cathode and oppositely four electron oxidation of oxide ion to molecular oxygen takes place at the anode under oxygen pumping conditions. Therefore, it is likely t h a t the

1224 active oxygen species, e.g. 02, O-, can be produced on the anode surface during oxygen pumping conditions. In addition, an electrochemical membrane reactor system has a lot of advantages to CnHm the hydrocarbon oxidation; (1) the rate of oxidation can be controlled 02 externally by varying the electric current or potential, (2) the CO, H2 contribution of the gas-phase CnHm.2 I "~~electrode reaction by 02 can be reduced, (3) CnHm O 1 ~ (catalyst) the system is safe from the view oxide ion conductor point of possible explosion and (4) a fuel cell reactor for energy Figure 1. Principle of the hydrocarbon generation can be constructed. oxidation using an electrochemical Thus, the application of membrane reactor. electrochemical cell using the YSZ for the oxidation of ethene [5], ethylebenzene [6] and methane [7-9] over metal or metal oxide catalysts as the anode has been studied. Furthermore, we have demonstrated that a ceria-based solid electrolyte (CeO2)0.s(Ln01.5)o.2 (Ln=Sm, Gd, Y) is an attractive candidates for use in an electrochemical reactor for the selective oxidation ofpropene at relatively low temperatures [10]. This is due to the high oxide ionic conductivity at the low temperature of 350~ compared with t h a t of YSZ. However, the complete oxidation of propene to carbon dioxide took place easily at the surface of CeO2, so that the selectivity to acrylaldehyde was lower than that using YSZ. Recently, K. Eguchi and H. Arai et al. reported t h a t a fuel cell with SDC coated with YSZ prepared by the ion plating method attained higher power density than that with YSZ alone at 800~ [11]. This is due to the achievement of constructing from a dense film of YSZ on the SDC to suppress the reduction by the fuel gas. In this paper, we report the investigation of a ceriabased solid electrolyte coated with YSZ to inhibit the complete oxidation of propene at the surface of Ce02 as well as the selectivity to partial oxidation to acrylaldehyde. 2. E X P E R I M E N T A L

The apparatus for partial oxidation of propene was constructed as shown in Figure 2. Gd and Sm doped ceria disk, (CeO2)0.s(Gd01.5)0.2 (GDC) and (CeO2)o.s(Sm01.5)0.2 (SDC) (0.8 mm thickness and 13 mm diameter, Anan-kasei Co.), respectively, were used as an electrolyte membrane. YSZ thin film was prepared on the disk by the spin-coating method with the YSZ slurry. The obtained specimen was sintered at 1000~ for 1 hr in air. The thickness of YSZ was controlled by the rotation rate of spin coating apparatus. A thin film of gold

1225

Au wire

5% C3H6--~~--~

t~[ll~ I

Quartz tube ~ ~ ] 1 ~

I[[[[ Ill I /Aluminatube 1111 ~ ]

. [

Glass seals

~111

::-

u lam 0

Sample (0.8 mm)~---I [i

I

~

---[1!~1--~ Thersnocoupm

1111

o lOO% o 2 - - - ~ - 1

Figure 2. Schematic diagram of the electrochemical membrane reactor. (1.0 ~m thick) was prepared as the anode on the YSZ film by a vacuum vaporization method. Porous silver cathode (2-3 ~m thick) was then prepared by painting silver paste on the other face of the disk, followed by annealing at 550~ for 1 h. The two electrodes (projected area; 0.5 cm 2) were connected withgold wires to an electrical circuit in order to control the oxygen transfer flux from the Ag cathode to the Au anode through the electrolyte. The disk was mounted between two vertical alumina tubes and sealed by low melting point glass. The reactor was placed in the electrical furnace to control the temperature of the electrochemical cell system. A gaseous mixture of propene (5%), nitrogen (5%)and helium (90%) was passed (1.5 l.h 1) over the anode at 450~ for testing the catalytic activity of the various electrolyte systems. Oxygen gas (1.8 l.h 1) was introduced into the cathode compartment. The products in the effluent gas were determined by gas chromatography using a thermal-conductivity detector and nitrogen as an internal standard. Porapak Q (GL. Science), PEG 6000 (GL. Science) and molecular sieve 13X (GL. Science) columns were used for analyzing and estimating the rates of hydrocarbons, oxygenated compounds and inorganic gases, respectively. The oxygen mass balance was estimated from the comparison between the electrochemical oxygen supply and the oxygen consumption calculated from the amounts of oxygenated compounds. The conductivity of the electrolyte membrane was measured by the ac impedance method [12].

1226 3. R E S U L T S AND D I S C U S S I O N 3.1. Characterization of the c e r i a - b a s e d solid electrolyte c o a t e d with YSZ SEM observations of the surface and the cross-section of the Gd doped ceria coated with YSZ (YSZ I GDC) show t h a t the thin YSZ film is porous and the YSZ particles are fairly uniform in size. The thickness of YSZ film was about 20 ~m. The conductivity of this composite membrane was 4 . 3 9 x 1 0 4 S'cm "1 at 500~ under air atmosphere, which was almost same order as t h a t of Sm doped ceria coated with YSZ (YSZ I SDC) (1.51 X 10 .4 S'cml). However, either value was lower than t h a t of GDC alone (4.03 x 10 -3 S.cm -1) or SDC alone (4.66 x 10 .3 S.cm-1). This may be due to the low conductivity of YSZ film. 3.2. Partial oxidation of p r o p e n e Propene oxidation was carried out by using an electrochemical reactor constructed from a Sm doped ceria electrolyte coated with YSZ (YSZ [ SDC) as a membrane. In a blank t e s t where nitrogen gas alone was passed over the Au anode instead of the reaction gas at 450~ it was confirmed t h a t the oxygen pumping was well controlled by the applied current, i.e., the amount of oxygen gas evolved at the anode coincided well the value calculated from the electric current by using Faraday's law. When the propene mixed gas was introduced into the anode space at 450~ the reaction cell gave a stable voltage of 480 mV (EMF) under open circuit condition. Applying the electric current to the reaction cell, propene oxidation took place, and the rate of oxidation increased with ~. 6increasing electric current (i.e., .~ YSZISDC increasing oxygen flux). No -~ evolution of molecular oxygen was ~k 5 observed in this case. Namely, -~ 4cch all oxygen species pumped ~ electrochemically through the o 3YSZI SDC membrane was CO consumed by the propene o 29 J A oxidation over the Au anode. ~ c~4o Acrylaldehyde, carbon monoxide and carbon dioxide were observed ~ o as the main products under the ~ 0 T" '~ ! 0 0.5 1.0 1.5 closed circuit conditions. The O formation rates of oxidation Current / mA products are plotted as a function of the current in Figure 3. The Figure 3. Partial oxidation ofpropene using formation rate of acrylaldehyde the YSZISDC electrochemical membrane increased linearly with increasing reactor. the electric current. This result

1227 indicates t h a t acrylaldehyde formation proceeds by the reaction of propene with the oxygen species pumped electrochemically through the YSZISDC membrane from the cathode to the anode. When the electric current was 1.1 mA, i.e., 10.2 ~mol~r of oxygen flux, the formation rates of acrylaldehyde, carbon monoxide and carbon dioxide were 0.33, 1.9 and 4.53 ~tmol~r, respectively, and the oxygen consumption rate of 9.02 ~tmol/hr was calculated by assuming water as another oxidation product. Mass balance calculated between oxygen evolution and consumption of propene oxidation was about 90%. 3.3. Effect of the YSZ c o a t i n g on the s e l e c t i v i t y to oxidation p r o d u c t s We have studied the partial oxidation of hydrocarbons with electrochemical reactor using an oxide ionic conductor, e.g. YSZ, SDC, etc. [2-4, 10]. In these studies, it was found t h a t a ceria-based solid electrolyte is useful for the propene oxidation to acrylaldehyde at relatively low temperature of 350~ [10]. In this case, however, the acrylaldehyde selectivity was lower than t h a t obtained with the electrochemical reactor constructed from YSZ. This may be due to the high activity ofceria surface for the complete oxidation of hydrocarbons [13,14]. The effect of the YSZ coating on the ceria-based solid electrolyte was shown in Table 1. When the YSZI SDC membrane was used as the solid electrolyte, selectivities to acrylaldehyde (ScHo) carbon monoxide (Sco) and carbon dioxide (Sco2) based on converted propene was 13.4%, 25.6% and 61%, respectively. Here, it should be emphasized that the selectivity to acrylaldehyde increased with YSZ coating compared with t h a t (ScHo =8.5 %) obtained by using SDC alone as a solid electrolyte. In addition, it was found t h a t carbon monoxide formation was observed in the present study, although its formation was not detected in the case of SDC alone. The same phenomena were observed, when the Gd doped Table 1.

Selectivity to oxidation products in the propene oxidation at 450~ Selectivity / %

Solid electrolyte

Thickness of YSZ

CO

C3H40

C02

SDC GDC

--- % --- %

8.5 % 12.0 %

91.5 % 88.0 %

YSZ/SDC

25.6 %

13.4 %

61.0 %

14.6~m

YSZ/GDC

20.7 %

18.2 %

61.1%

20.0~m

YSZ

10%

40.0%

50.0 %

1228

ceria coated with YSZ (YSZIGDC) was used as a electrolyte membrane for the electrochemical oxidation of propene. These results suggest that YSZ coating on the ceria-based solid electrolyte lead to inhibit complete oxidation of propene at the surface of SDC or GDC. However, taking into account the fact t h a t the acrylaldehyde selectivity in the either case of using YSZI SDC or YSZ I GDC was still lower than t h a t obtained with YSZ alone, it is likely that the surface of ceria was not completely covered by YSZ could not be obtained in this study. 3.4. R e a c t i o n site for t h e a c r y l a l d e h y d e p r o d u c t i o n

Generally, in the electrochemical membrane reactor, e.g. the AuIYSZIAg system, molecular oxygen is reduced to oxide ion by four-electron transfer at the Ag cathode, incorporated into and transferred though the YSZ as oxide ion, and then reoxidized to molecular oxygen by four-electron transfer at the Au anode. In addition, it is considered that oxygen evolution at Au anode occurs at the triple phase boundary between Au electrode-YSZ-gas phase [15,16]. Accordingly, we consider t h a t hydrocarbon, e.g. propene, reacts with the oxygen species appeared at the triple phase boundary through the YSZ and forms oxygenated compounds, e.g. acrylaldehyde. Furthermore, it is likely that the oxygen species pumped electrochemically at the triple phase is active for the partial oxidation of hydrocarbons, since the gold has no ability to dissociatively active molecular oxygen [17] and YSZ has no activity for the partial oxidation of hydrocarbons [4]. In order to clarify the reaction site of the partial oxidation of propene using the ceria-based solid electrolyte coated with YSZ as a membrane, we have studied the dependence of the selectivities to oxygenates on the thickness of YSZ. When the Sm doped ceria coated with YSZ (YSZISDC), each selectivity of the oxidation products did not dependent on the thickness of YSZ, as shown in Figure 4.

I 9.~

75 C02

0

O o 9~

50

op,.~

o o

25

4,

9 ~,4

n

l

-

CO C3H40

oP.,l

~

O~ 0

!

10

20

30

40

50

Thickness of YSZ / grn Figure 4. Dependence of the selectivity to oxidation products on the thickness of YSZ.

1229 This result indicates that the main reaction site is not inside the YSZ film, but its surface. Therefore, it is likely that the selective oxidation of propene occurred at the Au-YSZ-gas triple phase boundary by the oxygen species pumped electrochemically through the SDC and then the YSZ, as shown in Figure 5. However, it was found that acrylaldehyde selectivity in the case of using YSZISDC or YSZIGDC was still lower than that obtained with an electrochemical reactor constructed from YSZ alone, as shown in Table 1. Accordingly, a possible occurring of the propene oxidation at the surface of SDC and/or at the triple phase boundary between Au-SDC-gas cannot be denied in this study.

CO

CH2=CHCHO

C3H6

YSZ

Figure 5. Mechanism of the selective oxidation ofpropene using an electrochemical membrane reactor.

4. CONCLUSIONS An electrochemical cell system with ceria-based solid electrolyte coated with YSZ prepared by the spin coating method showed higher selectivity to acrylaldehyde than that with ceria-based solid electrolyte alone. This may be due to the fact that a film of YSZ on the ceria-based solid electrolyte to suppress the complete oxidation of propene. When the YSZISDC disk was used as an electrolyte membrane, selectivity of the oxidation products did not depend on the thickness of YSZ. This indicates that the selective oxidation of propene occurred at the Au-YSZ-gas triple phase bouridary by the oxygen species pumped electrochemically through the ceria-based solid electrolyte and the YSZ.

REFERENCES 1. 2. 3. 4. 5.

E.C. Subbarao, 'Solid Electrolytes and Their Applications,' Plenum Press, New York, 1980. T. Tsunoda, T. Hayakawa, K. Sato, T. Kameyama, K. Fukuda and K. Takehira, J. Chem. Soc., Faraday Trans., 91 (1995) 1111. A.P.E. York, S. Hamakawa, T. Hayakawa, T. Tsunoda, K. Sato, and K. Takehira, J. Chem. Soc., Faraday Trans., 92 (1996) 3579. S. Hamakawa, K. Sato, T. Hayakawa, A. P. E. York, T. Tsunoda, K. Suzuki, M. Shimizu and K. Takehira, J. Electrochem. Soc. 144, (1997) 1. M. Stoukides and C. G. Vayenas, J. Catal., 70 (1981) 137.

1230 6. 7. 8.

9. 10. 11. 12. 13. 14. 15. 16. 17.

J.N. Michaelis and C. G. Vayenas, J. Catal., 85 (1984) 477. D. Eng and M. Stoukides, Catal. Rev. -Sci. Eng., 33 (1991) 375. T. Hayakawa, K. Sato, T. Tsunoda, K. Suzuki, M: Shimizu and K. Takehira, J.Chem. Soc., Chem. Commun., (1994) 1743; K. Sato, J. Nakamura, T. Uchijima, T. Hayakawa, S. Hamakawa, T. Tsunoda and K. Takehira, J. Chem. Soc., Faraday Trans., 91 (1995) 1655. K. Otsuka, S. Yokoyama and A. Morikawa, Chem. Lett., (1985) 319. S. Hamakawa, T. Hayakawa, H. Yasuda, K. Suzuki, M. Shimizu and K. Takehira, J. Electrochem. Soc., 143 (1996) 1264. T. Inoue, T. Setoguchi, K. Eguchi and H. Arai, Solid State Ionics, 35 (1989) 285. T. Yajima, H. Iwahara and H. Uchida, Solid State Ionics, 47 (1991) 117. J.M. Deboy and R. F. Hicks, Ind. Eng. Chem. Res., 27 (1988) 1577. T. Hattori, J. Inoko and Y. Murakami, J. Catal., 42 (1976) 60. H. Yanagida, R. J. Brook and F. A. KrSger, J. Electrochem. Soc., 117 (1970) 593. J. Sasaki, J. Mizusaki, S. Yamauchi and K. Fueki, Bull. Chem. Soc. Jpn., 54 (1981) 1688. G.I. Golodets, 'Heterogeneous Catalytic Reactions Involving Molecular Oxygen,'Elsevier, New York, 1983.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

Vanadium pentoxide catalytic partial oxidation of 1-butene

membrane

1231

reactor

for

Sangjin Moon, Tayoon Kim, Seungdoo Park a, Jihoon J u n g b, Sukin Hong Dept. of Chemical Eng. Korea University, Seoul, Korea SKI petrochemical research center, Suwon-si, Korea b Dept.of Chemical Eng. Kyonggi University, Suwon-si, Korea a

ABSTRACT The vanadium pentoxide cata~tic membrane reactor was prepared by coating its sol inside the Vycor ~ tube membrane. After heat treatment of the prepared membrane, the [010] planes of vanadium pentoxide layer were grown largely which contributes to partial oxidation reaction of l-butene to maleic anhydride. The partial oxidation of l-butene to maleic anhydride was carried out in the catalytic membrane reactor. The maximum selectivity of 95% was obtained at 350~ when the surface velocity was 500cm/h. And at this condition, oxygen permeability was almost four times higher than the reaction had not occured.

1. I N T R O D U C T I O N One of the most industrially important reactions using vanadium pentoxide(V205) catalyst is the partial oxidation of l-butene to maleic anhydride [1]. Partial oxidation reactions are inherently unselective and often make by-products of little or no value. Oxygen-rich feeds result in low product selectivities and high hydrocarbon conversions [2]. Because partial oxidation and total oxidation always proceed competitively, the selectivity of maleic anhydride from l-butene is low. Though fixed bed reactors or fludized bed reactors have been used for partial oxidation for the past 30 years, the selectivity of maleic anhydride has not been obtained higher than 69% [3]. Some attempts have been reported on a new type of reactor to overcome the above limit. This is a membrane reactor which offers some advantages. A membrane reactor plays a

1232 dual role both catalyst and membrane, which is commonly regarded as a barrier capable of being selectively permeated by some components of a mixture or, at least, of changing the composition of a fluid stream that flows through it due to a certain driving force such as a pressure, concentration, or electric potential gradient.[4] First, one of the products is separated by the membrane and chemical equilibrium is shifted. As a result, the conversion is increased. This type of membrane can be used for water-gas shift reaction, dehydrogenation and so on [5]. Second, the formation of by-products is suppressed due to presence of a membrane between the two reactants, so that selectivity is improved. This second type can be applied for the partial oxidation a n d hydrogenation reaction that produces a lot of valuable by-products during the reaction, but lacks in study compared to the first type of membrane reactor. The aim of this research was to prepare the V20~-coated catalytic membrane reactor using sol-gel method, and to investigate reaction charateristics of that through the partial oxidation of 1-butene to maleic anhydride. And also, we investigated the enhancement of oxygen permeability through V2Os-coated catalytic membrane during partial oxidation reaction occurs.

2. EXPERIMENTAL 2.1 Preparation of a V205-coated catalytic m e m b r a n e by sol-gel method V205 sol was made by melting its powder at 900~ and then pouring it into deionized water. After that the mixture was stirred by magnetic stirrer and then it was filtrated. This procedure uses the property that amorphous V205 dissolves easily in water [6]. As a support for the membrane reactor, Vycor | glass with quartz tube connected to either end was used. The length and the diameter of Vycor are 100mm and 7 ram, respectively. In order to coat inside the Vycor with V20~ sol one end of the Vycor was connected to syringe and the other end of it was put into V205 sol. After the sol was sucked in, it was allowed to stay within the Vycor for a certain time to form thin layers on the inside surface of the Vycor, and then the remaining sol was pushed out. The thickness of the thin layers was controlled by the length of time the V205 sol remaining in the Vycor, and by the viscosity of sol. The coated Vycor was dried at room temperature for 24 hours and, in order to crystallize the V205, the heat t r e a t m e n t was carried out by raising the furnace temperature by 1 ~ per minute and keeping it at 450~ for an 0hour. And the V205 membrane was also prepared by dip coating the V205 sol in YSZ disk to characterize the membrane surface and its cross section.

1233

2.2 Reaction of 1-butene and p e r m e a t i o n of oxygen through the membrane The VzO5-coated catalytic membrane connected to the quartz tube was fixed within the 1/2 inch stainless tube by using CAJON ~ ultra-torr fitting. Both nitrogen as the transport gas and 1-butene were passed through the catalytic membrane under the application of oxygen to the outside of the catalytic membrane. The oxygen permeated through the V2Os-coated catalytic membrane and reacted with 1-butene on the coated thin layer. At this time the concentration of the mixture of 1-butene and nitrogen was sustained at 1%, and the surface velocities (volumetric flow rate/surface area of reactor) of 1-butene and nitrogen was set to 500cm/h and 200cm/h, respectively. Maleic anhydride and unreacted compounds were analyzed by on-line gas Chromatography. The permeabilities of oxygen through the membrane were calculated by measuring the decreasing oxygen pressure using a pressure transducer connected to a recorder. To investigate the reaction effects on permeabilities of oxygen, the pure nitrogen was fed into inside the Ve05-coated catalytic membrane, and uncoated Vycor was also used under the same reaction conditions as above. Fig.1. shows the reaction apparatus.

6

1

2

3

.....

1 . 0 2 2. butene 3. N2 4. MFC 5. furnace 6. pressure transducer 7. A/D converter 8. thermocouple 9. membrane

Fig.1. The a p p a r a t u s of catalytic membrane reactor system.

2.3 Characterization Both the specific area and the pore size distribution were measured by adsorbing nitrogen while the existence of macro pores was confirmed by mercury porosimeter. The crystal structure of V205 according to the crystallization temperature and time was investigated by X-ray diffractometer. The morphology of VzO5 sol was observed by TEM and the thickness of the thin layer coated on the Vycor was measured by

1234 SEM. To analyze the unreacted 1-butene and maleic anhydride in membrane reactor, gas chromatography equipped with FID Tenax-GC | column was used. Also to analyze the CO and CO2, chromatography equipped with TCD and molecular sieve 5A Porapak Q column was used.

the and gas and

3. R E S U L T 3.1 S u r f a c e

area

and

pore

diameter

The specific surface area of vanadium pentoxide coated membrane was 1.1 me/g, and the average pore diameter of it, measured by nitrogen adsorption, was 30A. Macro pores identified by mercury intrusion did not exist, and the porosity was approximately 9%. From results of isotherm adsortion-desorption curve,as shown in Fig.2, V2Oscoated membrane has slit-shaped pore structure. Hence the pores mostly exist between the V205 layers, and the surface of those are nearly non-porous.

13. !-" 03 "~

4O

40

35

35 D

25

E := o

20-

~

10

o -o

5

>

30

30

15

0

-

0.0

25

L

20

L O

15

c~

10

l

I i 1 , I , I , 0.2 0.4 0.6 0.8 1.0

~ 0.0

l , l 0.2 0.4

i 0.6

0.8

Relative Pressure (P/P o)

Relative Pressure (PIP o)

(a)

(b)

Fig.2 Adsorption and desorption curve of V 2Os-coated layer with nitrogen (a) before and (b) after calcination, respectively ( [-];adsorption

9

; desorption )

1.0

1235

3.2 X-Ray Diffraction Analysis The X-ray diffraction reasults as a function of the heat treatment of coated membrane are shown Fig.3. The V20~-coated membrane prior to calcination was an amorphous structure without any characteristic peak. When the membrane was heated to 200~ the crystallization was slightly developed, and at 300~ and 400~ the crystallization continued to occur. But at temperature over 400~ the crystal never grew.

[0101

10

20

(a)

30

40

10

20

10

I

40

[010] [1011

20

30 (b)

[010] [2001

[101] [400]

[1011

[4001

I I 30

(c)

[4001

[200]

40

10

20

I

30

40

(d)

Fig.3. The XRD pattern of vanadium pentoxide layer with calcination conditions. (a) before calcination (b) 25"C-200~ (c) 25~ (d) 25"0-450"C

Therefore the heat treatment temperature shoul be higher than 400~ The characteristic peak of vanadium pentoxide occured at 28=15.4 ~ 20.3 ~, 22 ~, 25.6 ~ 26.2 ~ and 31 ~ Each peak shows [200], [010], [110], [210], [101], [400] plane. In the case of the membrane heated up to 200~ [200] plane, [010] plane, [110] plane and [400] plane started to occur but the rate of crystal growth is low. The membrane that was heated upto 300~ shows only the great growth of [010] plane while the [210] plane and [010] plane grew slightly. The membrane which was heated at 450~ showed the great growth in [200] plane, [010] plane, [110] plane and [400] plane, but the crystalization rates on other planes were negligibly small. Of the four crystalline structure, the active sites of oxidation reaction are [010] plane at 20.3 ~ and [101]

1236 plane at 26.2 ~, where [010] plane is the main active site of partial oxidation [13]. Therefore, the higher the rates of intensity of [010] plane and [101] plane are, the more favorably the partial oxidation occurs. The intensity rates are expressed by the morphological factor which is defined by 11OlOl/111o11. In the case of the vanadium pentoxide coated membrane, this factor is ten times higher than those of the ordinary vanadium pentoxide based catalysts. 3.3 P a r t i a l O x i d a t i o n R e a c t i o n The conversion in the membrane reactor was comparatively low. The maximum conversion was only 22% and 33%, respectively, when surface velocity was 500 cm/hr and 200 cm/hr. On the other hand, the selectivity of maleic anhydride in the membrane reactor was very high as shown in Fig.4. When the surface velocity was 200 cm/hr and 500 cm/hr, the maximum selectivity was 80% and 95% at 350~ respectively.

100

100

~

80

:~ 60

~

6o

40

g~ 40

20

2o 00

200

300

400

oo

Temperature ('(3) Fig.4. S e l e c t i v i t y of m a l e i c a n h y d r i d e over V205-coated membrane reactor ( 0 2 pressure at 15 psi) 9 Z~

surface 9 velocity 500 cm/h and 200 cm/h,respectively

200

300

400

Temperature ( ~ ) Fig.5. Selectivity of maleic a n h y d r i d e with 0 2 pressure over V205-coated m e m b r a n e reactor (surface v e l o c i t y 200cm/h) 9B •

" 02 pressure 15, 30 and 45 ,respectively.

The total oxidation occurs when 1-butene reacts to the gas phase oxygen, where the partial oxidation occurs when 1-butene reacts to the lattice oxygen. In the membrane reactor, 1-butene and gas phase oxygen are separated by the membrane, and 1-butene and oxygen in the gas phase cannot react directly, and therefore the total oxidation do not occur. On the other hand 1-butene adsorbed on the V20~ membrane reacted only with the lattice oxygen in the VeO~ membrane, and the desired lattice oxygen was supplied from the other side of V205 membrane. Therefore the partial oxidation occured only at the membrane reactor and the selectivity of maleic anhydride was ncreased.

1237 The conversion in the membrane reactor was very low compared to that in the fixed bed reactor because the amount of V205 catalyst coated on the membrane reactor was only 0.01cc. This restricts the reaction to a small area. The conversion at the surface velocity of 500 cm/hr was lower t h a n that at 200 cm/hr, because the amount of oxygen through V205 membrane wall do not changed with surface velocity of 1-butene. However, the selectivity of the former case was higher in the reverse because the contact time between 1-butene and catalyst is reduced. When contact time is long, a portion of maleic anhydride reacts with the catalyst again, and tends to be converted into CO2 and other by-products and, therefore, lowers the selectivity of maleic anhidride. Because butadiene and acetic acid are produced under 250~ and a portion of lattice oxygen is used for total oxidation above 400~ the selectivities of maleic anhydride under 250~ and above 400~ are relative low. In order to increase the conversion in the membrane reactor, the membrane reactor should be made in hollow fibers, or the unreacted 1-butene should be recycled. 1-butene can be condensed easily because it has a high liquefaction temperature. If the membrane process through recycling is developed, both the selectivity and the yield could be increased. When the oxygen partial pressure applied from the outside was changed to 15, 30 and 45 psi, the effect of oxygen pressure was observed. As shown in Fig.5, the selectivity decreases as the oxygen partial pressure increases. Most notably when the pressure increases from 15 psi to 30 psi the decrease in the selectivity was slight, but when pressure was raised to 45 psi the selectivity showed a substantial drop. This is due to the fact that, the permeation rate of oxygen through membrane was too low as the oxygen pressure was low. The oxygen was supplied only by the adsorption of oxygen on the V20~ membrane. When the pressure was above a certain value, the oxygen was permeated through the membrane and reacted directly with 1-butene, which causes the total oxidation. The permeabilities of oxygen through the Vycor were constant with the pressure difference and decreased with the square root of the temperature (figure not included). So the main mechanism for gas permeation is Knudsen diffusion. However, the permeabilities of oxygen through the V20~-coated catalytic membrane were deviated from Knudsen diffusion. As the partial oxidation proceeded, the permeabilities of oxygen through the V2Os-coated catalytic membrane were increased slightly with temperature. But at the temperature of maximum selectivity obtained, the permeability of oxygen was enhanced almost four times higher than that of the V20~-coated catalytic membrane that did not participated in the reaction. (see, to Fig.6 when the feed gas was nitrogen only.) The enhanced permeability of oxygen through V2Os-coated catalytic membrane at 350~ was most likely due

1238 to the partial oxidation reaction based on Redox mechanism of 1-butene. But further studies must be conducted for a more precise explanation regarding the enhancement of permeability. 12 o 10

g

m

6

•

~ 4 9

E o.

2

_

V2Os-coated Vycor

B

I

o

0

1oo

200

I

I

300

400

Temperature (~ Fig.6 P e r m e a b i l i t y of 0 2 with t e m p e r a t u r e at p r e s s u r e d i f f e r e n c e 15psi ( feed gas

9 /X I I , 1-butene and N2 mixture

OO

,N2

)

4. CONCLUSION The V2Os-coated catalytic membrane was applied to the partial oxidation reaction to produce maleic anhydride from 1-butene. The maximum selectivity was 95% at 350~ and the permeability of oxygen through the V2Os-coated catalytic membrane was enhanced. More careful studies are needed about this permeability enhancement.

REFERENCES 1. 2. 3. 4. 5. 6.

G. A. F. G. A. J.

Centi and F. Trifiro, Chem. Rev., 88(1988) 55 Tonkovich et al., Chem. Eng. Sci., 51(1996) 789 Cavani, Ind. Eng. Pro. Res. Dev., 22(1983) 565 Saracco and Vito Specchia, Catal. Rev. Sci. Eng. 36(1994) 305 Champagnie et al., Chem. Eng. Sci., 45(1990) 2423 Bullot et al., J. Non-Cryst. Solids, 68(1984) 123

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

Peroxidase a n o n

1239

of phenol by catalase immobilized on carbon materials

E.Horozova, N.Dimeheva and Z.Jordanova Department of Physical Chemistry, University of Plovdiv; 24, Tsar Assen St, Plovdiv - 4000, Bulgaria SUMMARY The peroxidase activity of immobilized catalase on the oxidation of phenol has been studied. The immobilization was carded out from eatalase solutions with pH < 3,5 on two kinds of soot differing in the average size of the particles building them up. The effect of the initial concentration of phenol on the rate of its peroxidase oxidation by catalase immobilized on the soot of finer-grained structure has been studied. The relationships obtained are described by the equation of Michaelis-Menten. The kinetic parameters (the constant of Miehaelis - Km, the maximum reaction rate - V, the rate constant - k and the activation energy - E a of the process were calculated. It was found that catalase adsorbed on the soot of larger globular particles does not take part in the peroxidase oxidation of phenol. INTRODUCTION In bioeatalytic systems, eatalase is mainly used in immobilized state. High activity of immobilized catalase was achieved on its sorption immobilization on cellulose [6], on silica gel modified with fatty acids or phospholipids [7] as well as on activated carbon fibres and fabrics/tissues/[8]. Biocatalytie activity of catalase immobilized on cellulose was also studied in nonaqueous solvents [9,10]. In [9] it was found that unlike the enzyme dissolved in waterdimethylformamide medium, on the oxidation of o-dianisidine in the presence of dimethylformamide the immobilized eatalase does not show any peroxidase activity. It was used [10] for working out an organic-phase amperometric biosensor by immobilizing the enzyme in a polymeric film on a glass-carbon surface. The effect of the polyacrylamide pad for immobilization on the thermodynamic activity of immobilized eatalase was studied in [ 11]. Catalase is also used in co-immobilization with other enzymes such as glucose oxidase [12,13], lactate dehydrogenase [14], peroxidase [15], for creating enzyme membranes needed on the determination of the respective substrates. To study the possibility for peroxidase activity of catalase initiated by the process of its immobilization on carbon materials is he goal of the present work. The peroxidase function of eatalase in solutions was well studied on the oxidation of ethyl alcohol, phenol [ 16,17] and arornatie amines [18]. However, there is no data about peroxidase activity of catalase in immobilized state. Such data is greatly needed because eatalase monomers are stable in acid solutions in which the use ofperoxidase is impossible due to its deactivation.

1240 MATERIALS AND METHODS The catalase used in this work was EC 1.11.1.6 from Penicillium chrysogenum 245 (Biovet -Bulgaria), M, = 244 000. -! The specific activity of the enzyme is 1000 U x m g . The reagents for the solutions: Na2HPO4.12H=O, citric acid, phenol and H202, all with analytical grade qualification "pure for analysis". The solutions were prepared with bidistilled water. Carbon materials: soot "NORIT" with fine-grained structure - the average size of particles being 5 x 104 - 45 x 104A, and "PM-100" built up of larger globular particles with an average size of 21 x 104 - 340x 104A. The adsorption of catalase on both types of soot was performed by an adsorption method under static conditions from 1 ml solution of catalase with concentration of the enzyme C=lx 10"4M in citrate buffer (pH =3.02) per 10 mg soot in 24 hours. The amount of catalase was determined spectrophotometrically by the decrease in the concentration of the catalase in the solution after the adsorption. The spectrophotometric measurements were carried out on a Specord UV VIS (Carl Zeiss, Jena, Germany). The amount of the catalase in the solution was determined on the basis of a calibration graph (Fig. l-b) for the maximum at ~,~ = 280nm. The value of the extinction coefficient was ees0 = 0.9x 1051.mol'~.crff1. The peroxidase activity of catalase in solution and in immobilized state on both types of soot was studied on the oxidation of phenol. The kinetics of the enzyme reaction was monitored spectrophotometricaUy by the decrease in the concentration of the substrate at = 270 nm A

- -

.

I

.

C"

.

"

"-

III

II

1

A

II

0.7

0.7

0.6

0.6

0.5

0.5

0.4

0.4

0.3

0.3

0.2

0.2

O.l

0.1

.

280

300

~,, nm

.

.

.

.

: . .

__

: -

:-.r_-

::

~.-.~

-.:

.-

b

/.

260

.

.

i

I

..

. J.

2

. . . ..

t .

3

.

i

4

_.

.

._

1

. .

5

.l

6

..

l_

?

C x 10', mol.l ~ Figure 1. Catalase absorption spectra (a) and a calibration graph (b) for determination of the catalase concentration in the solution; a: catalase concentration in citrate buffer (pH = 3.02) l x l 0 "4M; 1 cm cell.

1241 RESULT AND DISCUSSION On the adsorption of catalase from solutions with a concentration of C = 1x lO4M in citrate buffer with pH = 3.02, a dissociation of the enzyme to subunits takes place [2,18]. No matter how the catalase dissociation is initiated, the dissociation process is always accompanied by a certain loss of catalase activity and optic changes - a shift towards the shortwave region [ 19]. The catalase tetramer dissociation to monomers [ 17] runs according to scheme E4 r 4E. For dissolved catalase and for eatalase immobilized on soot "NORIT" the dependence of the initial oxidation rate of phenol on the initial concentration of the substrate was studied (Fig.2), where A is the absorbanee corresponding to the current concentration of the substrate (phenol). The obtained relationships are described by the equation of Michaelis-Menten :

v..~iSl v = K. +lSl

(])

where V is the rate of the enzymatic reaction, and Vm~- the maximum rate, i.e. the rate at which the ~ e - s u b s t r a t e complex react; K ~ - Michaelis constant, equal to such a concentration of the substrate at which the rate of the enzymatic reaction is half of its maximum; [S] - the concentration of the substrate. A transformation of these relationships by the method of Lineweaver - Burk brought about to equation:

1

K= 1

v-

v..~ s + v..--~-

1

(2)

used to built up a graph in coordinates 1/V - 1/[S] (Fig.3). The obtained straight line intersects K. the X and Y axes. The slope gives ( V=u )' the intercept of the Y-axis ( 1/Vm~ ), the intercept of the X-axis (-1/K~). The values obtained for K~ and V=,~ are given in Table 1.

Table 1. Kinetic parameters ofper0xidase oxidation ofphen01 by eatalase Catalase

Kin, M

V~,

l~t 285 K

E,, kJ.mol q

,S "1

303 K

In solution 5.48x104 1.43 0.9x10 "4 1.3xlO "4 14.56 C=1.36x lO'SM Immobilized on "NORIT" 2.73• .3 16.70 7.3x10 "4 9.5x10 "4 10.95 -~ ~ 0 . 3 7 m g . . . . . . ................... . ........................................ Immobilized DOES NOT SHOW PEROXIDASE ACTIVITY on "PM- 100"

1242 ~ :

A 0.7

=~-

_

.

.

.

.

.

.

.

.

.

.

.

.

.

:

.

-~

.~

......

V-I

1

V-~<103

2

0.6

1.2

6

0.5 0.4

4

0.8 1

0.3 0.2

2

0.4

0.1 . . . .

0

i.

1

.

i

2

.

i

3

I . . . . .

4

~ ....

i_

5

6

i

,,,

7

1

-2

0

2

Cx 10*,tool.l"

4 6 S -txl0 "~,mol~.l

8

Figure 2. Dependence of the rate of peroxidase oxidation of phenol on the start concentration of the substrate: 1. catalase in solution C = 1.36• 10"5M; 2. catalase immobilized on "NORIT", enzyme amount g = 0.37 rag. Concentration of H, O2 C= 2x 109 M; Citrate buffer pH = 3.02

Figure 3. Lineweaver- Burk plot of peroxidase oxidation of phenol by catalase 1. catalase in solution with concentration C = 1.36x 105 M; 2. catalase immobilized on "NORIT", enzyme amount g = 0.37 mg Concentration of H202 C= 2• 10.9 M; Citrate buffer pH = 3.02.

The study of phenol oxidation with catalase immobilized on soot "PM-100" showed that the catalase adsorbed on that carder did not indicate peroxidase activity, it follows that the catalase subunits obtained on its dissolving in the citrate buffer with pH=3.02 are deactivated and lose their peroxidase activity when immobilized on that type of soot. This behaviour of catalase adsorbed on "PM-100" can be explained with the larger average size of the particles (21 x 104 - 340x 104A) of that soot. Probably, on the adsorption of catalase monomers on those larger particles, favourable conditions for monomers bonding set in and it brings to the recovery of the quaternary structure of the enzyme which only has a catalase function. The finer-gained structure of the ~NORIT" soot with an average size of the particles 5x 104 45x 104 A) on the one hand is conducive to the adsorption of the catalase subunits and on the other to a change of the conformation and the structure of the protein undissociated on its immobilization. Thus, on the immobilization of catalase on "NORIT" soot, there is a possibility of a partial liberation of catalase subunits [9] which are adsorbed to the carder. Therefore the "NORIT" soot samples with adsorbed catalase show peroxidase activity. The rate constants (Table 1.) for the peroxidase oxidation of phenol at various temperatures were calculated by the kinetic equation for a first order reaction (Fig.4). From the data in Table 1 is seen that on peroxidase oxidation of phenol, the catalase immobilized on "NORIT" soot shows a higher catalytic activity than the catalase in solution. It can be explained first, with the fact that the values in Table 1 are not the values for the specific rate constants, and secondly, as the catalase peroxidase function is characteristic of its subunits, with the complete or partial dissociation ofcatalase on its immobilization on "NORIT" soot. In this

1243

-In A 0.5

\

0.6

0.7 I 0.8

......... 10 20 30 t, min Figure 4. Relationship lnA - t for peroxidase oxidation phenol by catalase immobilized on "NORIT" soot. Enzyme amount g = 0.41 rag., temperature 30 ~ Citrate buffer pH = 3.02.

way, not only the catalase monomers obtained on its dissolving in a citrate buffer with pH=3.02 take part in the oxidation of phenol, but also those obtained on its adsorption. Probably this is the reason for the increase in the peroxidase activity of catalase in immobilized state. An increase in temperature by 18oc for both catalase in solution and catalase in immobilized state causes the same increase in the rate of the peroxidase reaction - about 1.3 1.4 times. The activation energy Of the process (Table 1) was calculated by the equation

km2 - E.(I Ink~1 R -T-1-1-Tzl

(3)

The values obtained for Ea fully correspond to the values of the rate constants. For catalase adsorbed on "NORIT", the rate of the peroxidase oxidation of phenol is higher because the activation energy Ea of this process is lower. On the basis of E, values we can make a conclusion about the diffusion factors which are some of the most complicated points concerning catalysis with immobilized enzymes. The value for the activation energy on peroxidase oxidation of phenol with catalase immobilized on "NORIT" soot is E, -10.95 kJ.mol~ which is an indication that the process takes place under diffusion regime. The latter means that the enzymatic reaction rate is determined by the mass tranfer of substrate to the surface of the carrier particles and its diffusion into the carrier.

REFERENCES 1. D.I. Metelitsa (A.E Shilov-ed), Modelirovanie okislitelno-vostanovitelnih fermentov, Publ. Nauka i Technika, Minsk, Russia, 1984 2. VD. Artemchik, V.P. Kurchenko and D.I. Metelitsa, Biochimiya, 50 (1985) 826

1244 3. V.D. Artemchik, V.P. Medveentsev and D.I. Metelitsa, Biochimiya, 51 (1986) 1355 4. M. Dixon and E. Webb, (A.I. Oparin-ed) Fermenty, Publ. "Mir", Moskva, Russia, 1966 5. M.N. Hughes (G. Shklaeva-ed), Neorganicheskaya chimiya biologicheskih protsessov, Publ. "Mir", Moskva, Russia, 1983 6. I.V. Berezin, V.E. Antonova and K.M. Martinska (eds), Immobilizovannie fermenty. Sowemennoe sostoyanie i perspektivy, Publ. Moskovskogo gossudarstvennogo universiteta, Russia, 1976 7. M.N. Veselova, E.C. Chuhray and O.M. Poltorak, J. physicheskoy chimiy, 48 (1974) 2019 8. A.V. Litvinchul~ A.A. Morozova, M.I. Savenkova and D.I. Metelitsa, Prikladnaya biochimiya i microbiologiya, 30 (1994) 572 9. A.N. Eremirg C.V. Omtskiy and D.I. Metelitsa, Kinetica i kataliz, 36 (1995) 847 10. J. Wang, G. Rivas and J. Lin, Anal. Letters, 28 '13 (1995) 2287 11. B. Jang and Y. Zhang, Europ. Polym J., 29 (1993) 1251 12. C.C. Lirg L.B. Wingard, S.K. Wolf,son, S.J. Jao, A.L. Drach and J.G. Scheller, Bioelectrochem Bioenerg., 6 (1979) 19 13. L.B. Wingard Jr, L.A. Candin and J.F. Castner, Biochim Biophys. Acta, 21 (1983) 748 14. F. Scheller, N. Seigbahn, V. Danielsson and K. Mosbach, Anal. Cshem., 57 (1985) 1740 15. T. Tatsuma, T. Watanabe and S. Tatsuma, Anal. Chem., 66 (1994) 290 16. U.M. Arizov, G.I. Mihtenstein and A.P. Purmal, Biochimiya, 37 (1972)620 17. M.N. jones, P. Manley, P.J. Midley and A.E. Wilkinson, Biopolymers, 21 (1982) 1435 18. H. Stmd, K. Weber and E. Molbert, Europ. J. Biochem., 1 (1967) 400

1245

AUTHOR INDEX

A Abon, M . - 209 Agaskar, P.A. - 219 Ahn, J . H . - 957 Ai, M . - 527 Albonetti, S. - 403 Alejo, L. - 623 Alibrando, M. - 693 Al'kaeva, E.M.- 939 Andersen, A.G. - 701 Andersson, A. - 413 Andrushkevich, T . V . - 939 Ang, S.S. - 471 Arena, F . - 347 Arends, I.W.C.E.- 1031 Aubry, J.M. - 883 Augugliaro, V. - 663

B Bacherikova, I.V.- 337, 787 Bagnasco, G. - 643 Balkus, Jr., K.J.- 999, 1007 BaSares, M . A . - 295 Barbaux, Y . - 383 Barrault, J. - 375 Batiot, C . - 375, 1097 Batista, J. - 633 Bebelis, S.I.- 77 Behrens, P. - 817 Belkhir, I.- 577 Belousov, V . M . - 1175 Ber~,i~,, G. - 633 Bere, K.E.- 209 Bernd, S. - 103 Bernier, A. - 777 Beschmann, K . - 929 Besson, M. - 615 Bethell, D . - 535 Bethke, G.K. - 453 Beziat, J.C. - 615 Bini, G . - 721 Blanchard, G. - 403 Bogutskaya, L.V.- 337, 787 Bonchio, M. - 865 Bordes, E . - 177

Bratinova, S . - 909 Br6geault, J.-M. - 545 Breckner, A. - 919 Brungs, A.J. - 711 Bueno, J.M.C. - 453 Bulgakov, N.N.- 1155 Bulman Page, P.C.- 535 Burattin, P. - 403 Burgina, E.B.- 939, 1203 Burgina, L.B.- 1155 Burrows, A. - 209 Busca, G . - 643, 989 Buttrey, D.J. - 219 Bychkov, V . Y u . - 327 C Cai, Y. - 255 Canesson, L. - 893 Carneiro da Rocha, M.G. 767 Carofiglio, T. - 865 Carraz~n, S.R.G.- 197 Cassidy, F.E.- 1097 Castillo, G. - 737 Cavani, F . - 19 Centi, G. - 893 Chagas, A.L. - 1193 Chen, SongYing - 829 Chen, Y . - 903 Ching, Chi-Bun - 683 Chornaja, S.S.- 585 Ciambelli, P. - 285 Claridge, J.B.- 711 Coenraads, R. - 443 Coluccia, S . - 347, 663 Cortes Corberan, V . - 747 Courtine, P . - 177

D Dai, Q.-L.- 1107 Dalai, M.K. - 1165 Dejoz, A. - 443 Delmon, B. - 43, 185, 197, 481,517, 777, 1213 Deloffre, A. - 545 Detusheva, L.G. - 1203 Devillers, M. - 517

DeVos, D.E.- 973, 1051 Dick, S. - 1061 Dimcheva, N . - 1239 Doi, T . - 245, 835 Doren, D. - 1081 Dowell, D . - 199 Doyle, A.M. - 1097 Dur~cu, S. - 615

E Emig, G . - 593, 847 Estenfelder, M . - 981

F Fache, E. - 577 Fajula, F . - 577 Fan, YanZhen - 829 Fan, Y. - 903 Farinha Portela, M. - 807 Fazzini, F. - 893 Fedotov, M . A . - 1203 Fierro, J.L.G. - 295, 623, 721 Finocchio, E. - 989 Finol, D . M . - 737 Firth, P . G . - 535 Flid, M.R. - 305 Folgado, J.V. - 747 Frei, H . - 1043 Frety, R. - 767 Frusteri, F . - 347 Fuchs, S . - 929 Fukui, K . - 227

G Gaffney, A.M. - v Gaigneaux, E . - 777 Gaigneaux, E.M. - 185, 197 Gallezot, P. - 615 Ganne, M. - 375 Gao, X. - 295 Gaube, J . - 393 Genet, M. - 1061 Germain, A. - 577 Ghenne, L . - 197 Gipe, R.K. - 1117

1246

Gleason, N.R. - 235 Gleaves, J.T.- 199 Gonzalez, S. - 1213 Goralski, Jr., C.T. - 491 Grange, P . - 777 Grasselli, R.K. - v, 219, 357, 817 Green, M.L.H. - 711 Grigorjeva, I.A.- 585 Grobet, P.J.- 973 Gross, M.J. - 857 Groves, J.T. - 865 Guo Xiu-Mei - 683 Gutman, A. - 315 Guzhnovskaya, T.D.- 305 H

Haber, J . - 1,337 H~fele, M. - 847 Hahn, T . - 929 Hamakawa, S . - 701, 1223 Hancock, F. - 535 Hannour, F.K. - 919 Hansen, S. - 413 Haruta, M . - 123, 965 Hayakawa, T . - 701, 1223 Hayashi, H. - 673 Hayashi, T . - 965 Herein, D . - 103 Herskowitz, M. - 315 Hidajat, Kus - 683 Hiltner, H . - 593 Hodges, G. - 653 Hodnett, B.K.- 1097, 1129, 1137 Hong, Sukin - 1231 Horozova, E . - 1239 Hou, Z.-Y. - 1107 Huff, M.C. - 501 Hutchings, G.J.- 209, 535 I

Idriss, H . - 265 Ihm, S.K. - 957 Inumaru, K. - 35 Isupova, L.A.- 1155 Ivkova, I.V.- 731 Iwasawa, Y . - 227

Jacobs, P.A.- 973, 1051,

1061 Jalowiecki-Duhamel, L.383 John, J . - 1165 Jones, C . W . - 603 Jordanova, Z . - 1239 Jovanovi6, N.N. - 1145 Juif, S . - 615 Jung, Jihoon - 1231 K

Kaliya, M.L.- 315 Kalvachev, Y.A. - 965 Kanai, H. - 673 Karakas, G. - 471 Kashuba, E.V.- 1175 Khanmamedova, A.K. 999, 1007 Kharitonov, A.S. - 857 Kharlamov, A.I. - 337 Khavrutskii, I.V.- 947 Kholdeeva, O.A.- 947 Kiely, C.J. - 209 Kiennemann, A. - 807 Kim, J.C. - 957 Kim, Tayoon- 1231 King, F. - 535 Klemm, E. - 847 Knops-Gerrits, P.P. - 1061 KnOzinger, H. - 817 Kogan, L.O. - 315 Kolmiichuk, V.N.- 1155 Konecny, R. - 1081 Korchak, V.N. - 327, 757 Koscheev, S.V.- 1203 Koyano, G. - 35 Krepel, A. - 461 Kretschmer, A. - 461 Krylov, O.V. - 275 Kryukova, G.N. - 939, 1155 Kuang, W . - 903 Kubias, B. - 461 Kung, H.H. - 453 Kung, M.C. - 453 Kunkeler, P.J.- 1015 Kurlyandskaya, 1.1.- 305 Kustova, G.N.- 1155 Kuwert, G. - 423 Kuznetsova, L.I.- 1203 Kuznetsova, N.I.- 1203

Lago, R.- 623, 721

Lahousse, C . - 777 Landa-C~novas, A.R. - 413 Landau, M.V. - 315 Lee, D . F . - 535 Lee, Jong Koog- 1183 Lee, Wha Young - 1183 Lempers, H.E.B.- 557 Lepetit, C . - 545 Levec, J. - 633 Levin, D. - 367 Levina, A . B . - 585 Levinbuk, M.I.- 731 Li, W . - 433, 873 Li, X.Y. - 1061 Likholobov, V . A . - 1203 Lindsay Smith, J . R . - 653 Lintz, H.-G.- 981 Lisi, L. - 285 Loddo, V. - 663 Loginov, A.U.- 731 Lopes, M.F.S.- 1193 L0pez Nieto, J.M.- 443 Lorenzelli, V . - 989 Lu, D.-W.- 1107 Lu, X.H. - 35 LOcke, B. - 919 Lugo, H.J.- 737 Lyashenko, L.V.- 787, 1175 Lyons, J.E. - v

Madeira, L.M.- 797 Madhavaram, H. - 265 Magaud, L. - 375 Makkee, M. - 1071 Marchese, L. - 663 Margolis, L.Ya.- 327 M~rquez-Alvarez, C. - 711 Martin, A . - 919 Martinez, E . - 747 Martra, G. - 347, 663 Masetti, S . - 403 Masilamani, D . - 1089 Matsumura, Y. - 673 Matsuura, I.- 245 McMorn, P. - 535 Misono, M . - 35 Miyake, T. - 835 Miyamae, T . - 245 Miyazaki, T . - 245 Mizuno, N. - 35 Moffat, J . B . - 673 Mohammedi, O. - 545

1247

Monnier, J.R. - 135 Moon, Sangjin- 1231 Motoda, K. - 227 Moulard, A. - 197 Moulijn, J.A. - 1071

N Nagy, A. - 103 Naidoo, A. - 423 Nardello, V. - 883 Naud, J . - 185 Neisius, T h . - 103 Ng Lee, Y . - 747 Nijhuis, T.A. - 1071 Nilsson, J. - 413

Oakes, J. - 653 Ohdan, K . - 527 O'Malley, A. - 1137 Otsuka, K.- 93 Oyama, S.T.- v, 873 Ozkan, U.S. - 471

P Palmisano, L. - 663 Panov, G.I. - 857 Pantaleone, M. - 663 Papp, H. - 461 Park, Gyo Ik- 1183 Park, Seungdoo- 1231 Parmaliana, A. - 347 Partenheimer, W. - 1117 Paukshtis, E.A.- 1155 Pavlov, M.L.- 731 Pellizon Birelli, M . - 1031 Peluso, G. - 643 Pefia, M.A.- 623, 721 Pereira, J.A.F.R.- 1193 Pillep, B. - 817 Pintar, A. - 633 Pires, E.L.- 1025 Ponchel, A. - 383 Popva, Z . - 909 Portela, M.F.- 797 Prati, L.- 509

R Ram, R.N. - 1165 Ramis, G . - 643, 989 Raybold, T.M. - 501

Reinaud, O.M.- 1081 Reitzmann, A. - 847 Renard, C. - 517 Richter, F. - 461 Rocha, Jr., A.O. - 1193 Rodkin, M.A. - 857 Roland, U. - 197 Romannikov, V.N.- 947 Rossi, M . - 509 Rouxhet, P.- 1061 Rozentuller, B.V.- 275 Ruiz, P.- 185, 197, 481, 517, 777, 1213 Ruoppolo, G. - 285 Russo, G . - 285, 643

Sadykov, V.A.- 1155 Sajip, S . - 2O9 Salles, L.- 545 Sapifia, F.- 747 Savin, E.M.- 731 Savkin, V.V.- 757 SchedeI-Niedrig, Th.- 103 Schiavello, M. - 663 Schilf, B.T. - 471 SchlOgl, R.- 103 Schmidt, L.D. - 491 Schubert, U.A. - 817 Schuchardt, U.- 1025 Sels, B.F. - 1051 Sergejeva, O.N.- 585 Sexton, A.W. - 1129 Shalyaev, K.V. - 865 Sheldon, R.A.- 151,557, 567, 1031 Sherrington, D.C.- 957 Shi, J . - 999 Shimizu, M . - 701, 1223 Shiozaki, R.- 701 Sinev, M.Yu.- 327, 757 Smirnov, V.K.- 731 Sobalik, Z. - 481, 1213 Sokolovskii, V. - 347 Song, In Kyu- 1183 Spengler, J. - 817 StankoviE, M.V.- 1145 Stein, B.- 393 Stern, D.L.- 357 Stoch, J.- 337 Subba Rao, Y.V.- 973 Sugiyama, S. - 673 Sun, T a o - 1029 Sun, YuHan - 829

SurlY, L.S.V.S.- 1193 Surkova, L.V.- 731 Suzuki, K . - 701, 1223

Tai, N.T.- 275 Takehira, K . - 701, 1223 Tan, K . - 1015 Teran, N.- 737 Theopold, K.H. - 1081 Tikhov, S.F.- 1155 Tkachev, A.V. - 947 Treger, Yu.A.- 305 Trifirb, F . - 19, 403 Trusov, S.R. - 585 Tsang, S.C. - 711 Tsubota, S. - 965 Tsybulya, S.V. - 939, 1155 Tuel, A. - 209, 893, 909 Tulenin, Yu.P. - 757 Turco, M. - 643 LI Ueda, W. - 433 Uriarte, A.K. - 857 Usachev, N.Y.- 731 V Van Bekkum, H . - 1015 Van den Oosterkamp, P.F. -315 Van der Waal, J.C. - 1015 Van Klaveren, M . - 567 Van Langeveld, A.D. - 1071 Van Steen, E. - 423 Vayenas, C.G.- 77 Vayssilov, G.N. - 909 Vazquez, M.I. - 443 V~drine, J.C. - 61 Veniaminov, S.A. - 1155 Vieira Soares, A.P. - 807 Villasmil, L.- 737 Volta, J.C. - 209, 285 W Wachs, I.E. - 255, 295 Wallau, M . - 1025 Wang, C-B. - 255 Wang, D. - 453 Wang, YuanYang - 829 Weimer, C . - 393

1248

Weiss, A. - 1061 Wenkin, M. - 517 Werner, H . - 103 White, B. - 219 Williams, M. - 423 Wilson, S.L. - 603 Wolf, E.E. - 693 Wouters, B. - 973 Wu, X.-Q. - 1107

Ying, J.Y.- 367, 1029 York, A.P.E. - 711

Zaera, F. - 235 Zaikovskii, V . I . - 1155 Zamaraev, K.I. - 947 Zazhigalov, V . A . - 337, 787 Zenkovets, G . A . - 939 Zhang, L e i - 1029 Zhang, W. - 903 Ziani-Derdar, F . - 545

STUDIES IN SURFACE SCIENCE AND CATALYSIS Advisory Editors: B. Delmon, Universite Catholique de Louvain, Louvain-la-Neuve, Belgium J.T. Yates, University of Pittsburgh, Pittsburgh, PA, U.S.A. Volume

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Volume 6 Volume 7 Volume 8 Volume 9 Volume 10 Volume 11

Volume 12 Volume 13 Volume 14 Volume 15

Preparation of Catalysts I.Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings of the First International Symposium, Brussels, October 14-17,1975 edited by B. Delmon, P.A.Jacobs and G. Poncelet The Control of the Reactivity of Solids. A Critical Survey of the Factors that Influence the Reactivity of Solids, with Special Emphasis on the Control ofthe Chemical Processes in Relation to Practical Applications by V.V. Boldyrev, M. Bulens and B. Delmon Preparation of Catalysts I1. Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings of the Second International Symposium, Louvain-la-Neuve, September 4-7, 1978 edited by B. Delmon, P. Grange, P.Jacobs and G. Poncelet Growth and Properties of Metal Clusters. Applications to Catalysis and the Photographic Process. Proceedings ofthe 32nd International Meeting ofthe Societe de Chimie Physique, Villeurbanne, September 24-28, 1979 edited by J. Bourdon Catalysis by Zeolites. Proceedings of an International Symposium, Ecully (Lyon), September 9-11, 1980 edited by B. Imelik, C. Naccache, Y. Ben Taarit, J.C. Vedrine, G. Coudurier and H. Praliaud Catalyst Deactivation. Proceedings of an International Symposium, Antwerp, October 13- 15,1980 edited by B. Delmon and G.F. Froment New Horizons in Catalysis. Proceedings of the 7th International Congress on Catalysis, Tokyo, June 30-July4, 1980. Parts A and B edited by T. Seiyama and K. Tanabe Catalysis by Supported Complexes by Yu.l. Yermakov, B.N. Kuznetsov and V.A. Zakharov Physics of Solid Surfaces. Proceedings of a Symposium, Bechyhe, September 29-October 3,1980 edited by M. Lazni~ka Adsorption at the Gas-Solid and Liquid-Solid Interface. Proceedings of an International Symposium, Aix-en-Provence, September 21-23, 1981 edited by J. Rouquerol and K.S.W. Sing Metal-Support and Metal-Additive Effects in Catalysis. Proceedings of an International Symposium, Ecully (Lyon), September 14-16, 1982 edited by B. Imelik, C. Naccache, G. Coudurier, H. Praliaud, P. Meriaudeau, P. Gallezot, G.A. Martin and J.C. Vedrine Metal Microstructures in Zeolites. Preparation - Properties- Applications. Proceedings of a Workshop, Bremen, September 22-24, 1982 edited by P.A. Jacobs, N.I. Jaeger, P.Jin3 and G. Schulz-Ekloff Adsorption on Metal Surfaces. An Integrated Approach edited by J. Benard Vibrations at Surfaces. Proceedings of the Third International Conference, Asilomar, CA, September 1-4, 1982 edited by C.R. Brundle and H. Morawitz Heterogeneous Catalytic Reactions Involving Molecular Oxygen by G.I. Golodets

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Preparation of Catalysts III. Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings ofthe Third International Symposium, Louvain-la-Neuve, September 6-9, 1982 edited by G. Poncelet, P. Grange and P.A. Jacobs Spillover of Adsorbed Species. Proceedings of an International Symposium, Lyon-Villeurbanne, September 12-16, 1983 edited by G.M. Pajonk, S.J. Teichner and J.E. Germain Structure and Reactivity of Modified Zeolites. Proceedings of an International Conference, Prague, July 9-13, 1984 edited by P.A. Jacobs, N.I. Jaeger, P.JiE=,V.B. Kazansky and G. Schulz-Ekloff Catalysis on the Energy Scene. Proceedings of the 9th Canadian Symposium on Catalysis, Quebec, P.Q., September 30-October 3, 1984 edited by S. Kaliaguine and A. Mahay Catalysis by Acids and Bases. Proceedings of an International Symposium, Villeurbanne (Lyon), September 25-27, 1984 edited by B. Imelik, C. Naccache, G. Coudurier, Y. Ben Taarit and J.C. Vedrine Adsorption and Catalysis on Oxide Surfaces. Proceedings of a Symposium, Uxbridge, June 28-29, 1984 edited by M. Che and G.C. Bond Unsteady Processes in Catalytic Reactors by Yu.Sh. Matros Physics of Solid Surfaces 1984 edited by J. Koukal Zeolites: Synthesis, Structure, Technology and Application. Proceedings of an International Symposium, Portoro~-Portorose, September 3-8, 1984 edited by B. Dr~aj, S. Ho~evar and S. Pejovnik Catalytic Polymerization of Olefins. Proceedings of the International Symposium on Future Aspects of Olefin Polymerization, Tokyo, July 4-6, 1985 edited by T. Keii and K. Soga Vibrations at Surfaces 1985; Proceedings of the Fourth International Conference, Bowness-on-Windermere, September 15-19, 1985 edited by D.A. King, N.V. Richardson and S. Holloway Catalytic Hydrogenation edited by L. Cerveny New Developments in Zeolite Science and Technology. Proceedings of the 7th International Zeolite Conference, Tokyo, August 17-22, 1986 edited by Y. Murakami, A. lijima and J.W. Ward Metal Clusters in Catalysis edited by B.C. Gates, L. Guczi and H. Kn6zinger Catalysis and Automotive Pollution Control. Proceedings of the First International Symposium, Brussels, September 8-11, 1986 edited by A. Crucq and A. Frennet Preparation of Catalysts IV. Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings ofthe Fourth International Symposium, Louvain-la-Neuve, September 1-4, 1986 edited by B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet Thin Metal Films and Gas Chemisorption edited by P.Wissmann Synthesis of High-silica Aluminosilicate Zeolites edited by P.A. Jacobs and J.A. Martens Catalyst Deactivation 1987. Proceedings of the 4th International Symposium, Antwerp, September 29-October 1, 1987 edited by B. Delmon and G.F. Froment Keynotes in Energy-Related Catalysis edited by S. Kaliaguine

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Methane Conversion. Proceedings of a Symposium on the Production of Fuels and Chemicals from Natural Gas, Auckland, April 27-30, 1987 edited by D.M. Bibby, C.D. Chang, R.F. Howe and S. Yurchak Innovation in Zeolite Materials Science. Proceedings of an International Symposium, Nieuwpoort, September 13-17, 1987 edited by P.J. Grobet, W.J. Mortier, E.F. Vansant and G. Schulz-Ekloff Catalysis 1987. Proceedings ofthe 10th North American Meeting ofthe Catalysis Society, San Diego, CA, May 17-22, 1987 edited by J.W. Ward Characterization of Porous Solids. Proceedings of the IUPAC Symposium (COPS I), Bad Soden a. Ts., April 26-29,1987 edited by K.K. Unger, J. Rouquerol, K.S.W. Sing and H. Kral Physics of Solid Surfaces 1987. Proceedings of the Fourth Symposium on Surface Physics, Bechyne Castle, September 7-11, 1987 edited by J. Koukal Heterogeneous Catalysis and Fine Chemicals. Proceedings of an International Symposium, Poitiers, March 15-17, 1988 edited by M. Guisnet, J. Barrault, C. Bouchoule, D. Duprez, C. Montassier and G. Perot Laboratory Studies of Heterogeneous Catalytic Processes by E.G. Christoffel, revised and edited by Z. Paal Catalytic Processes under Unsteady-State Conditions by Yu. Sh. Matros Successful Design of Catalysts. Future Requirements and Development. Proceedings ofthe Worldwide Catalysis Seminars, July, 1988, on the Occasion of the 30th Anniversary of the Catalysis Society of Japan edited byT. Inui Transition Metal Oxides. Surface Chemistry and Catalysis byH.H. Kung Zeolites as Catalysts, Sorbents and Detergent Builders. Applications and Innovations. Proceedings of an International Symposium, WL~rzburg, September 4-8,1988 edited by H.G. Karge and J. Weitkamp Photochemistry on Solid Surfaces edited by M. Anpo and T. Matsuura Structure and Reactivity of Surfaces. Proceedings of a European Conference, Trieste, September 13-16, 1988 edited by C. Morterra, A. Zecchina and G. Costa Zeolites: Facts, Figures, Future. Proceedings of the 8th International Zeolite Conference, Amsterdam, July 10-14, 1989. Parts A and B edited by P.A. Jacobs and R.A. van Santen Hydrotreating Catalysts. Preparation, Characterization and Performance. Proceedings of the Annual International AIChE Meeting, Washington, DC, November 27-December 2, 1988 edited by M.L. Occelli and R.G. Anthony New Solid Acids and Bases. Their Catalytic Properties by K. Tanabe, M. Misono, Y. Ono and H. Hattori Recent Advances in Zeolite Science. Proceedings ofthe 1989 Meeting ofthe British Zeolite Association, Cambridge, April 17-19, 1989 edited by J. Klinowsky and P.J. Barrie Catalyst in Petroleum Refining 1989. Proceedings of the First International Conference on Catalysts in Petroleum Refining, Kuwait, March 5-8, 1989 edited by D.L. Trimm, S. Akashah, M. Absi-Halabi and A. Bishara Future Opportunities in Catalytic and Separation Technology edited by M. Misono, Y. Moro-oka and S. Kimura

New Developments in Selective Oxidation. Proceedings of an International Symposium, Rimini, Italy, September 18-22, 1989 edited by G. Centi and F.Trifiro Olefin Polymerization Catalysts. Proceedings of the International Symposium Volume 56 on Recent Developments in Olefin Polymerization Catalysts, Tokyo, October 23-25, 1989 edited by T. Keii and K. Soga Volume 57A Spectroscopic Analysis of Heterogeneous Catalysts. Part A: Methods of Surface Analysis edited by J.L.G. Fierro Volume 57B Spectroscopic Analysis of Heterogeneous Catalysts. Part B: Chemisorption of Probe Molecules edited by J.L.G. Fierro Introduction to Zeolite Science and Practice Volume 58 edited by H. van Bekkum, E.M. Flanigen and J.C. Jansen Heterogeneous Catalysis and Fine Chemicals I1. Proceedings of the 2nd Volume 59 International Symposium, Poitiers, October 2-6, 1990 edited by M. Guisnet, J. Barrault, C. Bouchoule, D. Duprez, G. Perot, R. Maurel and C. Montassier Chemistry of Microporous Crystals. Proceedings of the International Symposium Volume 60 on Chemistry of Microporous Crystals, Tokyo, June 26-29, 1990 edited by T. Inui, S. Namba and T. Tatsumi Natural Gas Conversion. Proceedings of the Symposium on Natural Gas Volume 61 Conversion, Oslo, August 12-17, 1990 edited by A. Holmen, K.-J. Jens and S. Kolboe Characterization of Porous Solids I1. Proceedings of the IUPAC Symposium Volume 62 (COPS II), Alicante, May 6-9, 1990 edited by F. Rodriguez-Reinoso, J. Rouquerol, K.S.W. Sing and K.K. Unger Preparation of Catalysts V. Scientific Bases for the Preparation of Heterogeneous Volume 63 Catalysts. Proceedings of the Fifth International Symposium, Louvain-la-Neuve, September 3-6, 1990 edited by G. Poncelet, P.A. Jacobs, P. Grange and B. Delmon New Trends in CO Activation Volume 64 edited by L. Guczi Catalysis and Adsorption by Zeolites. Proceedings of ZEOCAT 90, Leipzig, Volume 65 August 20-23, 1990 edited by G. (~hlmann, H, Pfeifer and R. Fricke Dioxygen Activation and Homogeneous Catalytic Oxidation. Proceedings of the Volume 66 Fourth International Symposium on Dioxygen Activation and Homogeneous Catalytic Oxidation, BalatonfLired, September 10-14, 1990 edited by L.I. Sim,~ndi Structure-Activity and Selectivity Relationships in Heterogeneous Catalysis. Volume 67 Proceedings of the ACS Symposium on Structure-Activity Relationships in Heterogeneous Catalysis, Boston, MA, April 22-27, 1990 edited by R.K. Grasselli and A.W. Sleight Catalyst Deactivation 1991. Proceedings of the Fifth International Symposium, Volume 68 Evanston, IL, June 24-26, 1991 edited by C.H. Bartholomew and J.B. Butt Zeolite Chemistry and Catalysis. Proceedings of an International Symposium, Volume 69 Prague, Czechoslovakia, September 8-13, 1991 edited by P.A. Jacobs, N.I. Jaeger, L. Kubelkova and B. Wichterlova Poisoning and Promotion in Catalysis based on Surface Science Concepts and Volume 70 Experiments by M. Kiskinova Volume 55

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Volume 78

Volume79 Volume80 Volume81 Volume82

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Volume85 Volume86 Volume87

Catalysis and Automotive Pollution Control II. Proceedings of the 2nd International Symposium (CAPoC 2), Brussels, Belgium, September 10-13, 1990 edited by A. Crucq New Developments in Selective Oxidation by Heterogeneous Catalysis. Proceedings of the 3rd European Workshop Meeting on New Developments in Selective Oxidation by Heterogeneous Catalysis, Louvain-la-Neuve, Belgium, April 8-10, 1991 edited by P. Ruiz and B. Delmon Progress in Catalysis. Proceedings of the 12th Canadian Symposium on Catalysis, Banff, Alberta, Canada, May 25-28, 1992 edited by K.J. Smith and E.C. Sanford Angle-Resolved Photoemission. Theory and Current Applications edited by S.D. Kevan New Frontiers in Catalysis, Parts A-C. Proceedings of the 10th International Congress on Catalysis, Budapest, Hungary, 19-24 July, 1992 edited by L. Guczi, F. Solymosi and P.Tetenyi Fluid Catalytic Cracking: Science and Technology edited by J.S. Magee and M.M. Mitchell, Jr. New Aspects of Spillover Effect in Catalysis. For Development of Highly Active Catalysts. Proceedings ofthe Third International Conference on Spillover, Kyoto, Japan, August 17-20, 1993 edited by T. Inui, K. Fujimoto, T. Uchijima and M. Masai Heterogeneous Catalysis and Fine Chemicals II1. Proceedings ofthe 3rd International Symposium, Poitiers, April 5-8, 1993 edited by M. Guisnet, J. Barbier, J. Barrault, C. Bouchoule, D. Duprez, G. Perot and C. Montassier Catalysis: An Integrated Approach to Homogeneous, Heterogeneous and Industrial Catalysis edited by J.A. Moulijn, P.W.N.M. van Leeuwen and R.A. van Santen Fundamentals of Adsorption. Proceedings of the Fourth International Conference on Fundamentals of Adsorption, Kyoto, Japan, May 17-22, 1992 edited by M. Suzuki Natural Gas Conversion II. Proceedings ofthe Third Natural Gas Conversion Symposium, Sydney, July 4-9, 1993 edited by H.E. Curry-Hyde and R.F. Howe New Developments in Selective Oxidation I1. Proceedings of the Second World Congress and Fourth European Workshop Meeting, Benalmadena, Spain, September 20-24, 1993 edited by V. Cortes Corberan and S. Vic Bellon Zeolites and Microporous Crystals. Proceedings of the International Symposium on Zeolites and Microporous Crystals, Nagoya, Japan, August 22-25, 1993 edited by T. Hattori and T. Yashima Zeolites and Related Microporous Materials: State of the Art 1994. Proceedings ofthe 10th International Zeolite Conference, Garmisch-Partenkirchen, Germany, July 17-22, 1994 edited by J. Weitkamp, H.G. Karge, H. Pfeifer and W. H61derich Advanced Zeolite Science and Applications edited by J.C. Jansen, M. St6cker, H.G. Karge and J.Weitkamp Oscillating Heterogeneous Catalytic Systems by M.M. Slin'ko and N.I. Jaeger Characterization of Porous Solids III. Proceedings of the IUPAC Symposium (COPS III), Marseille, France, May 9-12, 1993 edited by J.Rouquerol, F. Rodriguez-Reinoso, K.S.W. Sing and K.K. Unger

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Volume90 Volume91

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Volume93 Volume94 Volume95 Volume96

Volume97 Volume98

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Catalyst Deactivation 1994. Proceedings of the 6th International Symposium, Ostend, Belgium, October 3-5, 1994 edited by B. Delmon and G.F. Froment Catalyst Design for Tailor-made Polyolefins. Proceedings of the International Symposium on Catalyst Design for Tailor-made Polyolefins, Kanazawa, Japan, March 10-12, 1994 edited by K. Soga and M. Terano Acid-Base Catalysis I1. Proceedings of the International Symposium on Acid-Base Catalysis II, Sapporo, Japan, December 2-4, 1993 edited by H. Hattori, M. Misono and Y. Ono Preparation of Catalysts VI. Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings of the Sixth International Symposium, Louvain-La-Neuve, September 5-8, 1994 edited by G. Poncelet, J. Martens, B. Delmon, P.A. Jacobs and P. Grange Science and Technology in Catalysis 1994. Proceedings of the Second Tokyo Conference on Advanced Catalytic Science and Technology, Tokyo, August 21-26, 1994 edited by u Izumi, H. Arai and M. Iwamoto Characterization and Chemical Modification of the Silica Surface by E.F.Vansant, P.Van Der Voort and K.C. Vrancken Catalysis by Microporous Materials. Proceedings of ZEOCAT'95, Szombathely, Hungary, July 9-13, 1995 edited by H.K. Beyer, H.G.Karge, I. Kiricsi and J.B. Nagy Catalysis by Metals and Alloys by V. Ponec and G.C. Bond Catalysis and Automotive Pollution Control II1. Proceedings of the Third International Symposium (CAPoC3), Brussels, Belgium, April 20-22, 1994 edited by A. F~ennet and J.-M. Bastin Zeolites: A Refined Tool for Designing Catalytic Sites. Proceedings of the International Symposium, Quebec, Canada, October 15-20, 1995 edited by L. Bonneviot and S. Kaliaguine Zeolite Science 1994: Recent Progress and Discussions. Supplementary Materials to the 10th International Zeolite Conference, Garmisch-Partenkirchen, Germany, July 17-22, 1994 edited by H.G. Karge and J. Weitkamp Adsorption on New and Modified Inorganic Sorbents edited by A. Da,browski and V.A. Tertykh Catalysts in Petroleum Refining and Petrochemical Industries 1995. Proceedings of the 2nd International Conference on Catalysts in Petroleum Refining and Petrochemical Industries, Kuwait, April 22-26, 1995 edited by M. Absi-Halabi, J. Beshara, H. Qabazard and A. Stanislaus 1lth International Congress on Catalysis - 40th Anniversary. Proceedings ofthe 1lth ICC, Baltimore, MD, USA, June 30-July 5, 1996 edited by J. W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell Recent Advances and New Horizons in Zeolite Science and Technology edited by H. Chon, S.I. Woo and S. -E. Park Semiconductor Nanoclusters - Physical, Chemical, and Catalytic Aspects edited by P.V. Kamat and D. Meisel Equilibria and Dynamics of Gas Adsorption on Heterogeneous Solid Surfaces edited by W. Rudzinski, W.A. Steele and G. Zgrablich Progress in Zeolite and Microporous Materials Proceedings ofthe 1lth International Zeolite Conference, Seoul, Korea, August 12-17, 1996 edited by H. Chon, S.-K. Ihm and Y.S. Uh

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Hydrotreatment and Hydrocracking of Oil Fractions Proceedings ofthe 1st International Symposium / 6th European Workshop, Oostende, Belgium, February 17-19, 1997 edited by G.E Froment, B. Delmon and P. Grange Natural Gas Conversion IV Proceedings of the 4th International Natural Gas Conversion Symposium, Kruger Park, South Africa, November 19-23, 1995 edited by M. de Pontes, R.L. Espinoza, C.P. Nicolaides, J.H. Scholtz and M.S. Scurrell Heterogeneous Catalysis and Fine Chemicals IV Proceedings of the 4th International Symposium on Heterogeneous Catalysis and Fine Chemicals, Basel, Switzerland, September 8-12, 1996 edited by H.U. Blaser, A. Baiker and R. Prins Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis. Proceedings ofthe International Symposium, Antwerp, Belgium, September 15-17, 1997 edited by G.E Froment and K.C. Waugh Third World Congress on Oxidation Catalysis. Proceedings of the Third World Congress on Oxidation Catalysis, San Diego, CA, U.S.A., 21-26, September 1997 edited by R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons Catalyst Deactivation 1997. Proceedings ofthe 7th International Symposium, Cancun, Mexico, October 5-8, 1997 edited by C.H. Bartholomew and G.A. Fuentes Spillover and Migration of Surface Species on Catalysts. Proceedings ofthe 4th International Conferenceon Spillover, Dalian, China, September 15-18, 1997 o edited by Can Li and Qin Xin

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