11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis 101 11TH INTERNATIONAL CONGRESS O N CATALYSIS 40TH ANNIVERSARY

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Studies in SurfaceScience and Catalysis Advisory Editors: B. Delrnon and J.T. Yates Vol. 101

IITH INTERNATIONAL CONGRESS ON CATALYSIS = 40TH ANNIVERSARY PART A Proceedingsof the 1I t h ICC, Baltimore, MD, USA, June 30 -July 5,1996 Editors

Joe W. Hightower Department of Chemical Engheering, Rice University, Houston, TX 77251- 1892,USA

W. NicholasDelgass School of Chemical Engineering, Purdue University, West Lafayette, IN 47907,USA

Enriquelglesia Alexis T. Bell Department of Chemical Engineering, University of California, Berkeley, CA 94720-9989,USA

1996 ELSEVIER Amsterdam - Lausanne- New York -Oxford

-Shannon -Tokyo

ELSEVIER SCIENCE B.V. Sara Burgerhartstraat 25 PO. Box 21 1,1000 AE Amsterdam, The Netherlands

ISBN 0-444-81947-9

0 1996 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, PO. 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 referredtothecopyright owner, Elsevier Science B.V., unlessotherwisespecified. 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

V

TABLE OF CONTENTS Part A Preface J.W. Hightower

xix

Plenary Lectures P- 1

P-2

P-3

P-4

P-5

Driving Forcesfor Innovation in Applied Catalysis I.E. Maxwell

1

Constrained Geometry and Other Single Site Metallocene Polyolefin Catalysts: A Revolution in Ole@ Polymerization J.C. Stevens

11

Characterizationand Chemical Design of Oxide Surfaces Y . Iwasawa

21

Photocatalysis:State of the Art and Perspectives K.I. Zamaraev

35

Towards Molecular Design of Solid Catalysts 51

A. Baiker

40th Anniversary Lectures

R- 1

R-2

A Retrospective View of Advances in Heterogeneous Catalysis: 1956-1996,Science R.L. Bunvell, Jr.

63

A Retrospective View of Advances in Heterogeneous Catalysis: 1956-1996 Technology H . Heinemann 69

Catalysis on Carbides, Nitrides, and Sulfides A- 1

An Example of Novel Basic Catalysts: The Aluminophosphate Oxynitrides or “AIPONs” A. Massinon, E. Gueguen, R. Conanec, R. Marchand, Y. Laurent, and P. Grange

The figures before the articles indicate the numbers used during the Conference

77

vi A-2

A-3

A- 5

A-6

Reaction Kinetics of the Hydrodenitrogenation of Decahydroquinoline over NiMo(P)/AI,O, Catalysts M. Jim and R. Prins

87

Effect of Spillover Hydrogen on Amorphous Hydrocracking Catalysts A.M. Stumbo, P. Grange, and B. Delmon

97

Hydrodesulfirization of Benzothiophene Catalyzed by Molybdenum Surfde Cluster Encapsulated into 5eolites M. Taniguchi, S. Yasuda, Y. Ishii, T. Murata, M. Hidai, and T. Tatsumi

107

Role of Adsorbed Hydrogen Species on Ruthenium and Molybdenum Suljdes. Characterization by Inelastic Neutron Scattering, Thermoanalysis Methods and Model Reactions M. Lacroix, H. Jobic, C. Dumonteil, P. Afanasiev, M. Breysse, and S. Kasztelan

117

General Papers A-7

A-8

A-9

A-10

Characterization of a Zeolite Membranefor Catalytic Membrane Reactor Application A. Giroir-Fendler, J. Peureux, H. Mouanega, and J.-A. Dalmon

127

Catalytic Reduction of SO, Stored in So, Transfer Catalysts - A TemperatureProgrammed Reaction Study G. Kim and M.V. Juskelis

137

Effect of Tunnel Structures of BaTi,O, and Na,Ti,O,, on PhotocatalyticActivity and Photoexcited Charge Separation M . Kohno, S . Ogura, K. Sato, and Y. Inoue

143

Organochromium Complexes in Homogeneous Olejn Polymerization G. Bhandari, J.L. Kersten, R.R. Kucharczyk, P.A. White, Y. Liang, and K.H. Theopold

153

Synthesis of Fine Chemicals A-1 1

A-I2

A-I3

A- 14

Selective Catalytic Oxidation with Air of Glycerol and Oxygenated Derivatives on Platinum Metals P. Fordham, R. Garcia, M. Besson, and P. Gallezot

161

Selective Methylation of Catechol: Catalyst Development and Characterisation L. Kiwi-Minsker, S. Porchet, P. Moeckli, R. Doepper, and A. Renken

171

Selective Oxidation with Copper Complexes incorporated in Molecular Sieves R. Raja and P. Ratnasamy

181

Discovering the Role of Au and KOAc in the Catalysis of VinylAcetate Synthesis W.D. Provine, P.L. Mills, and J.J. Lerou

191

vii A-15

A-16

A-17 A-18

A-19 A-20

A-2 1 A-22

A-23

Formation of Citraconic Anhydride by Vapor-Phase Decarboxy-Condensation of Pyruvic Acid M. Ai and K. Ohdan

201

Heterogeneous Enantioselective Dehydration of Butan-2-01 S . Feast, D. Bethell, P.C.B. Page, M.R.H. Siddiqui, D.J. Willock, G.J. Hutchings, F. King, and C.H. Rochester

21 1

Enantioselective Hydrogenation Catalysed by Palladium T.J. Hall, P. Johnston, W.A.H. Vermeer, S.R. Watson, and P.B. Wells

22 1

Stereochemical Studies of the Enantio-diferentiating Hydrogenation of Various Prochiral Ketones over Tartaric Acid-Mod$ed Nickel Catalyst T. Sugimura, T. Osawa, S. Nakagawa, T. Harada, and A. Tai

23 1

Enantio-diflerentiation over Heterogeneous Catalysts. The Shielding Eflect Model J.L. Margitfalvi, M. Hegedus, and E. Tfirst

24 1

Racemization of (IS)-(-)-exo-2,4-Dideuteroapopineneover Pd: Evidence for an Intramolecular I,3-Deuterium Shij G.V.Smith, B. Rihter, A. Zsigmond, F. Notheisz, and M. Bartok

25 1

Fe/MgO Catalysts for the Selective Hydrogenation of Nitriles G. Bond and F.S. Stone

257

Selective Synthesis of Ethylenediaminefrom Ethanolamine over Modified H-Mordenite Catalyst K.Segawa, S . Mizuno, M. Sugiura, and S. Nakata

267

Competitive Reaction Pathways in Propane Ammoxidation over V-Sb-Oxide Catalysts: An IR and Flow Reactor Study G. Centi and F. Marchi

277

Research on Model Catalysts

A-25

A-26

A-27

A-28

‘Seeing’the Active Site in Catalysis. STM and Molecular Beam Studies of Surface Reactions M. Bowker

287

Catalytic Formation of Carbon-Carbon Bonds in Ultrahigh Vacuum: Cyclotrimerization of Alkynes on Reduced TiO, Surfaces K.G.Pierce, V.S.Lusvardi, and M.A. Barteau

297

A New Approach to Understanding the Rochow Process: Synthesis of Methylchlorosilanesfrom CH, + C1 Monolayers on Cu,Si in Vacuum D.-H. Sun, A.B. Gurevich, L.J. Kaufman, B.E. Bent, A.P. Wright, and B.M. Naasz

307

Ruthenium as Catalystfor Ammonia Synthesis M. Muhler, F . Rosowski, 0.Hinrichsen, A. Homung, and G. Ertl

317

viii A-29

A-3 0

Surface-Structure-DependentReaction Pathways of Methyl Groups on Ni(lO0) and Ni(ll1) Surfaces R.B. Hall, M. Castro, C.M. Kim, and C.A. Mims

327

Normalization by oxygen Uptake of the Rates of Oxidative Dehydrogenation of Methanol and Ethanol S . Tanabe, H.E. Davis, Jr., D. Wei, and R.S. Weber

337

Environmental Catalysis A-3 1

A-32

A-33

A-34

Some Evidences of a Bifunctional Mechanismfor the Reduction of NO on Pd Based Catalysts H. Praliaud, A. Lemaire, J. Massardier, M. Prigent, and G. Mabilon

345

Characterization of Pd-based Automotive Catalysts R.W. McCabe and R.K. Usmen

355

Process Development for the Selective Hydrogenolysis of CCIF, (CFC-12) into CHZF2 (HFC-32) A. Wiersma, E.J.A.X. van de Sandt, M. Makkee, H. van Bekkum, and J.A. Moulijn

369

Catalytic Fluorination over Chromium Oxides. Preparation of Hydrofluorocarbons S. Brunet, B. Boussand, and J. Barrault

3 79

Oxidation of Methane A-35

A-36

A-37

Oxidative Coupling of Methane to Ethylene with 85% Yield in a Gas Recycle Electrocatalytic or Catalytic Reactor-Separator M. Makri, Y . Jiang, I.V. Yentekakis, and C.G. Vayenas

3 87

The Active w g e n for Direct Oxidation of Methane to Methanol in the Presence of Hydrogen Y . Wang and K. Otsuka

397

Photocatalytic Production of Methanol and Hydrogenfrom Methane and Water C.E. Taylor, R.P. Noceti, J.R. D’Este, and D.V. Martello

407

Oxidation Catalysis A-38

A-39

A-40

Role of A- and B-Cations in Catalytic Property of Substituted Hexaaluminate (ABAI,,O,, a) for High Temperature Combustion K. Eguchi, H . Inoue, K. Sekizawa, and H. Arai

417

AdMetal Oxidesfor Low Temperature CO Oxidation G. Srinivas, J . Wright, C.-S. Bai, and R. Cook

427

Photocatalytic Oxidation of Air Contaminants by Chlorine (CL!) or Hydroxyl (OH) Radicals or Holes @+): Mechanistic Correlations 0.d’Hennezel and D.F. Ollis

43 5

ix

A-4 1

A-42

A-43

A-44

A-45

A-46

A-4

A-24

Partial Oxidation of Methane to Synthesis Gas over Ru/TiO, Catalysts Y. Boucouvalas, Z.L. Zhang, A.M. Efstathiou, and X.E. Verykios

443

The Oxidative Transformation of Methane over the Nickel-Based Catalysts Modified by Alkali Metal Oxide and Rare Earth Metal Oxide Q. Miao, G. Xiong, S. Sheng, W. Cui, and X. Guo

453

Design of Stable Catalystsfor Methane-Carbon Dioxide Reforming J.A. Lercher, J.H. Bitter, W. Hally, W. Niessen, and K. Seshan

463

Comparison of Perovskite and Hexaaluminate-type Catalysts for CO/H,-fieled Gas Turbine Combustors C . Cristiani, G. Groppi, P. Forzatti, E. Tronconi, G. Busca, and M. Daturi

473

Hydrocarbon Activation and Oxidation on Transition Metal Mixed Oxides. Ft-IR and Flow Reactor Studies E. Finocchio, R.J. Willey, G. Ramis, G. Busca, and V. Lorenzelli

483

Biomimetic Oxidation on Fe Complexes in Zeolites (3.1. Panov, V.I. Sobolev, K.A. Dubkov, and A S . Kharitonov

493

A Molecular Approach to Synergy Generation in Co-Mo Binary Sulfide Catalysi‘s for Hydrodesulfirization Y. Okamoto and H. Katsuyama

503

Electrochemical Promotion of NO Reduction by CO and by Propene A. Palermo, M.S. Tikhov,N.C. Filkin, R.M. Lambert, I.V. Yentekakis, and C.G. Vayenas

5 13

Catalysis on Acids and Bases B- 1

B-2

B-3

B-4

B-5

B-6

Promotion of Molecular Hydrogen on Solid Acid Cracking Activity T . Shishido, T. Nagase, K. Higo, J. Tsuji, and H. Hattori

523

Selective Isomerization of Alkanes on Supported Tungsten Oxide Acids E. Iglesia, D.G. Barton, S.L. Soled, S. Miseo, J.E. Baumgartner, W.E. Gates, G.A. Fuentes, and G.D. Meitzner

533

Tungsta and Platinum-Tungsta Supported on Zirconia Catalysts for Alkane Isomerization G. Larsen, E. Lotero, and R.D. Parra

543

n-Butane Isomerization on Ni-promoted Sulfated Zirconia Catalysts W.E. Alvarez, H. Liu, E.A. Garcia, E.H. Rueda, A.J. Rouco, and D.E. Resasco

553

Rare Earth Modified Silica-Aluminas as Supportsfor Bifunctional Catalysis S.L. Soled, G. McVicker, S. Miseo, W. Gates, and J. Baumgartner

563

What NMR Has Told Us about Solid Acidity J.F. Haw and J.B. Nicholas

573

X

B-7

B-8

Novel Microporous Solid “Superacids”:CsjY3J W,,O,, (2sxs3) T. Okuhara, T. Nishimura, and M. Misono

581

Comparison of the Reactivities of H,P W,,O,, and HJi W,,O,, and their K’,NH,’ and Cs’ Salts in Liquid Phase Isobutane/Butene Allylation N. Essayem, S. Kieger, G. Coudurier, and J.C. Vedrine

59 1

B-9

Coupling of Alcohols to Ethers: the Dominance of the Surface SN2Reaction Pathway K. Klier, Q. Sun, O.C. Feeley, M. Johansson, and R.G. Herman 601

B-10

Characterization of Two Different Framework Titanium Sites and Quantijication of Extra-famework Species in TS-I Silicalites L. Le Noc, D. Trong On, S. Solomykina, B. Echchahed, F. Beland, C. Cartier dit Moulin, and L. Bonneviot

61 1

Environmental Catalysis

B-11

B-12 B-13 B-14

B-15 B-16

B-17

B-18

B-19

Injluence of Suljiu Dioxide on the Selective Catalytic Reduction of NO by Decane on Cu Catalysts F. Figueras, B. Coq, G. Mabilon, M. Prigent, and D. Tachon

62 1

Copt Clusters in NaMordenites as Catalystsfor SCR of No, L. Gutierrez, A. Ribotta, A. Boix, and J. Petunchi

63 1

Decomposition of Nitrous Oxide over ZSM-5 Catalysts F. Kapteijn, G. Mul, G. Marbin, J. Rodriguez-Mirasol, and J.A. Moulijn

64 1

On the Role of Free Radicals NO, and 0, in the Selective Catalytic Reduction (SCR) of No, with CH, over CoZSM-5 and HZSM-5 Zeolites D.B. Lukyanov, J.L. d’Itri, G. Sill, and W.K. Hall

65 1

An Infared Study of NO Reduction by CH,Over Co-ZSM-5 A.W. Aylor, L.J. Lobree, J.A. Reimer, and A.T. Bell

66 1

Precious Metal Loaded In/H-ZSM-5for Reduction of Nitric Oxide with Methane in the Presence of Water Vapor M. Ogura and E. Kikuchi

67 1

Interfacial RhOJCeO, Sites as Locationsfor Low Temperature N,O Dissociation J. Cunningham, J.N. Hickey, R. Cataluna, J.-C. Conesa, J. Soria, and A. Martinez-Arias

68 1

The Activity of VOJZrO, for the Selective Catalytic Reduction of NO V . Indovina, M. Occhiuzzi, P. Ciambelli, D. Sannino, G. Ghiotti, and F. Prinetto

69 1

Selective Reduction of NO, by Propene over Au/y-Al,O, Catalysts M.C. Kung, J.-H. Lee, A. Chu-Kung, and H.H. Kung

70 1

x1

B-20

The Role of Surface-Generated Gas-Phase Methyl Radicals in the Reduction of NO by CH, over a Sr/La,O, Catalyst S. Xie, T.H. Ballinger, M.P. Rosynek, and J.H. Lunsford

71 1

Part B Catalysis on Zeolites and Microporous Solids B-2 1

B-22

B-23

B-24

B-25

B-26

Isomerization and Hydrocracking of Decane and Heptadecane on Cubic and Hexagonal Faujasite Zeolites and Their Intergrowth Structures E.J.P. Feijen, J.A. Martens, and P.A. Jacobs

72 1

Solid State Ion Exchange of Alkali Metal Cations into Zeolite Y: Physicochemical Characterization and Catalytic Tests J. Weitkamp, S. Emst, M. Hunger, T. Roser, S. Huber, U.A. Schubert, P. Thomasson, and H. Knozinger

73 1

Role of Bronsted and Lewis Acidity in the Conversion of n-Pentane on Dealuminated H-Y, H-Mordenite and HZSM-5 V . Gruver, Y . Hong, A.G. Panov, and J.J. Fripiat

74 1

The Montmorillonite Catalyzed Production of Phenol and Acetone from Cumene Hydroperoxide W .A. de Groot, E.L.J. Coenen, B.F.M. Kuster, and G.B. Marin

75 1

Activation of Alcohols and Ketones by Surface Hydroxyls of Strong Solid Acids L. Kubelkova, J . Kotrla, J. Floriin, T. Bolom, J. Fraissard, L. Heeribout, and C. Doremieux-Morin

761

Heterogeneous Catalystsfor the Direct, Halide-fiee Carbonylation of Methanol B. Ellis, M.J. Howard, R.W. Joyner, K.N. Reddy, M.B. Padley, and W.J. Smith

77 1

B-27

Activation of Ethane on Modified ZSM-5 Zeolites Studied under Transient Conditions 78 1 A. Hagen, O.P. Keipert, and F. Roessner

B-28

Imaging of n-Hexane in Zeolites by Positron Emission Profiling (PEP) R.A. van Santen, B.G. Anderson, R.H. Cunningham, A.V.G. Mangnus, J. van Grondelle, and L.J. van IJzendoorn

79 1

Ethylene Dimerization in Nickel Containing MCM-41 and AIMCM-41 Studied by Electron Spin Resonance and Gas Chromatography M. Hartmann, A. Poppl, and L. Kevan

801

Alkane Partial Oxidation with Iron N,N'-bis(2-Pyridinecarboxamide)Complexes Encaged in Zeolite Y P.P. bops-Gerrits, M. L'abbe, W.H. Leung, A.-M. Van Bavel, G. Langouche, I. Bruynseraede, and P.A. Jacobs

81 1

B-29

B-30

xii Catalyst Characterization B-3 1

B-32

B-33

B-34

B-35 B-36

B-37

B-3 8

B-39

B-40

B-4 1

B-42

Physicochemical Characterization and Catalytic Properties of Solid Superacids Based on Surfated Zirconia Modified with Supported Platinum L.M. Kustov, T.V. Vasina, A.V. Ivanov, O.V. Masloboishchikova, E.G. Khelkovskaya-Sergeeva, and P. Zeuthen

82 1

Brensted Acid Strength of Solids Studied by 'HNMR: Estabhhing the Scale; Infruence of Lewis Acid Sites L. Heeribout, V. Semmer, P. Batamack, C. Doremieux-Morin, R. Vincent, and J. Fraissard

83 1

The Development of Strong Acidity by Non-Framework Aluminium in H-USY Determined by A1 XAFS Spectroscopy D.C. Koningsberger and J.T. Miller

84 1

Mesoporous MCM-41 Aluminosilicates as Model Silica-Alumina Catalysts: Spectroscopic Characterization of the Acidity F. Di Renzo, B. Chiche, F. Fajula, S. Viale, and E. Gmone

85 1

Elucidating the Nature of the Cobalt Centres in CoAPO-I8 Acid Catalysts L. Marchese, G. Martra, N. Damilano, S. Coluccia, and J.M. Thomas

86 1

Highly Dispersed Titanium Oxide on Silica: Preparation, Characterizationby XAFS, and Photocatalysis S . Yoshida, S. Takenaka, T. Tanaka, H. Hirano, and H. Hayashi

87 I

An Advance in Raman Studies of Catalysts: UltravioletResonance Raman Spectroscopy C. Li and P.C. Stair

88 1

The Dynamics of Surface-Catalyzed Reactions Studied by Infrared Chemiluminescence of the CO and CO, Products K. Watanabe, H . Uetsuka, H. Ohnuma, and K. Kunimori

89 1

Electron Availability and the Surface Fermi Level Local Density of States. An Alternative Way to See CatalyticActivity of Metals Y.Y. Tong, A.J. Renouprez, G.A. Martin, and J.J. van der Klink

90 1

A New Characterization Method for Adsorbed Hydrogen on Supported Pt Particles K. Asakura, T. Kubota, N. Ichikuni, and Y. Iwasawa

91 1

Hydrogen Chemisorption and Mobility on RdSiO,, WRdSiO, and Ru-Ag/SiO, R.L. Narayan, N . Savargaonkar, M. Pruski, and T.S. King

92 1

The Interaction of Rh with Intrinsic Defects of Reduced RwCeO, Catalysts: A Comparative XPS/UPS and 'H-NMR Study A. Pfau, J. Sanz, K.D. Schierbaum, W. Gopel, J.P. Belzunegui, and J.M. Rojo

93 1

...

XU1

B-43

In Situ Characterizationof the VanadiumSilicalite Catalyst (73-2) and its Photocatalytic Reactivity M. Anpo, S.G. Zhang, and H. Yamashita

94 1

Catalyst Deactivation B-44

Promotion and Deactivation of V,O,/TiO, SCR Catalysts by SO, at Low Temperature W.S. Kijlstra, N.J. Komen, A. Andreini, E.K. Poels, and A. Bliek 95 1

B-45

Structural Aspects of Activation and Deactivation of Cobalt Catalysts in Hydrogenation of Carbon Dioxide B. Klingenberg, G. Frohlich, F. Grellner, U. Kestel, G. Meyer, M. VoB, D. Bor,gnann, G. Wedler, J. Lojewska, T. Lojewski, and R. Dziembaj

96 1

Nature of H-Mordenite Deactivation Phenomena During SCR of No, M . Lezcano, A. Ribotta, E. Miro, E. Lombardo, J. Petunchi, C. Moreaux, and J.M. Dereppe

97 1

B-46

Oxidation Catalysis

c-1 c-2

c-3

c-4

c-5

C-6

c-I C-8

Eflects of Cs and V on Heteropolyacid Catalysts in Methacrolein Oxidation L.M. DeuBer, J.W. Gaube, F.-G. Martin, and H. Hibst

98 1

The Oxidation of C, Molecules on VanadylPyrophosphate Catalysts V.V. Guliants, J.B. Benziger, and S . Sundaresan

99 1

Direct Oxidation of Isobutane into Methacrylic Acid over Cs, Ni, and V-substitutedH,PMo,,O,, Heteropoly Compounds N . Mizuno, W. Han, T. Kudo, and M. Iwamoto

1001

Role of Water on the Properties of FePO Catalysts used for the Oxidative Dehydrogenation of Isobutyric Acid J.M.M. Millet, M. Forissier, D. Rouzies, P. Bonnet, and J.C. Vedrine

1011

Gas-Phase 0, Oxidation of Alkylaromatics with Ag'-Doped CVD Fe/Mo/DBH J.S. Yo0 and C. Choi-Feng

1021

Mechanistic Approach of the Oxidative Dehydrogenation of Propane over VMgO Catalysts by in situ Spectroscopic and Kinetic Techniques A. Pantazidis and C. Mirodatos

1029

Mechanochemistry in Preparation and Modification of Vanadium Catalysts V.A. Zazhigalov, J. Haber, J. Stoch, L.V. Bogutskaya, and I.V. Bacherikova

1039

Oxidative Dehydrogenation of Propane on Rare Earth Vanadates. Influence of the Presence of CO, in the Feed B. Zhaorigetu, R. Kieffer, and J.-P. Hindermann

1049

xiv C-9

C- 10

Structure Sensitivity of Oscillating Partial Oxidation of Propene on Widely Dispersed Platinum Catalysts M. Kobayashi, T. Kanno, H. Takeda, and S. Fujisak

1059

Catalytic Oxidation of Propane to Aclylic Acid with Molecular Oxygen Activated over Reduced Heteropolymolybdates W. Ueda, Y . Suzuki, W. Lee, and S. Imaoka

1065

Catalysis on Metals C- 1 1

C- 12

C- 13

Ru, Pt and Co Clusters in Zeolite Micropores; EX4FS/FTIWTPD Characterizationand Catalytic Behaviors in Methane Homologation M. Ichikawa, T . Tanaka, W. Pan, T. Ohtani, R. Ohnishi, and T. Shido

1075

The Relation between Pre-treatment of Promoted Copper Catalysts and their Activity in Hydrogenation Reactions D.S. Brands, E.K. Poels, and A. Bliek

1085

Cobalt Intermetallic Compoundsfor Selective Hydrogenation of Acetylene T . Komatsu, M. Fukui, and T. Yashima

1095

C- 14

Supported Pd-Cu Catalysts Preparedfrom Bimetallic Organo-metallicComplexes: Relation Between Surface CompositionMeasured by Ion Scattering and Reactivity A.J. Renouprez, K. Lebas, G Bergeret, J.L. Rousset, and P. Delichkre 1105

C- 15

Synergetic Effect of Pd and Ag Dispersed on MgO in the Reduction of NO by H2 at Room Temperature S. Naito and Y. Tanaka

1115

Heterogeneous Palladium Catalystsfor the Oxidation of Propylene to Propylene Glycol Acetates: Effect of Platinum and Rhodium as Promoters E.V. Gusevskaya, V.A. Likholobov, A.V. Karandin, A.I. Boronin, and E.M. Moroz

1125

Probing the Limits of Strucutre Insensitivity: Size-dependent Catalytic Activity of Al,O,-supported Iridium Clusters and Particlesfor Toluene Hydrogenation F.-S. Xiao, W.A. Weber, 0.Alexeev, and B.C. Gates

1135

The Kinetic Isotope Effectfor Alkane Dehydrocyclization B. Shi and B.H. Davis

1145

High-Thiotolerant Pt-Ge/Al,O, Naphtha Reforming Catalysts by in-situ Alloying T.F. Garetto, A. Borgna, and C.A. Apesteguia

1155

Metal-Support Interactions in Supported Platinum Catalysts: Zeolites and Amorphous Supports B.L. Mojet, M.J. Kappers, J.T. Miller, and D.C. Koningsberger

1165

The Role of Gaseous and Surface Isocyanates in the Reaction of Mixtures of H,, NO and CO over Supported Platinum Catalysts R. Diimpelmann, N.W. Cant, and A.D. Cowan

1175

C- 16

C-I7

C- 1 8

C- 19

C-20

C-2 1

xv

c-22

C-23

Reactions of Isobutane and Isobutylene over Silica- and L-Zeolite-Supported Pt/Sn and Pt/Sn/K Catalysts R.D. Cortright, E. Bergene, P. Levin, M. Natal-Santiago, and J.A. Dumesic

1185

Effect of Tungsten on Supported Platinum Catalysts J.L. Contreras and G.A. Fuentes

1195

General Papers C-24

C-25

C-26

On the Role of Microstructure of VanadiumPhosphorus Oxidesfor Propane Oxidation to Acrylic Acid N . Gribot-Perrin, J.-C. Volta, A. Burrows, C. Kiely, and M. Gubelmann-Bonneau

1205

Hydroconversion of Heavy Crude Oils Using Soluble Metallic Compounds in the Presence of Hydrogen or Methane A. Morales, A. Salazar, C. Ovalles, and E. Filgueiras

1215

Hydrodenitrogenation of Indole over NiMo Sulfide Catalysts L. Zhang and U.S. Ozkan

1223

Application of Theoretical Methods in Catalysis C-27

C-28

C-29

C-30

Quantum-chemicalStudy of the Nonclassical Carbonium Ion-like Transition States in Isobutane Cracking on Zeolites V.B. Kazansky, M.V. Frash, and R.A. van Santen

1233

Chemisw of Su1f.r Oxides on TransitionMetal Surfaces: A Bond Order Conservation - Morse Potential Modeling Perspective H . Sellers and E. Shustorovich

1243

Chemisorption and Decomposition of C, and C, Hydrocarbons on a Pd(l1 I ) Surface: A Periodic Density Functional Study J.-F. Paul and P. Sautet

1253

Density Functional TheoT Studies of Zeolite Acidity and Reactivity J.B. Nicholas

1263

Catalyst Synthesis

c-3 1 Control of Bulk and Surface Compositionof Doped Sm,Sn,O- Pyrochlore.

C-32

Relation between Formation of 0-Ba-CI Grajiings and C,-Selectivity in the Oxidative Coupling of Methane A,-C. Roger, C. Petit, S. Libs, and A. Kiennemann

1273

The Direct Room-Temperature Synthesis of Ce0,-based Solid Solutions: A Novel Route to Catalysts with a High oxygen Storage/Transport Cupacity F. Zamar, A. Trovarelli, C. de Leitenburg, and G. Dolcetti

1283

XVi

c-33

c-34

c-3s

C-36

c-37 C-3 8

c-39 C-40

Studies on Supported Metal Oxide-OxideSupport Interactions (An Incorporation Model) Y . Chen, L. Dong, Y.S. Jin, B. Xu, and W. Ji

1293

Zeolite-ConfinedMn Complexes of Cyclic Amines: New Selective Catalysts for Hydrocarbon Oxidation D.E. De Vos, J. Meinershagen, and T. Bein

1303

Platinum Indium Bimetallic in Silicalite: Preparation, Characterizationand Use in the VinylcyclohexeneTransformation P. MCriaudeau, A. Thangaraj, and C. Naccache

1313

Promotion of y-Alumina Dissolution by Metal Ions During Impregnation. Thermal Stability of the Formed Coprecipitates J.-B. d’Espinose de la Caillerie and 0. Clause

1321

Supported Catalysts Based on Carbon Fibrils M.S. Hoogenraad, M.F. Onwezen, A.J. van Dillen, and J.W. Geus

1331

Molecular Sieve Ti-UTD-I:A Novel Oxidation Catalyst K.J. Balkus, Jr., A. Khanmamedova, A.G. Gabrielov, and S.I. Zones

1341

Metal Composite Membranes: Synthesis, Characterization and Reaction Studies K.L. Yeung, R. Aravind, J. Szegner, and A. Varma

1349

Influence of the Process Parameters on the Extrusion of Ceramic Catalysts D. Ballardini, L. Sighicelli, C. Orsenigo, L. Visconti, E. Tronconi, P. Forzatti, A. Bahamonde, E. Atanes, J.P. Gomez Martin, and F. Bregani

1359

Synthesis Gas Conversion

c-4 1

C-42

c-43 c-44

c-45

Study on Catalytic Synthesis of Methanolfrom Syngas via Methylformate in Heterogeneous “One-Pot” Reactions H. Zhang, H . Li, G. Lin, Y. Liu, and K.R.Tsai

1369

The Effect of Zinc Oxide in Raney Copper Catalysts on Methanol Synthesis, Water Gas Sh$, and Methanol Steam Reforming Reaction D. Wang, L. Ma, C.J. Jiang, D.L. Trimm, M.S. Wainwright, and D.H. Kim

1379

Model Studies of Methanol Synthesis on Copper Catalysts J. Nakamura, I . Nakamura, T. Uchijima, T. Watanabe, and T. Fujitani

1389

Key Reactionfor Formation of Isobutene over ZrO, and Isoprene over CeO, in CO Hydrogenation K. Maruya, M . Hara, J. Kondo, K.Domen, and T. Onishi

1401

Promotion Effects in Methanol Synthesis over MgO-SupportedFe-Ir Catalysts Preparedfrom Mixed-Metal Clusters S . Marengo, R. Psaro, C. Dossi, S. Calmotti, and R.Della Pergola

1411

xvii C-46

Nanoscale Atfrition During Activation of Precipitated Iron Fischer-Tropsch Catalysts: Implicationsfor Catalyst Design A.K. Datye, M.D. Shroff, Y. Jin, R.P. Brooks, J.A. Wilder, M.S. Harrington, A.G. Sault, and N.B. Jackson

1421

Author Index

143 1

Subject Index

1435

This Page Intentionally Left Blank

xix

PREFACE

It seems highly appropriate that the Eleventh International Congress on Catalysis be held in Baltimore, USA, less than 200 km from the birthplace of these quadrennial events that began in Philadelphia in 1956. Planning for this 40th Anniversary Meeting has been coordinated by Gary L. Hailer, with the support of the Organizing Committee comprised of John N. Armor, Alexis T. Bell, W. Curtis Conner, Jr., Dady B. Dadyburjor, W. Nicholas Delgass, Sergio Fuentes, Richard D. Gonzalez, W. Keith Hall, Joe W. Hightower, Enrique Iglesia, Leo E. Manzer, James Maselli, Daniel E. Resasco, Kathleen C. Taylor, M. Albert Vannice, and Bohdan Wojciechowski. The PROCEEDINGS contain 145 papers - 7 plenary lectures and 138 submitted papers selected for oral presentation. The plenary lectures include five overviews of vital research areas by highly respected researchers and two overviews of advances in the science and technology of catalysis made during the last 40 years. The first group explores the forces that drive innovation in catalysis, constrained geometry in metallocene olefin polymerization, characterization and design of oxide surfaces, photocatalysis, and factors required in the molecular design of catalysts. Two others are presented by researchers who attended the first ICC meeting 40 years ago and who have been substantive contributors to science and engineering developments that have occurred since then. The 138 submitted papers were selected in the following manner. From a total of 521 submitted two-page abstracts, 156 were identified by peer review and evaluation of the Program Committee to be expanded into 10-page (maximum) camera-ready manuscripts. Submitted manuscripts were then peer-reviewed by at least two experts in the field according to standards comparable to those used for archival journals. Diversity in country origin was also considered, and an attempt was made to minimize multiple publications for individual research groups. Consequently, the 138 papers included herein should be considered as peer-reviewed publications that represent the worldwide state-of-the-art in catalysis research. These PROCEEDINGS of the International Congress on Catalysis differ from those published previously in two important ways. First, the papers were published PRIOR to the meeting for distribution to all delegates who attended the meeting in Baltimore. Second, none of the discussion at the meeting is included. With publication costs skyrocketing, we have elected to abandon the tradition of including the discussion, realizing that in so doing a valuable part of the meeting will be lost forever to posterity. However, it does have the advantage of smaller size (only two volumes, < 1,600 pages) which should make the books more attractive to libraries and other repositories of research literature.

XX

Finally, I would like to thank my co-editor, Nick Delgass, and Gary Heller for assistance in processing the paper revisions and to specially acknowledge Alex Bell and Enrique Iglesia, chair and co-chair of the Program Committee, for their excellent and timely management of the difficult but crucial task of paper selection. I am also grateful to Drs. Huub Manten of Elsevier Science for the fantastic cooperation the company provided in getting these two volumes printed. Most important, I wish to thank all the authors of the 145 papers appearing in these PROCEEDINGS for their diligence in faithfully meeting the exceedingly short deadlines that were necessary to get the material into print prior to the meeting. Joe W. Hightower, Editor

Houston, TX (USA)

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

DRIVING FORCES FOR I N N O V A T I O N IN APPLIED C A T A L Y S I S lan E. Maxwell Koninklijke/ShelI-Laboratorium, Amsterdam (Shell Research B.V.), P.O. Box 38000, 1030 BN, Amsterdam, Netherlands 1. INTRODUCTION

Recent governmental sponsored studies in both the US and Europe have recognized the vital role of catalytic technologies for sustainable economic growth in the future. For example, it has been estimated that in developed countries catalysis contributes directly and indirectly through processes and products to some 20-30% of GDP (Gross National Product). Furthermore, catalytic environmental technologies such as automobile exhaust catalysts and the selective catalytic reduction (SCR) DeNOx systems in power plants have already significantly contributed to the reduction of environmentally harmful emissions into the lower atmosphere. In addition, these studies have identified catalysis as not only being pervasive but also offering significant scope for further innovative development of new and improved technologies for environmentally acceptable processes and products in the future. The spectrum of process industries which are directly impacted by catalysis include for example, oil refining, natural gas conversion, petrochemicals, fine chemicals and pharmaceuticals. Environmental catalytic technologies also play an important role in emission control systems for power generation, fossil fuel driven transportation, oil refining and chemical industries. Catalytic technologies typically embrace a wide range of disciplines such as heterogeneous and homogeneous catalysis, materials science, process technology, reactor engineering, separation technology, surface science, computational chemistry and analytical chemistry (Figure 1). Innovation in this field is therefore very often achieved by lateral thinking across these different disciplines. This presentation will attempt to develop this theme further by means of examples from recent commercial successes and from this platform provide some guidelines for multi-disciplinary approaches at the academic and industrial interface to further enhanced innovation in catalytic technologies in the future. 2. OIL REFINING AND NATURAL GAS CONVERSION

The discipline of materials science, for example, has a major impact on innovation in catalysis. Developments in the field of porous solids have led to some

tal

Figure 1. Disciplines of Prime Importance to Catalytic Technologies new catalytic processes in the refining area based on novel shape selective microporous materials. Two such new processes which have recently been commercialized based on these types of materials include the selective isomerization of nbutene to iso-butene [1] (MTBE precursor) and iso-dewaxing of lubeoils [2]. Interestingly, both groups of industrial researchers (Shell and Chevron groups, respectively) involved in these developments combined the disciplines of computational chemistry, materials science andheterogeneous catalysis to gain an in-depth understanding of the relationships between the detailed topology of the micro-porous materials and the shape selective catalytic performance. In the case of n-butene isomerization it was demonstrated (Figure 2) that the ideal micro-pore topology led to retardation of the C8 dimer intermediate and that the catalyst based on the ferrierite structure was close to optimal in this respect [ 1 ]. For selective isodewaxing a one-dimensional pore structure which constrained the skeletal isomerization transition state and thereby minimized multiple branching such as the SAPO-1 1 structure was found to meet these criteria. Clearly, these are ideal systems in which to apply computational chemistry where the reactant and product molecules are relatively simple and the micro-porous structures are ordered and known in detail. Another recent new application of a microporous materials in oil refining is the use of zeolite beta as a solid acid system for paraffin alkylation [3]. This zeolite based catalyst, which is operated in a slurry phase reactor, also contains small amounts of Pt or Pd to facilitate catalyst regeneration. Although promising, this novel solid acid catalyst system, has not as yet been applied commercially.

80 O

E

I

70I

\\

FER

60

(3

13) t-

5O

TON

40

MFI

.>_ 30 .1-, m (~

cr

20F 10

~ 1

:

-

2

3

MOR

= 4

5

Position TMP vs smallest zeolite ring [Angstrom] Figure 2. Comparison of Calculated Diffusional barriers for 2,4,4-trimethyl-3pentane (TMP) in Various Zeolites and Molecular Sieves A non-acidic isomerization catalyst system has unexpectedly emerged from recent studies by French workers [4] in the area of Mo-oxycarbides. Although at an early stage of development, these new materials exhibit high selectivities for the isomerization of paraffins such as n-heptane. An alternative non-carbenium ion mechanistic route to achieve isomerization of higher alkanes could potentially overcome some of the limitations of conventional solid acid based catalyst systems. Novel combinations of heterogeneous catalysis, reactor technology and separation technologies have also led to major innovations. Examples include catalytic distillation which is now widely applied for the manufacture of MTBE with other potential applications under development [5]. Another example of multidisciplinary synergy in this context which was recently commercialised is the socalled Synsat process [6] developed jointly by the Criterion and Lummus companies for enhanced deep hydrogenation and desulphurization of diesel fuels. This process employs a multiple catalyst bed system in a single reactor shell with intermediate by-product gas removal and optional counter-current gas/liquid flow in the bottom catalyst bed (Figure 3). Government legislation related to aromatics and sulphur contents of diesel fuels has become more stringent and global in recent years such that the development of improved catalytic hydrotreating processes is most timely. Catalytic membranes, which combine the disciplines of heterogeneous catalysis, separation technology, materials science and reactor engineering, which have for some time held considerable promise now appear to be gradually emerging as viable technologies. Promising potential applications include propane to aromatics [7] and catalytic oxidation of methane to synthesis gas using air as the oxidant [8]. In the former example, Japanese workers [7] applied a Pd-alloy membrane reactor (PMR)

Fresh feed ...... =i Make-up H2 ~

CatalystA

Recycled ~

CatalystB

liquid

CatalystC

~ ~

Reactor tc p / 1

Vaporto vapor/liquid separation/ recycle

Make-up H 2 - ~ J

ppH2

Reactor bottom

Dieselproduct Figure 3. The Synsat Process for Deep Hydrotreating of Diesel Fuels 85 8

9 o >

-

7s

J

65

m'~

55 ....... 20

I

40

PMR,~,,'4 ,,'" s// sss

.................

\~" I

60

I

80

1O0

Conversion of Propane[%] Cat. Ga-H-ZSM5

773K

Figure 4. Comparison of Propane Aromatization Performances of a Palladium Membrane Reactor (PMR) and a Conventional Reactor (CR) using a Ga-H-ZSM-5 Catalyst to shift the equilibrium for the dehydroaromatization of propane. This PMR system resulted in a significant improvement in the selectivity towards the desired mixed aromatic products (Figure 4). For the partial oxidation of methane a catalytic membrane system under development by researchers at the Argonne National laboratory [8] effectively

separates oxygen from air which is then passed through the membrane in an anionic form to react with methane in the presence of catalyst. High conversion and selectivity levels to synthesis gas have been claimed, although space velocities have not yet been published. Such new catalytic processes can potentially reduce the costs of synthesis gas production and therefore positively impact the overall economics of natural gas conversion technology. New developments in the field of ceramic foam monoliths could also potentially provide new catalytic process technology for the conversion of methane into synthesis gas. For example, workers at Minnesota University [9] have achieved high synthesis gas yields at both high temperatures and space velocities using rhodium supported on a ceramic foam. 3. CHEMICALS

New materials are also finding application in the area of catalysis related to the Chemicals industry. For example, microporous [10] materials which have titanium incorporated into the framework structure (e.g. so-called TS-1) show selective oxidation behaviour with aqueous hydrogen peroxide as oxidizing agent (Figure 5). Two processes based on these new catalytic materials have now been developed and commercialized by ENI. These include the selective oxidation of phenol to catechol and hydroquinone and the ammoxidation of cyclohexanone to ecaprolactam. It was soon recognized that the TS-1 system has limitations particularly due to the small pore system which imposed restrictions on the molecular size of the OH

OH

ArH

,

ArOH

OH R ArOH

R,

R ,~:::0 + R'

R

H OH

HON,~

Figure 5. Range of Selective Oxidation Reactions Catalyze by the TS-1 Zeolite System Using Aqueous H202 as Oxidizing Agent

reactant molecules. More recently therefore titanium has been incorporated into larger pore zeolites [11] (e.g. beta and ZSM-48) and even mesoporous structures such as MCM-41 [12]. These larger pore materials also enable more bulky molecules such organic hydroperoxides to be used as oxidising agents. This field is still growing rapidly and would appear to hold promise for the development of new and improved heterogeneous catalyst systems for selective oxidation reactions of value to the chemicals industry. Base catalysis is another area which has received a recent stimulus from developments in materials science and microporous solids in particular. The Merk company, for example, has developed a basic catalyst by supporting clusters of cesium oxide in a zeolite matrix [13] . This catalyst system has been developed to manufacture 4-methylthiazole from acetone and methylamine. Heteropolyacids are also beginning to emerge from academic laboratories and find commercial applications. Showa Denko, for example, claim to have a process [14] for the direct oxidation of ethylene to acetic acid employing a bifunctional Pt/heteropolyacid catalyst system. The potential synergies between the disciplines of homogeneous and heterogeneous catalysis have also long been recognized but progress in the past has generally been frustrated by intangible technical problems. Particularly challenging is the goal of immobilizing homogeneous catalyst systems onto solid supports without incurring catalyst loss by leaching under reaction conditions. A particularly elegant approach to this problem involves the immobilization of a metal complex in a thin film of polar solvent (e.g. water) which is adsorbed on a high surface area hydrophilic support (e.g. silica). Such a system has been successfully applied in the laboratory [15] to immobilize a homogeneous water soluble chiral hydrogenation catalyst based on ruthenium (Figure 6). Using this catalyst a high degree of enantioselectivity was achieved for an important hydrogenation step in the synthesis of (S)-naproxen (an anti-inflammatory drug). 4. EMISSION CONTROL

Monolithic structures, often based on ceramic materials, are increasingly being applied in catalysis. The initial major thrust of monoliths was in the area of automobile exhaust catalyst systems where they are now applied exclusively. The demand for improved performance of these emission control systems, particularly under high temperature conditions is driving new developments such as ceramic foam technologies. Catalytic combustion, particularly for application in gas turbines, is another emerging field of technology where the developments in monolithic structures will be of growing importance. The Osaka Gas and Kobe Steel companies [16] have jointly developed a catalytic monolith which operates up to 1300 ~ and has been tested in a 160 kW gas turbine. New ceramic materials based on Mnsubstituted hexa-aluminates provide the high temperature stability required of catalytic monoliths for these demanding applications.

Ar2 CI P~ / P/Ru.~ Ar2

CI

" • •

Ru-BINAP(SO3Na), J = m - NaO~SC,~'-

-

H2 MeO

~ MeO

CO2H (S) - naproxen

Figure 6. Immobilization of Chiral Ruthenium Hydrogenation Catalyst in a Thin Hydrophilic Film on a Porous Glass Support Examples of multi-disciplinary innovation can also be found in the field of environmental catalysis such as a newly developed catalyst system for exhaust emission control in lean burn automobiles. Japanese workers [ 1 7] have successfully merged the disciplines of catalysis, adsorption and process control to develop a socalled NOx-Storage-Reduction (NSR) lean burn emission control system. This NSR catalyst employs barium oxide as an adsorbent which stores NOx as a nitrate under lean burn conditions. The adsorbent is regenerated in a very short fuel rich cycle during which the released NOx is reduced to nitrogen over a conventional three-way catalyst. A process control system ensures for the correct cycle times and minimizes the effect on motor performance. 5. FUTURE CHALLENGES

The above examples should serve to reinforce the multi-disciplinarity of catalytic technologies. However, to further exploit the significant potential of catalysis for innovation and renewal across a broad range of industries multi-discplinary approaches to problem solving will be vital to success. This ingredient for success is, in general, recognized within industrial research laboratories where multidisciplinary project teams are commonly deployed. However, this approach is traditionally less common in academic laboratories which generally tend to be more narrowly focussed in terms of disciplinary skills. Another element of concern at this interface is the perceived gap between academic basic research and industrial applied research. The recent trend towards shorter term goals within industrial research laboratories has further exacerbated

()

.~

oo

Sectors

Enabling Technologies

Multi-sector

Emerging Technologies

Multi-sector

Figure 7. Programme Model for Proposed UK National Institute of Applied Catalysis this situation whereby discontinuities are perceived to exist between basic and applied catalysis. Both these factors, which are likely retarding innovation and the potential synergies between industrial and academic laboratories in the field of catalysis, have been termed the "innovation gap". In Europe, particularly in the UK and the Netherlands, this mismatch has been recognized and some government supported initiatives are currently in progress. In the Netherlands, for example, a type of "virtual" organization called NIOK has been established to foster multi-disciplinary inter-university linkages and to strengthen the relationships with industry. More recently the UK has also launched a similar initiative with the intention of forming an organization based on both "virtual" and "hard core" components involving both academia and industry. This proposed new UK organization is termed NIAK (National Institute of Applied Catalysis). This NIAK organizational model is directly aimed at closing what is perceived to be a substantial "innovation gap" between industry and academia in the UK. A programme model has been recently developed for NIAC (Figure 7) which contains three separate components defined as emerging, enabling and sectors. The emerging and enabling components are envisaged to contain elements of common interest to all the industrial members whereas the sectorial programmes will be much more specifically oriented towards the individual needs of each industrial sector. Thus, in order to fully realize the potential of catalytic technologies not only will this require technical innovation in multi-disciplinary teams but also the appropriate organizational structures which maximize the synergies between academic and industrial research. The countries which recognize this potential and provide the

appropriate stimuli will likely have leading positions in catalytic technologies in the future. ACKNOWLEDGEMENTS

The author is most grateful for the valuable discussions with many of his Shell research colleagues in the field of catalytic technologies. In particular, as relate to this paper, discussions with Dr. G. Boxhoorn, Dr. K. de Jong, Ir. J. Naber and Dr. W. Stork are gratefully acknowledged. REFERENCES

1. H.H. Mooiweer, K.P. de Jong, B. Kraushaar-Czarnetzki, W.H.J. Stork and B.C.H. Krutzen, "Zeolites nd Related Microporous Materials: State of The Art 1994", Studies in Surface Science and Catalysis, Elsevier, 1994, Eds. J. Weitkamp et al, p2327 2. S.J. Miller, Microporous Mater., 1994, 2, 239-449 3. CM.A.M. Mesters, D. Peferoen, J.P. Gilson, C. de Groot, P.T.M. van Brugge and K.P. de Jong, submitted for publication to Appied Catalysis 4. M. Ledoux, C. Pham-Huu, H. Dunlop and J. Guille, "Proceedings lOth International Congress on Catalysis, eds. G.L. Solymosi and P. Tetenyi, Akademiai Kiado, Budapest, 1993, B, 955-967 5. V.J. D'Amico, P.H.O. Dixon & B.A. Strain, Paper AM-89-44, NPRA meeting, 1989, San Francisco. 6. A.J. Suchanek, E. L. Granniss, Presentation, AM-95-40, NPRA, Annual Meeting, March 19-21, 1995, San Francisco 7. Shokubai (Catalyst), 36 (4), 1994, 246 8. Argonne Research Laboratory Report, 1995 to be published. 9. D.A. Hickman and L.D. Schmidt, Science, 1993, 259, p343 10. B. Notari, Stud. Sur. Sci. Catal., 1988, 37, 413-425 11. C.B. Khouw, C.B. Dartt, J.A. Labinger, M.E. Davis, J. Catal., 1994, 149, p195-205 12. J.M. Thomas, Nature, Ti-MCM-41 13. F.P. Gortsema, B. Beshty, J.J. Friedman, D. Matsumoto, J.J. Sharkey, G. Wildman, T.J. Blacklock, S.H. Pan, Presentation at the 14th Conference on Catalysis of Organic Reactions, Albuquerque, April 27-28, 1992 14. M. Otake, Chemtech 1995, September, p36-41 15. K. Wan and M.E. Davis, Nature, 370, (1994), 449 16. H. Sadamori, T. Tanioka and T. Matsuhisa, "Proceedings of the International Workshop on Catalytic Combustion", Ed., H. Arai, Catalysis Society of Japan, (1994), p 158 17. S. Matsumoto, Pre-prints 2nd Japn-EC Joint Workshop on the Frontiers of Catalytic Science and Technology, 1995, vol 1, p39-42

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J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

11

C o n s t r a i n e d G e o m e t r y and O t h e r Single Site Metallocene Polyolefin Catalysts: A Revolution In Olefin P o l y m e r i z a t i o n James C. Stevens Polyolefins and Elastomers Research and Development Laboratories, The Dow Chemical Company, 2301 Brazosport Boulevard, Freeport, TX 77541 1. ABSTRACT The polyolefins industry is at a crossroads. A new generation of single-site catalyst technology promises to revolutionize this multi-billion pound per year industry. Single-site catalysts have recently moved from laboratory curiosities to commercial success. Newly developed single-site c a t a l y s t s allow u n p r e c e d e n t e d control of polymer molecular a r c h i t e c t u r e , which yields products having improved properties. The consistency and control of polymer structure is allowing new discoveries to be made in f u n d a m e n t a l polymer research. This paper will touch on all aspects of single-site catalyst technology, including m e t a l l o c e n e m e t a l complexes, a c t i v a t i n g c o c a t a l y s t s (e.g., alumoxanes), cationic catalysts, as well as single site polymers, focusing on recent developments at Dow Plastics. 2. INTRODUCTION The metal catalyzed production of polyolefins such as high density polyethylene (HDPE), linear low d e n s i t y polyethylene (LLDPE) a n d polypropylene (PP) has grown into an enormous industry. Heterogeneous transition metal catalysts are used for the vast majority of PE and all of the PP production. These catalysts fall generally within two broad classes. Most commercial PP is isotactic and is produced with a catalyst based on a combination of titanium chloride and alkylaluminum chlorides. 1 HDPE and LLDPE are produced with either a titanium catalyst or one based on chromium supported on silica. 2 Most commercial t i t a n i u m - b a s e d PE catalysts are supported on MgC12. One of the most exciting developments in the polyolefins industry in recent years has centered on the development of commercial homogeneous single-site catalysts. These single-site catalysts produce olefin polymers with properties that are different when compared with traditional thermoplastic polyolefins. Homogeneous single-site catalysts based on bis-cyclopentadienyl derivatives of titanium have been known since the 1950's, although the catalytic activity of these early catalysts was too low for commercial practicality. 3 The key discovery by K a m i n s k y and Sinn t h a t m e t h y l a l u m i n o x a n e (MAO, [MeA10]n ) i n conjunction with Cp2TiMe2 and Cp2ZrC12 afforded extremely active catalysts for

12 PE and atactic PP lead to the recent explosion of interest in single-site catalysts.4, 5 The most valuable feature of single-site catalysts is the ability to logically control the structure of the polymer from the design of the catalyst. The most commonly used families of single-site catalysts are based on metal complexes shown in Figure 1. The polymer that is produced with these catalysts is strongly influenced by the catalyst structure. Catalysts with structure 1, having C2v symmetry produce atactic PP. By adding a bridging group between the cyclopentadiene ligands, catalysts having chiral C2 symmetry such as 3 produce isotactic pp.6,7 Linking the cyclopentadienes together to give a catalyst having Cs symmetry as in 4 produces syndiotactic PP. These relationships have led to the development of catalysts which produce isotactic-b-atactic PP by rotation of a substituted indenyl catalyst through C2v and C2 symmetry.8

R

2

~MX2 ~'~R

TiX3

HDPE, LLDPE Atactic PP

Syndiotactic PS

Isotactic PP

1

2

3

MX 2

i

Syndiotactic PP 4

Isotactic-b-atactic PP 5

F i g u r e 1. Single-site catalysts for olefin polymerizations. The bis-cyclopentadienyl-based, or metallocene, single-site catalysts are generally activated with MAO in relatively large molar amounts. The catalytic activity increases with increasing A1 : M ratio. Typically, at least 500 - 1000 molar equivalents of aluminum are required for acceptable activity. The high levels of MAO are a problem commercially, due to the relatively high cost of MAO. In addition, very high levels of MAO leave a large amount of aluminumcontaining "ash" in the polymer which can affect the product properties. The

13 catalytic activity is also a function of the transition metal. In general, the order of activity is Zr > Hf > Ti. Polymer molecular weight is also a function of the transition metal, generally following the order Hf > Ti > Zr. The ability to control the polymer from the design of the catalyst, coupled with high catalytic efficiency has led to an explosion of commercial and academic interest in these catalysts. Exxon started up a 30 million lb/yr ethylene copolymer demonstration plant in 1991 using a bis-cyclopentadienyl zirconium catalyst of structure 1. The Dow Chemical Company (Dow) began operating a 125 million lb/yr ethylene/1-octene copolymer plant in 1993 and has since expanded production capacity to 375 million lb/yr. This paper will focus on the structure / property relationships of the catalysts used by Dow to produce single-site ethylene a-olefin copolymers.

3. C o m m ~

Polyethylene

Commercial polyethylene falls within three general classes, as shown in Figure 2. LDPE is a highly branched dendritic polymer containing a range of short and long-chain b r a n c h e s , which r e s u l t from v a r i o u s r a d i c a l recombination processes in a high temperature (250 - 300 ~ high pressure process of up to about 45,000 psi. The numerous long-chain branches impart high melt strength and excellent processability to the polymer. In contrast, linear HDPE and LLDPE are produced using coordination catalysts and are characterized by a linear backbone containing no long-chain branching. LLDPE is primarily produced as a copolymer of ethylene with C4-C8 a-olefins up to about 15 weight percent. As a result, the linear molecules impart good toughness and strength properties, but are relatively more difficult to process than LDPE. Processability can be improved by broadening the molecular weight distribution, which tends to increase the number of low molecular weight molecules. The increased processability is generally achieved at the expense of strength and other physical properties due to the reduced number of high molecular weight molecules. Narrow molecular weight linear polyethylenes are relatively difficult to process. Focusing on LLDPE, traditional chromium or titanium-based Ziegler/Natta catalysts produce a product with a broad distribution of individual polymer molecules, each of which contributes to the overall properties of the resin. It is believed that the active catalysts contain active sites of various oxidation states and coordination environments, each of which exhibits different rates of propagation, termination, and comonomer reactivity. 9 The catalyst sites which incorporate comonomer have higher rates of chain t e r m i n a t i o n , and consequently the polymer molecules which contain more comonomer are lower molecular weight. A significant fraction of the polymer contains little if any comonomer and is generally of high molecular weight. The "mixture" of polymer molecules which results represents a limitation of conventional heterogeneous polyolefin catalysts in that the ability to control the individual component polymer molecules is rather restricted.

14

radicals H2C---CH 2 high temperature

I~PE

Z/N Catalysts H2C---CH 2

n ttDPE

H2C---CH 2

R Z/N Catalysts

+

LLDPE F i g u r e 2. Classes of commercial polyethylene. Metallocene catalyzed LLDPE is characterized by polymer molecules t h a t are the result of a single active catalytic site. As a result, all of the polymer molecules can in theory be made with statistically the same comonomer distribution. The molecular weight distribution is quite narrow, in the range of about 2.0 Mw/Mn. As a result of the ability to tailor the individual polymer chains, material scientists now have the ability to produce targeted polymer species and control the properties of the resin to a high degree. Great advances in such areas as product strength, clarity, toughness, and melting behavior can be commercially realized using metallocene catalysts. Single site polymers having a uniform comonomer and molecular weight distribution produced with low efficiency v a n a d i u m catalysts have been commercially available from Mitsui (Tafmer resins) since the 1980's. However, these polymers are relatively expensive and the ability to tailor the product is limited.

15 4. Constrained Geometry Catalysts Dow has developed a new family of ethylene-based polyolefins using constrained geometry catalyst technology. The catalyst and process technology has been commercialized under the tradename INSITE TM. INSITE Technology utilizes a family of new constrained geometry catalysts t h a t allows the production of unique polyolefin polymers in a relatively low pressure solution process. An important feature of the solution process is the need to operate the polymerization reaction above the melting point of the polymer in order to keep the product in solution. Solution LLDPE processes generally run at very high ethylene conversion and with a short reactor residence time of only a few minutes. Unfortunately, metallocene catalysts such as structures 1-5 produce low molecular weight products under such conditions, due to relatively facile ~hydride elimination. Constrained geometry catalysts, on the other hand, allow for the production of high molecular weight ethylene copolymers in a high temperature solution process. The key catalyst features are shown in Structure 6. The catalysts are monocyclopentadienyl Group 4 complexes with a covalently attached amide donor ligand. The amide ligand stabilizes the metal electronically, while the short bridging group (B) has the effect of sterically opening up one side of the complex, producing a stable but highly open and reactive active site upon activation with a variety of cocatalysts.

Rn N.-MX2 ]

R M = Ti, Zr, H f B = SIR2, C2H4, etc. R = alkyl, aryl, etc. X = halide, CH3 S t r u c t u r e 6. General structure of Constrained Geometry Catalysts.

In general, the open nature of the catalytic site in the constrained geometry catalysts does not allow for much steric control of the polymerization reaction, and homopoly a-olefins are generally atactic, although it has been reported that a small degree of tacticity can be introduced in polypropylene using such catalysts by selection of the substituents and conducting the polymerization at r e l a t i v e l y low t e m p e r a t u r e s . 1~ The degree of tacticity obtained under commercially useful conditions is so low that the catalysts can be considered to be atactic. The s t e r i c a l l y u n e n c u m b e r e d c a t a l y s t active site allows the copolymerization of a wide variety of olefins with ethylene. Conventional heterogeneous Ziegler/Natta catalysts as well as most metallocene catalysts are much more reactive to ethylene than higher olefins. With constrained geometry catalysts, a-olefins such as propylene, butene, hexene, and octene are readily incorporated in large amounts. The kinetic reactivity ratio, rl, is approximately

16 4 for the copolymerization of ethylene with 1-octene, which is approximately 2 orders of m a g n i t u d e more reactive towards octene t h a n some MgC12-supported heterogeneous catalysts. In addition, non-traditional olefins such as s t y r e n e can be incorporated in high levels. Styrene / ethylene copolymers containing significant a m o u n t s of styrene and having a high molecular weight have not been available in the past, as conventional polyolefin catalysts will e i t h e r not copolymerize ethylene with styrene to any appreciable extent, or the molecular weight is too low to be u s e f u l .

5. C o ~ a t a l y s t s and Polymerization Behavior C o n s t r a i n e d geometry complexes and metallocenes in general r e q u i r e the addition of a cocatalyst in order to become catalytically active. When activated with a large excess of MAO, catalyst efficiencies between 150,000 and 750,000 g of polymer per g r a m of metal are obtained, depending on the reactor t e m p e r a t u r e , specific catalyst, MAO level and other process variables. In general, these efficiencies a r e lower t h a n can be o b t a i n e d w i t h b i s - c y c l o p e n t a d i e n y l m e t a l l o c e n e s a n d MAO. In solution p o l y m e r i z a t i o n s , however, h i g h Mw polymers are obtained, even at t e m p e r a t u r e s as high as 160 ~ with 1-octene as comonomer. The Mw decreases with i n c r e a s i n g t e m p e r a t u r e , as shown in Figure 3, as a result of increasing ~-hydride elimination to give u n s a t u r a t e d chain ends. The Mw is sufficiently high u n d e r solution conditions t h a t H2 can be used as a Mw control, giving two i n d e p e n d e n t controls over m o l e c u l a r weight. 200000 , .

150000-

D

100000 -

50000I

I

I

I

0

I

I

0 t""

Reactor T, ~ F i g u r e 3. Mw data for ethylene 1-octene copolymerization. Catalyst = [(C5Me4)SiMe2N(t-Bu)TiC12 / MAO, 450 psi ethylene, 10 minute reaction time. 11

17 The open nature of the catalytic site allows for the incorporation of extremely high levels of comonomer, producing elastomers with over 20 weight % 1-octene comonomer. Increasing levels of comonomer depress the density, leading to ultra-low density elastomers. High Mw elastomeric ethylene/octene resins with densities between 0.87 and 0.85 g/mL can be obtained with high efficiencies. Figure 4 shows the relationship between density and weight % 1octene, determined using 13C NMR for such elastomers. Prior to the advent of metallocene catalysis, such extremely low density copolymers were not commercially accessible at low cost.

0.88

0.87

[]

[]

0.86

0.85 30

40

50

60

Weight % Octene

Figure 4. Ethylene-co-l-octene density as a function of 1-octene content for elastomeric copolymers. Comonomer incorporation, molecular weight, and catalytic efficiency are sensitive to the nature of the group bridging the Cp ring and the substituent on the amide ligand. Shorter bridges constrain the cyclopentadienyl ligand and amide group to adopt a particularly open and reactive catalytic environment. 12 Unlike the bis-Cp metallocenes, the titanium constrained geometry catalysts generally show the highest activity, comonomer incorporation, and Mw. While the constrained geometry catalysts exhibit many unique properties, the catalytic efficiency using MAO cocatalyst is relatively low for commercial applications, on the order of 104 - 10 5 g polymer per g of Ti. In contrast, a variety of cationic constrained geometry catalysts can be prepared which show extremely high activity, exceeding 10 7 grams of polymer per gram of transition metal. Cationic catalysts can be prepared with ammonium salts (Figure 5a) 13, oxidation of a corresponding Ti(III) complex (Figure 5b) 14, or abstraction of a hydrocarbyl group using B(C6F5)3 (Figure 5c). 15

18

[R3NH] [B(C6F5)4] A)

e2SiNN~T1Me2

NaN ~

-

CH 4

t_B/

[Cp2Fe][B(C6F5)4] ...... ~

~

t.Bu]

.

~ 1 ,CH Me2Si\ iTi'" \2

Me2SiXNITi.

[B(CGF5)4] G Me

/

B)

t-Bu~ M e 2 ~ N ~

\

C)

~ ~ " Me2Si\NIT1Me 2 t'Bu]

/

Cp2Fe Me2Si( ~~.~,,,C.H2 [B(C6F5)47

N-'"[

B(C6F5)3 S. ~ / ~ ~ ~ Me2 1\NIT1 ~ t-Bu/

| [CH3B(C6F5)3]

Me

Figure 5. Formation of Cationic constrained geometry catalysts: The use of B(C6F5)3 as a cocatalyst is particularly useful for solution polymerization, as this cocatalyst is soluble in the hydrocarbon polymerization solvent. The polymers produced in a continuous solution polymerization using these constrained geometry catalysts possess the expected properties of narrow molecular weight distribution and uniform comonomer distribution across the entire molecular weight range. In general, a narrow molecular weight and comonomer distribution would be expected to improve physical properties at the expense of processability. The polyolefins produced using INSITE Technology in a continuous solution polymerization process at high temperatures and high ethylene conversion give a polymer with high shear sensitivity, low melt fracture, high melt strength, and easy processability. The unusual properties of these INSITE polyolefins is the result of small but significant levels of longchain branching in an otherwise linear molecule. 16 These long-chain branches are postulated to result from the reincorporation of vinyl-terminated polymer molecules according to Scheme 1. The conditions of high reactor temperature

19 which leads to high vinyl termination, high comonomer reactivity, and high conversion with the concomitant high polymer concentration and low ethylene and comonomer concentrations in the continuous solution process produces the conditions favorable to long-chain branch formation.

Ti--CH2CH2--polymer A

Ti

+

Polymer--CH-CH 2 +

TimH

+

Polymer--CH--CH 2

monomers

PolymermCH--Polymer I

Polymer S c h e m e 1. Mechanism for formation of long-chain branching: Ti = active constrained geometry catalyst. 6. Conclusions Metallocene catalysts allow for the targeted control of polymer molecular structure to a degree which has not been possible previously. Rational structure-property relationships of these homogeneous catalysts are allowing new polymers to be produced which are finding large commercial markets. The high catalytic productivity possible with metallocene catalysts enable the production of polymers at competitive prices, even though the catalysts and cocatalysts are complex and relatively expensive. For ethylene/a-olefin copolymers, constrained geometry catalysts allow the production of a unique family of olefinic polymers. The proper selection of the metal, bridging group, and other substituents allows the control of product properties in a high temperature solution process. With the proper selection of catalyst variables, products ranging from high molecular weight elastomers to high density polyethylene can be produced.

References

1. J. Boor, Jr., Ziegler-Natta Catalysts and Polymerizations, Academic Press, 1979, New York, p. 108-129. 2. Ibid, p. 8.

20 3. D. S. Breslow and N. R. Newberg, J. Am. Chem. Soc., 1959, 81, 81. 4. H. Sinn,and W. Kaminsky, Adv. Organomet. Chem. 1980, 18, 99. 5. W. Kaminsky, M. Miri, H. Sinn, R. Woldt, Makromol. Chem., Rapid Commun. 1983, 4, 417. 6. o

8. o

J.A. Ewen, J. Am. Chem. Soc. 1984, 106, 6355. W. Kaminsky, K. Kiilper, H. H. Brintzinger, F. R. W. P. Wild, Angew. Chem., Int. Ed. Engl. 1985, 24, 507. G.W. Coates and R.M. Waymouth, Science 1995, 267, 217-218. J. Boor, Jr., Ziegler-Natta Catalysts and Polymerizations, Academic Press, 1979, New York, p. 262-269.

10. J. A. Canich, U.S. Patent 5,026,798 (1991). 11. J. C. Stevens, et al., European patent application 416,815 (1991). 12. J. C. Stevens, Stud. Surf. Sci. and Catal. 1994, 89, 277-284. 13. J. C. Stevens and D. R. Neithamer, US patents 5,064,802 (1991); 5,132,380 (1992). 14. R. E. LaPointe, et al., US patent 5,189,192 (1993). 15. R. E. LaPointe, et al., European patent application 520,732 (1991). 16. S.Y. Lai, et al., U. S. patent 5,272,236 (1993); 5,278,272 (1994).

j.w. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

21

C h a r a c t e r i z a t i o n and C h e m i c a l Design of O x i d e Surfaces Yasuhiro Iwasawa Department of Chemistry, Graduate School of Science, The University of Tokyo, Hongo, Bunkyo-ku, Tokyo 113, Japan

This paper attempts to review recent and representative work dealing with single crystals, flat surfaces, and even designed surfaces relevant to oxide catalysis. These surfaces can provide novel information on the key issues in catalytic research such as the structure and composition of active sites, ensembles and phases, behavior of adsorbed active species, electronic property participating in catalysis, etc. which are well characterized by recent in-situ spectroscopy and also by traditional spectroscopy. The paper is also devoted to selected topics in the field describing the detail. Organized assembly of the knowledge of characterizations integrated over both molecular catalytic chemistry of powder catalysts and catalytic surface science of model surfaces would allow us to move toward the ultimate goal of rational catalyst design.

1. I N T R O D U C T I O N Metal oxides find application in a variety of technologies where surface science is critical to success, including catalysis, gas sensors, photoelectrolysis, electronic ceramics, semiconductor devices, pigments, cosmetics, etc. Metal oxides are tractable materials not only for spectroscopic techniques such as FT-IR, Raman, X-ray absorption fine structure(XAFS), etc., but also for techniques that might be disrupted by charging effects, including XPS, scanning tunneling microscopy(STM), etc., while special devices for sample preparation and careful interpretation of the spectra are essential. More critical is the tendency of many oxide surfaces to undergo thermal fracture, reconstruction, and particularly faceting[ 1]. These phenomena arise mainly from the need for charge balancing and minimization of surface polarity and energy, which may be relevant to the properties of oxide catalysts pretreated at different temperatures and ambient conditions. Understanding and controlling oxide surfaces are the key issues for the development of industrial oxide catalysts, but oxide surfaces are in general heterogeneous and complicated, and hence have been little studied so as to put them on a scientific basis by traditional approaches. While studies of the structure of surfaces have focused on metals and semiconductors over the past thirty years, the application of surface science techniques to metal oxides has blossomed only within the last decade[ 1-3]. An important future goal of catalytic surface science is to monitor the structure of surfaces and adsorbates at the molecular level in situ under catalytic reaction conditions, to model the more complex technical catalysts, and to undertake the design and tuning of new catalyst surfaces.

22 2. ACID-BASE AND REDOX P R O P E R T I E S OF MODEL S U R F A C E S

Acid-base reactivity is an important property of oxide catalysts, and its control is of interest in surface chemistry as well as being of importance in industrial applications. The exposed cations and anions on oxide surfaces have long been described as acid-base pairs. The polar planes of ZnO showed dissociative adsorption and subsequent decomposition of methanol and formic acid related with their surface acid-base properties[3]. Further examples related to the topic of acidbase properties have been accumulated to date[ 1,4-6]. In contrast to the extensive studies of heterogeneous acidic oxides, less effort has been given to the study of heterogeneous basic oxides. The first study of heterogeneous basic catalysts, in which sodium metal dispersed on alumina acted as an effective catalyst for double bond migration of alkenes, was reported by Pines and Ipatieff[7]. Now, a number of materials have been reported to act as heterogeneous basic catalysts; alkaline earth oxides, alkali metal oxides, rare earth oxides, ZaO2, ZnO and TiO2, alkali ion-exchanged zeolites, alkali metal ions on oxides, hydrotalcite, chrysotile, sepiolite, KF supported on alumina, etc[8,9]. A superbasic catalyst(~,-Al203-NaOH-Na) is prepared by addition of NaOH to alumina followed by further addition of Na. The resulting catalyst which possesses basic sites stronger than H_=37 and a distorted 13-Na~O2 phase at the catalyst surface as characterized by solid NMR and XPS, is industrially employed, with nearly 100% yield in the commercial plant for synthesis of 5ethylidene-2-norbomene (additive to ethene-propene copolymer rubber) from 5-vinyl-2norbornene that is obtained from dicyclopentadiene[10]. Considering the tendency of Na to donate electrons, it seems natural that Na dispersed on alumina acts as a heterogeneous basic catalyst. It is, however, found that the basic property of the oxide surface is not naturally proportional to the quantity of Na deposited on the surface[ 11 ]. Atom-resolved STM has visualized the mechanism producing structural sensitivity in the reaction of CO2, a reaction probe at basic sites, with the Na-deposited "I~O2(110) surface[ 12]. The amount of adsorption to form carbonates varies with Na coverage exhibiting an S-shaped dependence. An STM image of the 0.1 ML-Na deposited TiO2(l 10) surface depicts a dispersed geometry due to repulsive forces between the Na atoms. The surface is converted to a nearly complete c(4x2) overlayer via p(4x2) order locally formed on the surface by increasing Na coverage. The Na atoms in the c(4x2) surface are reactive to CO2 as shown in the STM topography in which the c(4x2) order disappeared and the chains for CO32 were locally ordered in a p(3x2) symmetry which appeared along the [001] direction. In contrast, CO2 does not adsorb on a surface with randomly dispersed Na atoms. The genesis of strongly basic sites is thus suggested not to be linearly correlated with Na quantity, but to be correlated with the ordered structure or at least suitable ensembles of Na. Recently, it has been demonstrated that coordination vacancies on the surface metal cations are relevant to the unique redox reactivity of oxide surfaces[2]. Oxidation of formaldehyde and methyl formate to adsorbed formate intermediates on ZnO(0001) and reductive C-C coupling of aliphatic and aromatic aldehydes and cyclic ketones on "HO2(001) surfaces reduced by Ar § bombardment are observed in temperature-programmed desorption(TPD). The thermally reduced "HO2(110) surface which is a less heavily damaged surface than that obtained by bombardment and contains Ti cations in the +3 and +4 states, still shows activity for the reductive coupling of formaldehyde to form ethene[ 13]. Interestingly, the catalytic cyclotrimerization of alkynes on TiO2(100) is also traced in UHV conditions, where cation coordination and oxidation states appear to be closely linked to activity and selectivity. The nonpolar Cu20(111) surface shows a

23 maximum selectivity for complete oxidative dehydrogenation of CH3OD to CO, while the polar Cu+-terminated (100) surface shows a maximum selectivity for partial dehydrogenation to CH20 [14]. Methanol on ZrO2(100) decomposes near 630 K to produce CO and CH4, whereas on the (110) surface the primary methoxide decomposition pathway is oxidation to produce CH20. This difference in reactivity can be related to the local atomic structure of each surface. Chemical and catalytic trends in oxide surfaces where metal atoms are isolated by oxide ligands resemble those observed in homogeneous metal oxo-complexes except for surface phenomena such as surface restructuring and transformation, diffusion of atoms and adsorbed molecules, etc[2]. 3. C H E M I C A L T U N I N G OF A C T I V E S I T E S The key properties of oxide surfaces are the coordination environment, oxidation state and acidic or redox properties of surface cations, and the basicity of surface anions. The longer term challenge to oxide surface science is to address important issues in selectivity in catalytic oxidation and acid-base reactions, in particular the principle of "tuning of metal reactivity by oxide ligands". There are a number of clear-cut examples in catalysis by mixed oxides, including the selective oxidation of propane to acrylonitrile with V-Sb oxide[15-17], Bi-V-Mo oxide [18,19], Mol.0V0.4Te0.2Nb0.104.65 oxide[20], and the selective oxidation of butane to maleic anhydride with VPO catalysts[21 ], where selectivity toward the desired products relies on the limited availability of oxygen at active ensembles of the metal oxide component and also on the properties and structure of defects, besides the nature of surface oxygen atoms. Notwithstanding the interesting industrial outlook, relatively few data exist either on the characterization of such catalysts or on the reaction kinetics and mechanisms. In the development of new catalysts, new chemical concepts regarding composition or structure are conceived. The requirements and design of quantitative ensemble sizes represent important but as yet unaddressed challenges to the field. Although the efforts on the design of excellent catalysts have been acutely difficult challenges, recently molecular-level catalyst preparation has become realistic on the basis o'f modem physical techniques and accumulated knowledge of oxide surfaces[22-25]. A Series of chemical designs of Nb structures on SiO2 is introduced as an example of a onecomponent tuning catalyst. Niobium has been considered to be a poor catalyst, but recently has attracted much attention as a key element for industrially important processes such as amrr,oxidation of propane[20], oxidative dehydrogenation of propane[26], etc. A guideline for the Nb-structure design is seen in a theoretical concept. We consider and extend this theoretical prediction to the case of ethanol oxidation, via a Nb-OC2H5 intermediate, on NbS+-oxide sites supported on a SiO2 surface[27]. In ~ C H elimination on dS-metal ethoxide complexes, the orbital interactions have to take place in such a manner that the electron donation from a(CH) to g*(MO) and the back-donation from g(MO) to c*(CH) are required to form the MH g and CO n bonds and break the CH a and MO c bonds. Also in case of d~ complexes, the presence of a weak M--H agostic interaction is predicted by the theoretical calculation, but the Ti-~CH angle is unfavorable for overlap of the occupied Ti d-orbital and the CH antibonding orbital. Furthermore, there is formally no d-electron available for the promotion of the CH bond scission. The predicted key factors are electron density of the d state and overlap of the two orbitals. The former boundary is satisfied by attaching Nb do ions to the SiO2 surface through Nb-O-Si bonding. The electronic structure of a distorted tetrahedral dioxo-Nb monomer on SiO2 calculated by the DV-Xct cluster method shows that the component of the Nb 4d orbitals is hybridized with the higher occupied O 2p levels, enabling the ~ C H breaking. The support electronically modifies the metal oxide

24

species through chemical or ionic bonds and induces the structural change of the surface metal oxides needed for catalysis. It predicts new catalysis involving I],-elimination of the CH bond by coordinatively unsaturated tetrahedral Nb monomers chemically attached to the SiO2 surface. The latter boundary of the orbital overlap seems to be less rigid for the tetrahedral Nb monomer structure. The SiO2-attached Nb monomer catalyst with a four-coordinate structure was prepared by the use of Nb013-C3H5)4 as precursor and characterized by extended x-ray absorption fine structure(EXAFS), FT-IR, Raman, ESR, and XPS. The monomer catalyst(l) exhibits high activity and selectivity for the dehydrogenation of ethanol to form acetaldehyde and H2 as shown in Table 1[27]. The activity is much higher than that of a usual impregnation Nb catalyst and the selectivity is as high as 95-100%, whereas the impregnation catalyst is less active and unselective. The dehydrogenation reaction proceeds via the Nb-ethoxide intermediate, but the intermediate is very stable and is not decomposed until 600 K. It is dehydrated to ethene and water above 600 K. On the other hand, when the Nb-ethoxide intermediate is exposed to ethanol, the dehydrogenation of ethanol proceeds at much lower temperatures such as 423 K, with ca. 100% selectivity. Thus the switchover of the reaction path from dehydration(~,-CH bond break) to dehydrogenation(l~-CH bond break) by the second ethanol molecule adsorbed on the Nb-ethoxide is observed. Even if an adsorbed species at the surface is too stable in vacuum or before catalysis, the catalytic reaction is able to proceed through the same species by activation of the intermediate by the reactant(reactant-promoted mechanism)[28]. Thus, it may be critical for the dehydrogenation on Nb sites to create a vacant site with an appropriate conformation for the transition state on which electron donation-induced activation of ~hydrogen of C2H50 group is favorable. , 0.307nm, ;~

, '<

" "'1

........

I

O..."

YI'/5'~' ....

(1)

I

I

I.."

0.342nm, .~; ~

0.378nm.

,

~.:.."~tO.21 in m

"'~

.................... ~ ' 1 8 2 ' ................. (2)

......................... ,~ I'(~'2 ......................

(3)

Table 1 Catalytic performance of the Nb monomer, dimer, and monolayer supported on SiO2, and an impregnation Nb catalyst for dehydrogenation and dehydration(intra and inter) of ethanol Catalyst Initial rate(mmol/min/g-Nb) Selectivity(%) Total AA E+DE AA E DE Monomer* 1.25 1.20 0.049 96.1 2.8 1.1 Dimer* 0.18 0.004 0.176 2.1 24.2 73.7 Monolayer** 0.11 0.001 0.106 0.9 99.1 0.0 lmpreg.* 0.17 0.052 0.118 30.5 20.2 49.3 bulk** 0.0026 0.0004 0.0022 14.8 46.9 38.3 AA: Acetaldehyde, E: Ethene, DE: Diethyl ether, Ethanol: 3. I kPa, *: 523 K, **: 573 K. If a vacant site is occupied by another Nb atom, such that it is a dimer, new catalysts may be designed. The Nb dimer catalyst(2) was prepared by reaction of [Nb(rlS-CsHS)H-I.t-(rlS,rl 1C5H4)]2 with a SiO2 at 313 K, followed by treatment with 02 at 773 K. A proposed structure(2) was characterized by EXAFS, x-ray absorption near-edge structure(XANES), FTIR, UV-vis, and XPS, which shows Nb-Nb (coordination number: 0.9) and Nb-Si(2.3)

25 interatomic distances at 0.303 nm and 0.328 nm, respectively, besides Nb-O bonds[29]. The Nb dimer/SiO2 catalyst(2) shows high selectivity for the dehydration of ethanol irrespective of the presence or absence of the ambient ethanol(Table 1). The dehydrogenation observed on the Nb monomer catalyst is remarkably suppressed to 1/300 and the dehydration is promoted 4 times on the dimer, indicating that the Nb dimers on SiO2 have an acidic character. It is to be noted that the change of the number of Nb atoms at the active sites from one to two metal atoms gives rise to a complete reverse of basicity/acidity in the catalytic properties[29]. In the dimer(2) the access of a second ethanol molecule to the Nb atom coordinated with a first ethanol molecule in a preferable conformation is difficult unlike the case of the monomer(l). Furthermore, the Lewis acidity of the Nb atoms in the dimer is increased by the oxygen-bridge. The Nb-oxide monolayer(3) was successfully prepared by using a Nb(OC2HS)5 precursor. The monolayer growth was monitored mainly by Nb--Si bond(0.327 nm) formation characterized by EXAFS. The monolayer(3) shows two different Nb-Nb separations at 0.378 and 0.342 nm, both being considered to have one or two bridging oxygens(Nb-O: 0.211 nm), respectively. The monolayer(3) is active and selective for C2H4 formation from C2H5OH in the temperature range 373-573 K as shown in Table 1. The monolayer catalyst(3) always shows the selective intramolecular dehydration of ethanol, suggesting that the Lewis acid sites in the monolayer niobium oxides may be distributed in an isolated manner. The niobium oxide layer is somewhat distorted by the structural mismatch and the strong Nb-O-Si interaction between the Nb oxide overlayer and the SiO2 surface as proved by EXAFS. The distortion and mismatch should be released by the creation of the Lewis-acidic Nb sites. Nb-oxide monolayers on various oxide supports have been extensively investigated and characterized by Raman spectroscopy[30]. These structures are described as being made up of an octahedrally coordinated NbO6 structure with different degree of distortion. The Lewis acid sites may be dispersed basically in the NbO6 overlayer. This type of catalyst is applicable to intramolecular dehydration processes. These new and distinct materials and chemistry prepared stepwise in a controllable manner b~, using organometallic and inorganic complexes as precursors provide an opportunity for the development of efficient catalytic molecularly-organized surfaces en route to the ultimate catalyst technologies. The notable superiority of these catalytic systems is in single-site reaction, high selectivity, generality(wide applications), and tunability. The design and characterization of the Mo2/A1203 and Mo2/SiO2 are among this class of catalysts[23-25]. The Mo sites have been shown to be entirely different from those in MOO3. Principal modes of supporting precursors to inorganic surfaces are classified in the literature[24]. Applications of the principle to industrial catalytic systems are as follows; (1)vanadium(acetylacetonate)3/SiO2(with AI(C2H5)2CI) for ethene-propene copolymer rubbers[31 ], (2)CrOx/SiO2 and NiA,-AI203 by atomic layer epitaxy of CrO2Cl2 or Cr(acac)3, and Ni(acac)2, respectively, for ethene polymerization and toluene hydrogenation[32], and (3)Diphosphine-rhodium/isocyanate-alkyl-trialkoxysilane/SiO2 for enantioselective catalytic hydrogenation[33]. 4. C H A R A C T E R I Z A T I O N OF CATALYTIC P H E N O M E N A Characterization of catalytic phenomena at oxide surfaces includes: (1) characterization of established catalyst surfaces to improve the catalytic performance, (2)characterization of new catalysts in comparison with conventional catalysts, (3) characterization of specific model surfaces such as single crystals and epitaxial flat surfaces to transfer the knowledge so obtained to catalytic systems or even to create a new type of catalyst, and (4) characterization of catalysis

26 for the understanding of catalytic phenomena, optimizing catalytic operations and extracting essential factors for the genesis of the best catalysis. The most important objective in the characterization of metal oxide surfaces requires a depth of knowledge similar to that available in homogeneous catalysis. Recently, the characterization of oxide surfaces at the atomic and molecular level has received a great boost from the development of a variety of sophisticated techniques, including scanning probe microscopy(SPM) like STM and atomic force microscopy(AFM), XAFS, high resolution electron microscopy(HREM), solid-state magic angle spinning(MAS) NMR, reflection absorption infrared spectroscopy (RAIRS or IRAS), which provide valuable support for the mostly empirical approach to catalyst design. A highly detailed picture of a reaction mechanism evolves in-situ studies. It is now known that the adsorption of molecules from the gas phase can seriously influence the reactivity of adsorbed species at oxide surfaces[24]. In-situ observation of adsorbed molecules on metaloxide surfaces is a crucial issue in molecular-scale understanding of catalysis. The transport of adsorbed species often controls the rate of surface reactions. In practice the inherent compositional and structural inhomogeneity of oxide surfaces makes the problem of identifying the essential issues for their catalytic performance extremely difficult. In order to reduce the level of complexity, a common approach is to study model catalysts such as single crystal oxide surfaces and epitaxial oxide flat surfaces. STM has particularly great potential for in situ chemical studies. While our present knowledge of the atomic structure of catalyst surfaces is largely limited to those structures which are stable in ultra-high vacuum before and after reaction, STM may provide an insight into both adsorbate and catalyst surface structure in situ during the reaction. The following issues to be characterized by STM may be most relevant to characterization of catalysts and catalysis; (l)Identification of structural characteristics of the variety of non-equivalent surface sites and observation of site specificity to reactivity, (2)Study how the reaction at one particular surface site affects the local activity of neighboring sites, (3)Structural transformation and chemical modification of the surface caused by adsorbates and chemical reactions, (4)Detailed information about the mechanism of the surface chemical reaction, and (5)Surface diffusion and surface mobility. Under favorable conditions, a vertical resolution of hundredths of an ~gstrom and a lateral resolution of about one ~mgstrom can be reached. Therefore STM can provide real-space images of surfaces of conducting materials down to the atomic scale. The general principle of operation of STM is surprisingly simple. In contrast to other electron microscopes and surface analytical techniques using electrons, STM can be operated in air, liquids, vacuum, and reaction conditions because there are no free electrons involved in the STM experiments. Following the invention of the STM in 1981134,35], the surface structures of several transition metal oxides such as TiO2(110)[36-40], TIO2(100)[41], V205(001)[42,43], Rbl/3WO3(0001)[44], Fe304(001)[45], and Mo18052(100)[46] have been observed in real space with atomic resolution by the use of STM. Recently, it has been shown that in-situ STM studies at high pressures and temperatures are indeed feasible[47]. The experiments show the utility of variable temperature STM for the investigation of surface chemistry. A serious limitation of the STM technique so far is its lack of chemical sensitivity. Generally, STM is not specific for the elemental species in multi-component systems, though there are special cases where the direction of charge flow is well known as shown for the GaAs(110) surface. The surface area which one is looking at by STM is typically quite small. The problem of how representative the obtained 'tunnel vision' is, is at least partly solved by considerably increasing the total scan range of STM/SPM instruments.

27 XAFS is a most powerful technique in the characterization of structures and electronic states of any sort of catalytic material, though its analysis and interpretation should not be overestimated. XAFS can provide information not only in a static state but also in a dynamic state(in-situ conditions) that is not given by other techniques for characterization of catalysts and related surfaces[23,24,48-52]. Metal oxide phases in a highly dispersed form or thin layers on inorganic oxide supports are not naturally symmetric/isotropic. The chemical state and surface energy of the metal oxide, the texture and morphology of the support, the interaction between metal oxide and support, etc. are often asymmetric or anisotropic. The Mo dimer/SiO2 catalyst is a typical sample which shows an anisotropic change of 0.04 nm in a lateral direction(Mo-Mo separation) and 0.01 nm in a longitudinal direction(Mo-surface separation) during ethanol oxidation[23,24]. In the case of powder samples such as typical heterogeneous catalysts, however, the structures determined by XAFS are averaged structures in every direction of the sample. When flat surfaces like single crystals are employed as supports, characterization of metal sites on them can be achieved separately in three different bond directions parallel(two inplane directions) and normal(one direction) to the surface by polarized X-ray stemming from synchrotron radiation. For K-edge and LII, IN-edge EXAFS spectra of oriented samples, effective coordination numbers N* are given by N* = 3 ~:cosE0j and N* = 0.7 N + EcosE0j, respectively, where cosE0j is the angle between the electric-field vector of the incident X-ray and the bond vector[53]. Thus, when the polarization of the incident X-ray is normal to the surface(p-polarization), the X-ray absorber's neighbors which lie along lines parallel to the surface do not contribute to EXAFS signals. These bonds are preferentially observable when spolarized X-rays parallel to the surface are used. In reverse, the bonds normal to the surface are favorably detectable by p-polarized EXAFS. By changing the orientation of the sample to polarized synchrotron radiation and using the total-reflection fluorescence method(PTRFEXAFS), it is possible to determine the anisotropic or asymmetric structures on surfaces; e.g. Cu atom on ix-quartz(0001)[54], CoOx and Co304 on tx-A1203(0001)[55], Pt clusters on o~A1203 (0001)[56,57], and MoOx on TiO2(110)[57]. A schematic diagram of the PTRF-XAFS for in situ characterizations is shown in Fig.1. The bond orientation on surfaces can also be determined by in-situ PTRF X-ray absorption near-edge structure(XANES) in the same system as Fig.l; e.g. VOx/ZaO2(100)[58].

X-ray S l i l ~ M ~M i

,i 0 M'---

M---M

i

~M

'I

O'

M'~

i I

--M 0

I

,

I

M'--M'

---M---M

Iololol

M'- - - M:- - -M" - - M'

ml-o_l_o-l-o-I_ iiiiiiIinll

iiIIlllllrllllllllllmlllllll

Figure 1. Schematic diagram of polarization-dependent total-reflection fluorescence XAFS.

28 5. P R I N C I P L E S OF S U R F A C E R E A C T I V I T Y F O R OXIDE CATALYSIS Most of the experimental and theoretical studies on metal-oxide single-crystal surfaces have been focused mainly on the characterization of adsorbates and their noncatalytic reactions under vacuum by means of photoelectron spectroscopy and TPD. Few pioneering works examined catalytic performance on ZnO single crystal surfaces, where reaction sites were assigned into defects generated by the reduction of the surfaces[3]. It has been thought that the acid-base character is an intrinsic property of oxide substrates. The selectivity in the catalytic decomposition reaction of formic acid has been used to scale the acid-base property: dehydration over acidic oxide and dehydrogenation over basic oxide, though this classification is oversimplified vide infra. A new aspect of acid-base catalysis was found on TiO2(110), where a switchover of the reaction paths from unimolecular dehydration to bimolecular dehydrogenation, is seen when a formic acid molecule participates in the decomposition process of a surface formate. It is to be noted that acidic molecules induce basic catalysis[6]. The catalytic dehydration reaction is strongly suggested to involve the unimolecular decomposition of formate as a rate-determining step. In other words, the formate-surface interaction activates the unimolecular decomposition of formate(HCOO(a)) to preferentially yield CO(g) and OH(a). An acidic proton of a HCOOH molecule, which encounters the surface in a steady state, reacts with the resultant OH(a) to form H20. The two reaction steps form a catalytic dehydration cycle with an activation energy of 120 kJ/mol as shown in Scheme 1. At lower temperatures a formic acid molecule selectively reacts with the formate before the unimolecular decomposition to open a new basic-catalysis path with an activation energy of 15 kJ/mol by a bimolecular reaction mechanism. In this case the lateral nEO(g) HCOO(a) _..,I HC~O(a) + HCOOH(g)

~'-

HCOOH(g) _ _

]

~

CO(g) + OH(a)

~ CO2(g) + H2(g) + q

Scheme 1. Switchover of the reaction path by acidic molecules to basic catalysis. coulombic(acid-base) interaction between H ~i- of H C O O ( a ) a n d H § of HCOOH is more important than the vertical coulombic(acid-base)interaction between HCOO(a) and Ti 4§ Although TiO2 powder is a selective catalyst for the dehydration, the present study shows that "HO2(l10) favors the dehydrogenation at low temperatures, which might be an unexpected feature. It is possible on oxide single crystals to find new catalytic reaction paths by controlling the coverage and bond arrangement of intermediate and reactant, which are hardly obtained by powder catalyst surfaces. In fact the latter feature is achieved on TiO2(110) surface which has the alternative alignment of the exposed, five-fold coordinated Ti-row and the bridging O-ridge row, resulting in the characteristic anisotropy on this surface. Individual Ti ions on the clean surface are resolved as axial corrugations on the rows along [001] direction in the STM image(Fig.2a). Formate ions on TiO2(110) form a monolayer ordered in a (2x l) periodicity, when saturating the surface. The C-H bond of a formate is normal to the surface, while the O-CO plane is parallel to the [001] axis. These formate ions were removed from a scanned area by

29 tip fabrication in Fig.2b[59]. Time-resolved STM-observation of the void refilled by diffusing formate ions reveals molecular-scale kinetics of formate transport, which is related to the origin of the catalytic decomposition of formic acid on TiO2(110). The void shrinks with time as shown (b)-(e), and disappears in (f). The boundaries between the void and the migrating monolayer remains very clear. It was found that the rate of boundary shift was strongly anisotropic. The boundaries parallel to the [1i'0] direction advance toward the [001] or [00i] direction at a rate of 0.15 nm/min, whereas the boundaries parallel to the [001 ] axis shift at a much smaller rate, 0.02 rim/rain or less, toward the ['I 10] or [I 10] direction. The anisotropic transport suggests that the formate ions slide on Ti-rows with their O-C-O plane parallel to the row. Thus it is possible to control the surface reaction in an orientation-controlled condition through the highly regulated transport of adsorbed species.

Figure 2. Serial STM images(29x28 nm 2) for the formate migration on TiO2(110)(a) recorded at (b)15, (c)26, (d)35, (e)50, and (f)63 min after the rastering(14xl4 nm2). Recently, a surface/gas phase reaction on TiO2(110) was first visualized in situ by STM at high temperature[60]. Surface reaction of metal oxides especially under reactive atmospheres plays a crucial role in many catalytic processes. Dynamic observation during surface reaction is indispensable for understanding how those processes occur on the oxide surfaces. This kind of information also helps us to consider the relation of surface structure to reactivity, a fundamental issue in surface catalysis. When an 02 atmosphere(lxl0 -5 Pa) was introduced onto the TiO2(110)-(lxl) surface, many hills were randomly nucleated over the terraces. The random nucleation of hills suggests that the Ti n+ ions to be oxidized migrated vertically from the bulk to the surface, by hopping from one O-octahedron to another. The hills are then transformed to added Ti203 rows, showing spikes in the STM images which reflect vigorous movements of surface atoms. Added rows and new terraces shrink on average during 02 exposure. This process continues until the Tin+(n_~<3) ions from the bulk accessible to the surface have been exhausted. Ti3+ electron donor ions are formed by high temperature activation of TIO2161,62]. The ESR peak intensity of Ti 3§ and the 1-butene isomerization through a rt-allyl carbanion on the Ti 3§ sites show a maximum around 750 K. The dynamic features observed on TiO2(110) may be conceived to visualize the surface activation process of oxides.

30

Reactant-promoted mechanisms for reversible water-gas shift(WGS) reactions are briefly described, which involve activation of reaction intermediates(formates) with a drastic selectivity change by weakly coadsorbed reactant molecules under the catalytic reaction conditions[28]. It is to be noted that on MgO the formate produced from H20+CO is not decomposed forwardly to H2+CO2 in the absence of water vapor and the formate is quantitatively decomposed backwardly to H20 and CO again, whereas in the presence of weakly adsorbed water, the formation of H2 and CO2 is observed. Adsorbed water molecules do create a reaction path for the dehydrogenation of the formates. On ZnO and Rh/CeO2 surfaces, the rates of the formate decomposition are promoted by a factor of 10-100 by water vapor. Seventy % of the bidentate formates on ZnO backwardly decompose to the original H20+CO under vacuum, whereas in the presence of second water molecules almost 100% of the formates are able to decompose to H2+CO2. Note that the origin from which the formate is produced is "remembered" as a main decomposition path under vacuum, while the decomposition path for the formate is "forgotten" by coadsorbed H20. The WGS reaction is a reversible reaction, that is, it attains equilibrium with reverse WGS reaction. Thus the fact that the WGS reaction is promoted by H20(a reactant), in turn, implies that the reverse WGS reaction may also be promoted by a reactant, H2 or CO2. In fact the decomposition of the surface formates produced from H2+CO2 is promoted 8-10 times by gasphase hydrogen. The WGS and reverse WGS reactions can conceivably proceed on different formate sites of the ZnO surface unlike usual catalytic reaction kinetics, while the occurrence of the reactant-promoted reactions does not violate the principle of microscopic reversibility[63]. The observed synergy and the dynamic activation of intermediates may be associated with the genesis of catalysis. The phenomenon is similar to that observed for the ethanol dehydrogenation on the Nb monomer/SiO2 catalyst as stated above[27]. It has also been reported that the reactivity of the formate ions on a Ni/SiO2 catalyst was markedly increased by the presence of formic acid in the ambient gas compared to that at the same coverage under vacuum[64]. As described above, adsorbed molecules or impinging molecules under catalytic reaction conditions play an important role in surface catalytic reactions even if the adsorption of "promoter" is very weak or undetectable. This concept may have implications for a new way to prepare designed catalysts through in-situ activation of inactive or stable surface sites under the conditions of a target reaction[24,28]. 6. C H A R A C T E R I Z A T I O N AND CATALYSIS OF T H I N - F I L M O X I D E S The method of using an epitaxial oxide film on a conductive transition metal single crystal as a model surface for the corresponding oxide catalyst has been proposed as a model catalyst sample to solve the difficulties in obtaining oxide single crystals that we want to examine and in charging problems with insulator surfaces. The range of the surface characterization techniques can be extended further to model thin-film oxides, although these require an additional level of effort in preparation and characterization, and may not always be easy to prepare with adequate control of crystallinity or surface orientation. An obvious question about this method is the question of the structural and chemical similarity and difference between the model surface and the actual oxide surface. In some cases the produced oxide films may also be new phases with novel properties which are hardly formed and stabilized on the oxide single crystal bulk. Growth, nucleation, and structure of a variety of metal oxide overlayers have been characterized by STM, AES, XPS, ISS(LEIS), FIM/FEM, CO TPD, crystal current, etc. Representative catalytic samples are listed in Table 2.

31

Table 2 Oxide overlayers characterizedfor their surface structure and growth mode Overlayer Substrate Remarks Ref. TiO2 Pt(1119) Stranski-Krastanov growth (hydrogenation) (1) NbOx Pt(11 1) Stranski-Krastanov growth (distorted NbO6 octahedral) (2) TiO2 Ni(polycrystalline) Stranski-Krastanov growth (3) A1203 Fe(100),(110),(111) Three dimensional growth (NH3 synthesis) (4) MgO(ll 1) Mo(ll0) thin film(adsorption of [Re2(CO)I0] and [HRe(CO)5]) (5) MgO(100) Mo(100) thin film (acid/base properties) (6) Li/MgO(100) Mo(100) thin film (selective oxidation of methane to ethane) (7) ZnO Cu( 111 ) thin films (8) PbO Ag(polycrystalline) thin film (9) ZnTiOx Si thin film (catalyticdehydrogenaton of hydrocarbons) (10) thin films (unstable bulk surface, NO(a)+OH(a)-->HNO2)(11) CoO(I 11) Co(0001 ) TiOx Pt(polycrystalline) thin film (CO + 3H2 = CH4 + H20) (12) A1203 Ta(110) thin film (Epitaxialgrowth(0.5-4 nm)) (13) MgO(100) Mo(100) thin film (Li+-doped, [Li+O] and color centers, OCM ) (14) thin film (adsorption) (15) NiO(100) Mo(100) (1)M.E.Levin, M.Salmeron, A.T.Bell and G.A.Somorjai, J.Chem.Soc., Faraday Trans.1, 83(1987)2061; C.S.Ko and R.J.Gorte, Surf.Sci., 155(1985) 296. (2)L.Xie, D.Wang, C. Zhong, X.Guo, T.Ushikubo and K.Wada, Surf.Sci., 320(1994)62. (3)G.B.Raupp and J.A.Dumesic, J.Catal,, 95(1985)587; J.Phys.Chem., 88(1984)660. (4)M.E.Levin, M. Salmeron, A.T.Bell and G.A. Somorjai, J.Catal., 106(1987)401; J.P.S. Badyal, A.J.Gellman, R.W.Judd and R.M.Lambert, Catal.Lett., 1(1988)41; D.Chadwick, A.B.Christie and M.A. Karolewski, Vacuum, 31(1981)705. (5)J.S.Comeille, J.-W.He and D.W.Goodman, Surf.Sci., 306(1994)269; S.K.Purnell, X.Xu, D.W.Goodman and B.C.Gates, J.Phys.Chem., 98(1994) 4076. (6)M.-C.Wu and D.W.Goodman, Catal.Lett., 15(1992)1. (7)M.-C.Wu, C.M.Truong, K.Coulter and D.W.Goodman, J.Am.Chem.Soc., 114(1992)7565; M.-C.Wu, C.M.Truong and D.W.Goodman, Phys.Rev., B46(1992)12688. (8)C.T.Campbell, K.A.Daube and J.M.White, Surf.Sci., 182(1987)458. (9)D.Chadwick, A.B. Christie and M.A.Karolewski, Vacuum, 31 (1981)705; D.Chadwick and M.A.Karolewski, Appl. Surf.Sci., 9(1981)98. (10)Z.X.Chen, A. Derking, W.Koot and M.P.van Dijk, to be published. (ll)M.Hassel and H.-J.Freund, Surf. Sci., 325(1995)163. (12)R.A.Demmin, C.S.Ko and R.J.Gorte, J.Phys.Chem., 89(1985)1151. (13)P.J.Chen and D.W.Goodman, Surf.Sci., 312(1994)L767. (14)M.L.Burke and D.W. Goodman, Surf.Sci., 311(1994)17. (15)M.-C.Wu, C.M.Truong and D.W. Goodman, J.Phys. Chem., 97(1993)9425. 7. FUTURE PROSPECTS Metal oxide surface science is a rapidly expanding frontier area, while the development of formulations and techniques for producing ordered, well defined oxide surfaces with controlled structure, composition, oxidation state distribution, and morphology remain important challenges to researchers in this field. While unifying concepts on well defined surfaces may on occasion produce oversimplifications or omissions, they clearly provide a more useful framework for better understanding of catalysis and the design of catalysts. The trends that are particularly relevant to future progress in catalysis are the observation of catalytic performance on the surface in real space and real time in situ under reaction conditions and the molecular design of oxide catalyst surfaces on a more rational basis. With the further development of in-

32 situ STM methods, future experiments should enable morphological changes to catalyst surfaces to be followed in real time as the reaction proceeds in order to establish not only the details of the surface restructuring mechanisms but also accurate relationships between surface structure and catalytic performance. A breakthrough is needed for real time imaging for faster reactions. X-ray photoemission electron microscopy(XPEEM) can provide information on the composition and chemical state of surface species, and their spatial and temporal resolution during cross diffusion of multi elements at oxide surfaces, growth of mixed oxide layers, selective adsorption and catalytic reactions. This new technique is applicable to complex and inhomogeneous, mixed oxide catalysts of vital practical importance, which have been considered to be not well suited to fundamental studies because their surfaces are ill-defined and poorly characterized. Another promising future characterization technique to monitor surface reactions proceeding on surfaces in real time may be temperature-dependent electron stimulated desorption ion angular distribution (ESDIAD). Infrared-visible sum frequency generation(SFG) at interfaces has been demonstrated to be a new probe of the surface vibrational state. Surface SFG is applicable to gas/solid, liquid/solid, gas/liquid, and solid/solid interfaces. Surface tailoring has strategic advantages in the synthesis of catalysts with surface structures and compositions which are analogous or unusual in homogeneous systems, and acutely difficult to make in traditional heterogeneous systems. Chemically designed metal oxide catalysts may also contribute to a complete understanding of the origin of heterogeneous catalysis and hence basic implications and new strategies for the development of new types of promising catalysts. The role of a catalyst is to generate an active intermediate responsible for success of the process, while the intermediates for catalytic reactions are not necessarily reactive when they are located as isolated surface species, but are capable of acting as mediator in situ under the catalytic reaction conditions. This concept offers exciting prospects for the creation of new materials and novel chemistry by surface synthesis techniques. Despite the success in modeling catalysts with single crystals and well defined surfaces, there is a clear need to develop models with higher levels of complexity to address the catalytically important issues specifically related to mixed oxide surfaces. The characterization and design of oxide surfaces have not proven to be easy tasks, but recent progress in identification of the key issues in catalytic phenomena on oxide surfaces by in-situ characterization techniques on an atomic and molecular scale brings us to look forward to vintage years in the field.

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W.Lu, N.Nevins, M.L.Norton and G.S.Rohrer, Surf.Sci., 291 (1993)395. G.Tarrach, D.Btirgler, T.Schaub, R.Wiesendanger and H.-J.Giintherodt, Surf.Sci., 285(1993)1. G.S.Rohrer, W.Lu, R.L.Smith and A.Hutchinson, Surf.Sci., 292(1993)261. B.J.Mclntyre, M.B.Salmeron and G.A.Somorjai, Catal.Lett., 14(1992)263; Rev. Sci.Instrum., 64 ( 1993)687; J. Vac.Sci.Technol.A, 11(1993) 1964. D.C.Koningsberger and R.Prins, X-Ray Absorption: Principles, Applications, Techniques of EXAFS, SEXAFS and XANES, Wiley, New York, 1988. E W.Lytle, G.H.Via and J.H.Sinfelt, Synchrotron Radiation Research(H.Winick and S.Doniach eds.), Plenum, New York, 1980, Chap. 12. J.C.J.Bart and G.Vlaic, Adv.Catal., 35(1987)1. B.K.Teo, EXAFS: Basic Principles and Data Analysis, Springer, Berlin, 1986. YIwasawa (ed.), X-ray Absorption Fine Structure(XAFS) for Catalysts and Surfaces, World Scientific, Singapore, 1996. E.A.Stern, Phys.Rev., B 10( 1974)3027. M.Shirai, K.Asakura and YIwasawa, Chem.Lett., (1992) 1037. M.Shirai, T.Inoue, H.Onishi, K.Asakura and YIwasawa, J.Catal., 145(994)159. K.Asakura, M.Shirai and Y.Iwasawa, Catal.Lett., 20(1993)117. W.-J.Chun, K.Asakura and YIwasawa, Proc.13th Int.Vacuum Conf./9th Int. Congr.Solid Surf.Yokohama.1995, in press. M.Shirai, K.Asakura and YIwasawa, Catal.Lett., 26(1994)229. H.Onishi and YIwasawa, Jpn.J.Appl.Phys., 33(1994)L1338. H.Onishi and YIwasawa, Phys.Rev.Lett., in press. M.Che, C.Naccache and B.Imelik, J.Catal., 24(1972)328, M.Che, C.Naccache, B.Imelik and M.Prettre, Comptes Rendus, 264C(1967) 190 I. H.Hattori, M.Itoh and K.Tanabe, J.Catal., 38(1975) 172. T.Shido and Y.lwasawa, J.Catal., 140(1993)575. K.Takahashi, E.Miyamoto, K.Shoji and K.Tamaru, Catal.Lett., 1(1988)213.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

35

Photocatalysis: State of the Art and Perspectives Kirill I. Zamaraev; State Research Center "Boreskov Institute of Catalysis"; Novosibirsk 630090, Russia 1. INTRODUCTION Photocatalysis is a fundamental feature of life processes on our planet [1] (it provides photosynthesis in plants and bacteria) and of the chemistry of its atmosphere [2]. Work is under way to develop photocatalytic technologies for abatement of environmental problems [3, 4]. Photocatalysis is anticipated to become in the coming years important also for selective organic synthesis [4]. In a more distant future thermal catalytic processes induced by heating with solar radiation, together with photocatalytic processes may become important for environmentally friendly technologies of solar energy utilization [5-9]. Upon absorbing one or several photons, a photocatalyst (PC) by means of intermediate chemical interactions with the reaction components, provides transformations of the reagents into products via cyclic reaction pathways, being regenerated after each reaction cycle. A photon serves in photocatalytic reactions as an additional energy-rich reagent, which provides new reaction routes and can withdraw thermodynamic restrictions for reactions that would be strongly endergic in the absence of light. This reagent can be provided by sunlight at zero cost, however, it may become rather expensive when provided from electrically fed light sources. In the last two decades we have witnessed in photocatalysis, as a science, a continuous shift from phenomenological approaches to studies at the molecular level. With accumulation of information obtained in such studies, the accents in the work aimed at development of new photocatalysts and new photocatalytic reactions and technologies, are expected to more and more shift from empirical search to intentional design. In this paper the state of the art and perspectives will be discussed, of basic and applied photocatalysis and thermal catalysis induced by solar radiation, with the emphasis on the following areas: 1. Frontiers of basic reasearch on photocatalysis, including studies of simple molecular systems, organized molecular assemblies and semiconductors as photocatalysts. 2. Photocatalysis in selective organic synthesis. 3. Photocatalysis in abatement of environmental problems. 4. Present developments and perspectives for applications in the future, of photocatalysis and thermal catalysis induced by solar radiation, for utilization of solar energy. 5. The expected role of photocatalysis in the global chemistry of the atmosphere. Of the numerous works described in the literature, particular examples are selected, which hopefully will let a reader to feel "the flavour" of R&D, as well as potential applications in various areas of photocatalysis.

36

2. FRONTIERS OF BASIC RESEARCH Most of the so far designed photocatalytic systems can be divided into three categories: simple molecular ones, organized molecular assemblies and semiconductor systems. In this section typical photocatalytic behaviour and reaction mechanisms will be discussed for photocatalysis with systems of all these types.

2.1. Simple molecular systems To this category belong, e.g., homogeneous photocatalytic systems based on soluble metal complexes or organic dyes as photocatalysts. Instructive examples are photoreactions assisted by heteropolyacids (HPAs), transition meal complexes with carbonyl, phosphine or some other ligands, and metal porphyrins.

Photocatalytic dehydrogenation in the presence of heteropolyacids. Upon illumination with the UV quanta, heteropolyacid H4[SiWl2040 ] provides with the quantum yield q~ = 0.2 (at ~. = 333 nm) and remarkable (ca. 100%) selectivity dehydrogenation of ethanol into acetaldehyde in water-ethanol mixtures [ 10]: hv, HPA C2HsOH

> CH3CHO + H2, H2

AG -- 20 kJ/mol

H4[SiWI2040 ]

(1)

hv, C2HsOH

(pc)

(H2PC) H6[SiW1204~

CH3CHO

fast PC+H2PC -,- "- HPC Photocatalytic reaction (1) follows the pathway presented in Scheme 1 [10]. It consists basically of two steps: I - two-electron reduction of the excited state of the photocatalyst PC* (i.e. H4[SiW12040]*) by C2HsOH to form H2PC (i.e. H6[SiW12040]) and CH3CHO, and II two-electron step of H2 evolution from H2PC. Step II is the rate determining one [ 10]. For step I further studies are needed to clarify whether it is a single two-electron reaction ( C2HsOH + PC hv > CH3CHO + H2PC) or a sequence of one-electron reactions, (e.g., C2HsOH + PC hv > CH3C'HOH + HPC,

C2HsOH +HPC

CHsC'HOH + PC hv , CH3CHO + HPC, CH3C'HOH + HPC

hv ; CH3C'HOH + H2PC, hv> CH3CHO + H2PC) [10]. .....

In contrast, step II is clearly a single two-electron reaction which proceeds with a rather low activation energy E a = 25 kJ/mol, but a high-negative activation entropy, ASa ~ - 200 J/K-mol. Note, that identical values were found for the rate constant kli of reaction II in four independent experiments, namely, by measuring: (1) - the rate of H2 evolution (with GC-method) in the overall photocatalytic reaction of Scheme 1, and (2), (3) and (4) - the rates of H 2 evolution (also using GC), disappearance of H2PC (using UV spectroscopy) and formation of HPC (using EPR) in the absence of illumination from/in solutions of PC that had

37 been prereduced to H2PC on Zn amalgam [10]. The value obtained for k n ~ 3.10 -4 sec 1 is rather small due to the high negative value for ASa. However, step II can be notably accelerated under the action of visible (including red) light. This light is known to excite the bands of intervalence charge transfer in reduced forms of HPAs, which makes the electron transfer from the HzPC frame to its internal protons more facile.

Photocatalysis and photogenerated catalysis with metal carbonyl, phosphine and other ligandso Excellent reviews of this field are given in [11-12]. Mechanistically, two types of photoinduced catalysis with the participation of transition metal complexes can be distinguished, namely, classical photocatalysis and photogenerated catalysis. The difference between these two types of catalysis are illustrated in Scheme 2. Scheme 2a presents a classical photocatalytic pathway for the water-gas-shift reaction in aqueous THF solutions (see [12] and refs. therein). For this pathway photon directly participates as a reagent in the catalytic cycle of olefine hydrogenation. Scheme 2b (see [11 ] and refs. therein) presents an example of photogenerated catalysis. Here photons do not participate in the catalytic cycle itself, but provide preparation of catalyst active form H2Fe(CO)3(olefine) from Fe(CO)5 precursor via photochemical dissociation of two CO ligands from the metal atom. CO CO2 _ hv ~,HFe(CO)4 ~ ~ HFe(CO)3 HFe(CO)4

--'72__

HFe(CO)4 + OH-

OH"

/

~1

"- 2Fe(CO)5 + Fe(CO)4 ~ . ~ hv

Fe(CO)5 ~

CO

H2Fe2(CO)7

2-~ Fe2(CO)7 2CO

H2

\~

Fe(CO4) ~ /

[I--Fe(CO)4 ~Qhv, H~ v~ CO H I II-Fe(CO)3

I Scheme 2b[

H H2 ~1-- Fe(CO)3

C3H8~

H I Fe(CO)3

~

\X/

38 Besides dissociation of ligands, photoexcitation of transition metal complexes can facilitate: (1) - oxidative addition to metal atoms of C-C, C-H, H-H, C-Hal, H-Si, C-O and CP moieties; (2) - reductive elimination reactions, forming C-C, C-H, H-H, C-Hal, Hal-Hal and H-Hal moieties; (3) - various rearrangements of atoms and chemical bonds in the coordination sphere of metal atoms, such as migratory insertion to C=C bonds, carbonyl and carbenes, c~- and [3-elimination, o~- and [3-cleavage of C-C bonds, coupling of various moieties and bonds, isomerizations, etc. (see [11, 12] and refs. therein). Such a broad range of classical elementary reactions of homogeneous catalysis with metal complexes, that can be facilitated by photons, make illumination of reaction solution a very useful instrument for substantial increase of the possibilities of homogeneous metal complex catalysis in organic synthesis. Particular examples of light-assisted syntheses will be given in section 3. Redox reactions with metal porphyrins (MPs) as photocatalysts. A spectacular example here is the reaction that couples upon illumination with the stmlight, methanol oxidation to formaldehyde with the formation of hydrogen peroxide in benzene-methanol mixture (90:10) hv, PC CH3OH + 02 > H2CO + H202 (2) in the presence of oxoalkoxoporphyrinatomolybdenum(V) as the photocatalyst [13]. Due to a large amount of energy provided by photons, some redox photocatalytic reaction can easily proceed at very low temperature. A characteristic example is the reaction between C C I 4 and triethylamine that,proceeds smoothly in the presence of porphyrinatozinc(II) and magnesium(II) photocatalysts even at 77 K in vitrified ethanol solutions (Scheme 3) [14]. Photoinduced electron transfer from Et3N to MP* and subsequent photoinduced electron transfer from M P to CC14 proceed at this temperature via electron tunneling, and the reaction products Et3N+ and CCI 3, Cl are separated in space [14].

Et 3 +

MP-

hv, CC14

[ Scheme3 ]

Et3N, h v "

\ MP

CC14

~- CC13 + C1-

2.2. Photocatalysis in organized molecular assemblies A prominent example here are photocatalytic systems based on lipid vesicles for watersplitting into H2 and 02. Design of such systems is based on both functional and structural mimicing of natural photosynthesis. Typically, photocatalytic cleavage of water in organized molecular assemblies (Scheme 4) is designed as a sequence of three key steps: charge separation, that is, formation of sufficiently strong oxidant (D § and reductant (A-) from intermediate electron acceptor A and electron donor D under the action of solar light in the presence of a photocatalyst PC, and subsequent catalytic reactions of oxygen evolution from water by oxidant D § and of hydrogen evolution by reductant A (see [6, 7, 15, 16] and refs. therein). The last two steps, to be

39

~ehv

Ill

I

AG~

eo

O

PC tpc+

T oxygen evolution

Q y charge

H20/H2 kJ + 89 = 237 mol

Q

Scheme 41

~r separation

hgdro0en evolution

accomplished, do not necessarily require light. The overall process is described by the equation 2 H 2 0 hv, PC, D ' A --, 2H 2 + 0 2 cat H2 ' cat 02

(3)

The major problem in accomplishing water splitting via the pathway of Scheme 4 is how to suppress the back recombination reaction D + + A ---> D + A, which is a simple exothermic bimolecular process and therefore typically proceeds much more rapidly than complex catalytic reactions of H2 and O2 evolution. An attractive way to overcome this problem is to use microheterogeneous photocatalytic systems based on lipid vesicles, i.e. microscopic spherical particles formed by closed lipid or surfactant bilayer membranes (Fig. 1) across which it is possible to perform vectorial photocatalytic electron transfer (PET). This leads to generation of energy-rich one-electron reductant A- and oxidant D+, separated by the membrane and, thus, unable to recombine. As a result of such PET reactions, the energy of photons is converted to the chemical energy of spatially separated one electron reductant and oxidant. C15H31 C15H31 O=C

C=O

I ~

t o

k -- H2C--- CH t O

i -

Clt2

i

O=P-OI

O

I CH2 I CH2 L N+

CH3 t CH3 CH3

Fig. 1. Structure of a lipid vesicle. As an example, let us consider a system, which contains a photocatalyst Ru(bipy)~§ in the inner volume of the vesicle, an electron carrier, e.g. cetylviologen (CI6V2+), in the membrane, and

40 an electron acceptor, Fe(CN)63- in the outer volume of the vesicle. Illumination into the PC absorption band induces electron transfer from Ru(bipy) 32. across the membrane to Fe(CN) 36 9The quanttma yield of the transmembrane electron transfer, % is related to the quanama yield of Ru(bipy) 33++ CI6V§ formation, % = 15%, as follows: q~ = q ) o k t / ( k t + kr) , w h e r e k t is the rate constant of electron transfer across the membrane by C16V§ radical and k, is the rate constant of recombination between Ru(bipy) 33. and Cl6V§ (see Fig. 2). The recombination and between Ru(bipy) 33, viologen radicals nearby the inner surface of the vesicle is fotmd to be strongly inhibited as compared to homogeneous solution, presumably due to a fast extraction of hydrophobic viologen radicals into the depth Ru(bipy)2+Gk;2+ ,P=,Po kt+k r of the membrane. Two mechanisms of transmembrane electron Fig.2. Mechanism of PET across a lipid membrane in the transfer were elucidated: (i) Ru(bipy) 32. _CV2+_Fe(CN) 63- system. via the translocation of viologen radical across membrane and (ii) via reaction of electron exchange between the radical and oxidized viologen located in the different monolayers of a bilayer membrane (see [16] and refs. therein). Note, that the dication C~6V2+ can not penetrate across the central core region of the membrane, presumably because of its unsufficient hydrophobicity. For various carriers of viologen type the rate constants k t and kr were measured, and the efficiencies b = k t / (k t + kr) of the transmembrane charge separation reaction were determined to be up to several percent. The rate constant k t was proved to be dependent on the substituent in viologen, the phase state (gel or liquid crystal) and the electrical polarization of vesicle membrane. Another promising way of transmembrane PET includes intramolecular electron transfer along bridge molecule D-PC-A which spans the bilayer and contains PC, D and A fragments linked by covalent bonds [17]. The membrane-separated reductant and oxidant formed upon PET can be used for accomplishment of various catalytic redox reactions which provide conversion of the chemical energy of a (D+...A-) pair into the chemical energy of a pair of more stable species such, e.g., as H2 and 02 molecules. This stored energy can be released when necessary in the form of high potential heat or electricity via combustion of H2 + 1/2 02 mixture in a furnace or fuel cell. It proves possible to anchor catalysts of H2 evolution to the outer and inner surface of the vesicle membrane. These catalysts are finely dispersed (10-20 A in diameter) metal Pt or Pd particles formed via reduction of appropriate salts in vesicle suspension (see [ 15, 16] and refs. therein). Among the viologen-type electron carriers a promising one is p-bis (1,2,5-triphenyl4-pyridil)benzene which possesses reduction potential low enough for water reduction at neutral pH. Recently, using this mediator we succeeded in H2 evolution conjugated with PET ,

41 across membrane in the "sacrificial" system which is schematically represented in Fig. 3 (V is an electron carrier, EDTA - ethylenediaminetetraacetic acid).

EDTAox

hv , Ru(bpy>~+C~ Ru(bpy)

"v"

EDTA~

Ru(bpy)3+

(Vv_lK::+

Fig.3. Mechanism of PET across a lipid membrane, coupled with H 2 evolution in EDTA-Ru(bipy) 32 _V2+_pd / lipid vesicle system. Photocatalyzed H2 evolution inside vesicle cavity over tiny anchored Pt particles also has been reported (see [15] and refs. therein). The membrane-bound catalyst for water oxidation to 02 can be prepared via oxidation of Mn(II) and Co(II) salts to Mn(IV) and Co(Ill) hydroxides, respectively, in the presence of lipid vesicles. Using these catalysts and photogenerated Ru(bipy)33. complex as an oxidant, it is possible to oxidize water to 02 in vesicle systems. One of such systems for 02 evolution is schematically represented in Fig. 4. Thus, vesicles serve as chemical 1 / 2 0 2 + 2H + hv | / 2 S 2 O a2" microreactors with transparent and electron-conducting walls. Molecular Ru(bpy)+ + engineering of such microreactors nowadays is one of the most fascinating areas of basic research in supramolecular photocatalysis. However, one should not SO~ underestimate the difficulties in OH 2 Ru(bpy)3+ designing vesicular catalytic reactors. The major problem comes from the Cat - (OH),Co.'" O+..z,; n = 100 fragility of the walls. Indeed, the so far available surfactant materials for the Fig.4. Photocatalytic $20 82--Ru(bipy) 32. walls are still not stable enough and - (OH)x-Co iu O(3n-x)n / lipid vesicle decompose rather easily upon n system for 02 evolution. illumination in the presence of strong

42 oxidants. Moreover, nonpolar small molecules, such, e.g., as 02 produced on cat o2 (see the notation in Scheme 4) attached to the outer surface of the membrane wall (Fig. 4), easily migrate through this wall. Contact with 02 destroys the system for H2 evolution on cat H2 attached to the inner wall of the microreactor. For this reason, the closed photocatalytic cycle of water splitting into H 2 and 02 has not so far been accomplished with vesicular systems, despite the fact that separately photocatalytic transmembrane electron transfer and H2 or 02 evolution, have been accomplished [ 15, 16]. Better isolation between the oxidative and reducfive parts of the photochemical system, perhaps, may be provided by solid layered structures, where two different interlayer compartments containing the catalysts for evolution of H 2 and 02, can be separated by a rigid layer of framework atoms [18]. It might also occur possible to avoid the inhibiting influence of 02 on the molecular apparatus for H 2 production, by means of evolving dihydrogen and dioxygen from water consecutively rather than simultaneously. In this case one may benefit from the technique [7], which provides the separation in time of the two steps: (i) photogeneration of a strong oxidant, which in this case must be conjugated with the simultaneous process of dihydrogen evolution, and (ii) subsequent oxidation of water with this oxidant. From what we know today about PET in biological and synthetic membrane or layered systems, we may expect that the non-biological apparatus providing photogeneration of spatially separated one-electron reductant and oxidant is likely to be developed in a rather universal way and may be expected to accomplish in the future not only water cleavage, but also various other redox reactions, such e.g., as photochemical synthesis of ammonia via the reaction N2 + 3H20 hv_~ 2NH3 + 3/202, photochemical synthesis of various organic compounds from CO2 and H20 (e.g., CO2 + 2H20 hv ~ CH3OH + 3/202 ), etc. Other examples of organized molecular assemblies of interest for photocatalysis are: (1) PC-A, PC-D or D-PC-A molecules where PC, A and D fragments are separated by rigid bridges; (2) host-guest complexes; (3) micelles and microemulsions; (4) surfactant monolayers or bilayers attached to solid surfaces, and (5) polyelectrolytes [19].

2.3. Photoeatalysis with semiconductors Semiconductor photocatalysts in a foma of colloids, powders, porous granules, thin films or bulk solids including single crystals (used in model studies) provide both liquid phase and gas phase transformations. Comprehensive reviews in this field can be found in monographs [4] (Chapters by N.S.Lewis and M.L.Rosenbluth; M.Gr/itzel; M.Schiavello and A.Sclafani; P.Pichat and J.-M.Herrmann; G.A.Somorjai; T.Sakata; H.Tributsch; M.A.Fox; H.A1-Ekabi and N.Serpone; D.F.Ollis, E.Pelizzetti and N.Serpone); [8] (Chapter by Yu.A.Gruzdkov, E.N.Savinov and V.N.Parmon) and [3]. The nature and general pathway of the photocatalytic action of semiconductors are nowadays well established (see Fig. 5), though numerous structural, thermodynamic and mechanistic peculiarities often make the detailed pathways for the same PC in different reactions or different PCs in the same reaction, also somewhat different [3, 4, 6, 8]. As seen from Fig. 5, upon absoption of photons with the energy hv > Eg, an electron | and hole Q centres are formed. They migrate to different sites on the PC surface, thus becoming spatially separated. Note, that what solid state physisists call surface electron and hole centers, in fact are some definite chemical species with strong reducing and oxidizing

43 properties, respectively. E.g., for TiO2 semiconductor surface electron and hole centers are suggested

to

be,

respectively,

{-Ti(III)-}surface and

{Ti(IV-O'- -Ti(IV)}surface

or

{Ti(IV)-O2-Ti(IV)} ()H sites [20]. Chemically pure semiconducor materials can absorb only those photons, the energy hv of which exceeds the band gap Eg. Therefore, Eg value determines the "red" boundary of the light that is used in photocatalytic action of these materials. By way of example, Table 1 presents the values of Eg and the corresponding values of boundary wave length 9~o= hc/Eg (where e is the velocity of light) for some semiconductor and dielectric oxides [2]. However, a semiconductor PC can be sensitized to light with ~ > ~o by Fig.5.The nature and general pathway of the chemical modifications of its surface photocatalytic action of a semiconductor layer or adsorption of certain catalyst particle. Eg is the band gap, E e shows molecules on its surface, provided that the direction of change of the energy for such treatment creates additional full electrons or empty electron levels in the band gap of the semiconductor material. Yet another approach to sensitizing PCs to a broader light spectrum is to use composite materials with a heterojunction (Fig. 6) between a narrow band gap and wide band gap semiconductors. A particular Table 1. photocatalyst of Fig. 6 smoothly provides the reaction of H2S Band gaps Eo and "red" boundaries Lo of the optical absorption f~3rtypical semiconductor and dielectric decomposition oxides [2] H2 S hv,PC ~ H 2 + S Oxide E z/eV* Lo/nm * To reduce Aad s the bottom of NiO 0.93 1340 the PC conduction band must be Cr203 1.4 890 located above the electron level of CuO 1.7 735 Aads, while to oxidize Dad s, the top CdO 2.1 595 of the valence band must be located below the level of Daas. If Fe203 2.2 570 a particular semiconductor PC TiO2 3.0 420 does not match these conditions, ZnO 3.2 390 one can shift the position of the MgO 7.2 178 conduction band up and that of SiO 2 8.6 145 the valence band down by just A1203 9.0 138 decreasing the size of the PC particle to ca. 10 2 ~ or smaller *The value refers to the pure bulk material and does not [21]. take into account the particle size and impurities.

44 The efficiency of semiconductor PCs in some reactions (such as dehydrogenation of organics, splitting of H20 and HES, etc.) can be enhanced by depositing tiny islands of additional catalysts, which facilitate certain reactions stages that may not require illumination. For example, islands of Pt metal are deposited on the surface of the composite photocatalyst in Fig.6 with the aim to facilitate the step of H2 formation. Fig.6. Photocatalytic cleavage of H2S over a The list of reactions provided by platinized composite sulfide semicondispersed semiconductors as photocatalysts ductor particle with a heterojunction. is very broad [22]. It includes: (1) splitting of water into the H2 and 02; (2) dehydrogenation of alcohols into H 2 and aldehydes or ketones; (3) oxidation of inorganic substrates (water into H202, NO2" into NO[, C N into CNO); (4) partial oxidation of organic substrates with 02 (benzene into phenol, toluene into methylphenols, - ~(CH2)4 ~ [--'( CH2)4 i .~

R

Ph Ph Ph O. Ph Ph Ph ph)= ~ a h ) - - O + ph)z---x + ph~CHCHO " ph)CH2 ~ p h ) - - O , etc.); (5) complete oxidation (mineralization) of hydrocarbons, oxigenated and chlorinated hydrocarbons, amines, nitrocompounds; (6) reduction of CO2 with H20 into CH3OH and C2HsOH; (7) nitrogen fixation: N 2 + 3H20 -~ 2NH 3 + 3/202, (8) syntheses of amines, e.g. NH 3 + 3CHaOH -}(CH3)3N + 3H20; (9) synthesis of aminoacids, e.g. NH 3 + PhCH2COCOOH --} PhCH2CHNH2COOH; (10) oligomerization and polymerization, e.g.

(11 ) structural isomerization of hydrocarbons, e.g.

(12) hydrogenation and cracking of alkenes and alkines with H20 or alkohols, e.g. C3H6 + H20 --~ CH4, CIH6, C3H8; (13) deposition of metals on semiconductors, e.g., 4Au 3§ + 6H20 --~ 4Au ~ + 302 + 12H § For some of these reactions detailed mechanistic studies were carried out. As an example, in Scheme 5 the pathway suggested for 1,1-diphenylethylene oxidation with 02 into benzophenon [23] is shown.

45

TiO2 hv _ _ @ + @ ;

Ph

@+02

~O2

02 (or 0~)

Ph

I Scheme 51 r Ph

ph>--- + (~

Ph

O~O"

TiO2 Ph P

~

+

"x

_@ I -" O--O

P

,A

O--O

I

! i_

Adsorption/desorption [24-25], as well as mass transfer phenomena [26] were also proved to be important for photocatalysis. An important advatage of semiconductor PCs compared to molecular ones, is a far greater stability of the former PCs. This is particular true for TiO2 photocatalyst. However, a notable disadvantage of TiO2 and most of other stable semiconductor PCs is that they have rather large band gaps Eg, and thus, are sensitive to only UV light. 3. PHOTOCATALYSIS IN ORGANIC SYNTHESIS The lists of reactions provided by molecular, supramolecular and semiconductor PCs

(vide supra) suggest certain perspectives for their practical applications for fine organic synthesis. Note, that sometimes PCs act selectively producing a sole product in rather exotic reactions; in other cases, a whole set of exotic products can be produced from the same reagents (vide supra). As additional examples note formation of glycols from alkohols ZnS (CH3OH > HOCH2-CH2OH and ZnS or (CH3)3OH > (CHs)zC(OH)CHzCHzC(OH)(CH3)2) and RuO2/TiO2 ZnS diamines from amines ((C2Hs)3N > CH3CH HCCH 3 [27]), as well synthesis of 1~(C2H5)2 J(C2Hs)2 aminoacids from hydroxocarboxylic acids (Scheme 6) or ketocarboxylic acids in ammoniawater solutions on CdS, ZnS and TiO2, or CdS and ZnTSPP (TSPP is the tetrasulfonatophenylporphyrinato ligand), respectively, with remarkably high quantum yields upto 35%. The aminoacids Ala, Gly, Phe, Leu and Tyr are produced efficiently in this way [27].

RCH(OH)COOH

Scheme 6 ]

hv,PC / ~ " ' - " ~

2H" ~ ~ . ~ . ~

RCOCOOH NH3

RCH(NH2)COOH

RC(=NH)COOH+ H20

46 Scheme 7 illustrates the transformations of a bicycloheptadiene in the presence of transition metal complexes as photocatalysts or photogenerated catalysts [ 11 ]:

,••,•

v

~hv,

Ni(CO)4/ hv /

Fe(CO)5 hv

.f.

Cr(CO)~= \ \(Ph3 P)2Ni(CO)2

+

Thus, photocatalysis and photogenerated catalysis indeed open up rather reach opportunities in fine organic synthesis, including some new reactions and nontraditional pathways for some known reactions. More efforts should be made in engineering of appropriate photocatalytic reactors for such synthesis. 4. PHOTOCATALYSIS IN ABATEMENT OF ENVIRONMENTAL PROBLEMS Photocatalysts demonstrate a remarkable ability to provide at room temperature and atmospheric pressure, complete oxidation (mineralization) of both organic (hydrocarbons and their derivatives containing O ' N, S, Cl and other heteroatoms) and inorganic (S 2-, SO 23

'

NO ~, CN) pollutants, remaining stable enough under the reaction conditions. The latter statement refers first of all to TiQ. All this make photocatalysis attractive for future commercial scale treatment of gaseous and aqueous waste streams, as well as hydrocarbon spills [3]. Intensive work in this field of applied photocatalysis is underway in various countries. Pilot plants have been tested with both electrically fed lamps and solar light (nonconcentrated or concentrated with mirrors). Their schemes and photographs can be found in monograph [3]. Economic feasibility studies suggest that even at the present state of the art photocatalytic technology indeed can be competitive with the traditional carbon adsorption or incineration technologies in treatment of contaminated soil vapor extraction vents and small scale VOCcontaining vents [28]. Rapid progress in basic and applied research in photocatalysis suggests

47 that in the near enough future photocatalytic technologies can be substantially improved and can become competitive with the traditional technologies also in other applications for environmental control. 5. PHOTOCATALYSIS AND THERMAL CATALYSIS INDUCED BY SOLAR RADIATION IN UTILIZATION OF SOLAR ENERGY At present, thermal catalysis induced by solar radiation is more ready for potential practical use in the energy production industry of the-future, than photocatalysis [7,9,29,30]. Fig.7 illustrates the scheme of the pilot plant for thermocatalytic solar-to-chemical energy

.i. N

/

i ••

.---

\

/~Storage vessel~ ~ for m e t h a n e ) / t
Solar Catalytic Reactor for methane Heat to recuperati consumer

/"~'~#.

~

!

Catalytic Reactor for Methanation of syn-gas

consumer

[ Storage vessel ~ for syn-gas

Fig.7. Scheme of thermal catalytic conversion and utilization of solar energy based on reversible methane reforming / methanation reactions. conversion and utilization, which is based on a closed loop sequence of the strongly endothermic reaction of methane reforming CH4 + H20 --> CO + 3H2 and subsequent strongly exothermic methanation reaction CO + 3H2 --~ CH4 + H20. Methane reforming reaction is accomplished under the action of heat collected from mirror concentrator of solar light. The mixture of CO and H2 produced in this reaction can be stored and then, when necessary, converted into high-potential heat (with the temperature up to 950 K) in the methanation catalytic reactor. The efficiency of solar-tochemical energy conversion as high as

Fig.8. Pilot plant for solar energy conversion based on the scheme of Fig. 7.

48 43% (with respect to the stored enthalpy) was experimentally achieved for methane reforming reaction in the pilot solar catalytic plant with useful power ca. 2kW [7]. Fig.8 shows the photograph of this pilot plant. As shown in [30], the maximum achievable efficiency of solar-to-chemical energy conversion (calculated with respect to Gibbs free energy) is ca. 60% for the thermocatalytic method and ca. 30% for the photocatalytic method. The actual progress made to date in developing workable converters for the thermochemical catalytic method exceeds the progress made in developing photocatalytic converters. However, we do not suggest that photocatalytic method should be discarded, and priority by given solely to the thermochemical catalytic method. A more careful scrutiny shows that each method has both advantages and disadvantages. For instance, an efficient thermochemical conversion necessitates concentrated sunlight. At present, only the direct solar light can be concentrated; the scattered light cannot. It is worth noting here that since the earth's atmosphere contains a multitude of aerosols, the proportion of scattered light in the overall solar energy flux falling on the surface of the Earth may be several tens percent. Thermochemical converters are not designed at present to utilize this scattered light. By contrast, photocatalytic conversion can be equally efficient with both direct and scattered sunlight. Moreover, photocatalytic conversion can proceed at ambient temperatures, while the thermochemical conversion must proceed at elevated temperatures to reach the high efficiencies. This makes the photocatalytic method more advantageous in obtaining energy carriers, such as H2, from industrial wastewater and from nontraditional raw materials (e.g., H2S dissolved in sea water). The advances made so far in thermocatalytic conversion of solar energy, suggest that this method may, perhaps, become important for energy production in the XXI century. Whether this actually happens or not, will depend primarily upon two factors. The first factor is of economic character. At present all the existing methods of energy production via conversion of solar energy are far more expensive than the traditional ones, such as combustion of fossil fuels at power plants, nuclear fission, etc. In this situation the shift towards a wide use of solar energy is expected to occur gradually, following the trends of solar energy becoming cheaper with every new scientific and technological breakthrough, and traditional energy sources becoming more expensive due to their gradual exhaustion. The second factor relates to environmental issues. Much will depend on how dangerous will actually be global consequences of Earth pollution with manmade extra heat, chemicals, etc., associated with traditional types of energy production. Note, that nuclear fusion, which sooner or later is anticipated to be developed, also is expected to pollute Earth with extra heat. If such pollution occurs intolerable, the development and commercialization of solar power plants, which produce no extra heating of the Earth and in other respects also seem to be environmentally friendly, may obtain high priorities. 6. H E T E R O G E N E O U S CATALYSIS IN THE TROPOSPHERE: A POSSIBLE

I M P A C T ON THE GLOBAL CHEMISTRY OF THE A T M O S P H E R E Direct photochemical reactions in the stratosphere induced by far ultraviolet solar light are well known to be important for the chemistry of the atmosphere. However, only a small fraction of the solar energy flux is in this spectral region.

49 Recently was estimated an expected impact on the global chemistry of the atmosphere of the indirect heterogeneous photocatalytic reactions under the much more abundant near ultraviolet, visible and near infrared solar light [2]. As photocatalysts may serve atmospheric aerosols, i.e. ultrasmall solid particles that sometimes are embedded into liquid droplets. Aerosols are known to contain TiO2, Fe203, ZnO and other natural oxides, as well as metal sulfides of volcanic or antropogenic origin, that may serve as semiconductor photocatalysts (see Fig.5). Aerosols are known to be concentrated mainly in the air layers near the surface of the Earth, i.e. in the troposphere, rather than stratosphere. Direct experimental studies of photocatalysis and photoadsorption over real atmospheric aerosols are just starting. Therefore, today the estimates of their expected role in the global and local chemistry of the atmosphere can be made only on the ground of numerous in vitro studies of the reactions over certain model photocatalysts (first of all, TiO2, Fe203 and ZnO) and photoadsorbents (e.g., MgO [31 ]), and some known characteristics of the atmosphere as a global photocatalytic reactor (such as the average concentration of continental aerosols in the troposphere, their distribution over the altitude, average content of TiOz, Fe203 and ZnO in them; the spectrum and intensity of solar radiation in the atmosphere; the spectral and photocatalytic properties of semiconductor photocatalysts; the average natural concentrations of the main and trace gases in the atmosphere, their average residence time in the atmosphere and rates of removal (supply) from (to) the atmosphere) [2]. Estimates made in [2] suggest that the impact of photocatalytic reactions over the troposphere aerosols on the global chemistry of the atmosphere may be non-negligible. Such reactions are anticipated to affect the intensity of acid rains, concentrations of green-house and ozone-depleting gases, as well as clean the atmosphere and soil from harmful compounds, including oil spills. Thus, desert areas and volcanos, where a lot of continental dust is generated, may serve as "kidneys" of the Earth. 7. A C K N O W L E D G E M E N T The author is grateful to his coworkers in the area of photocatalysis from the Boreskov Institute of Catalysis: V.N. Parmon, V.I. Anikeev, Yu.I. Aristov, G.L. Elizarova, Yu.A. Gruzdkov, M.I. Khramov, V.A. Kirillov, S.P. Lymar, L.G. Matvienko, E.N. Savinov, E.R. Savinova, V.S. Zakharenko for their contributions to the studies at the Institute reviewed in this paper and fruitful discussions. 8. REFERENCES 1. 2. 3. 4. 5. 6.

Govindjee (ed.), Photosynthesis, Vols. 1 and 2, Academic Press, New York, 1982. K.I. Zamaraev, M.I. Khramov and V.N. Parmon, Catal. Rev.-Sci. Eng., 36 (1994) 617. D.F. Ollis and H. A1-Ekabi (eds.), Photocatalytic Purification and Treatment of Water and Air, Elsevier, Amsterdam, 1993. N. Serpone and E.Pelizzetti (eds.), Photocatalysis. Fundamentals and Applications. Wiley, New York, 1989. K.I. Zamaraev and V.N. Parmon, Catal. Rev.-Sci. Eng., 22 (1980) 261. M.Gr/atsel (ed.) Energy Resources through Photochemistry and Catalysis, Academic Press Inc. New York, 1983.

50 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.

K.I. Zamaraev and V.N. Parmon, in E. Pelizzetti and M. Schiavello (eds.), Photochemical Conversion and Storage of Solar Energy, Kluwer, Netherlands, 1991, p. 393. K.I. Zamaraev and V.N. Parmon (eds.), Photocatalytic Conversion of Solar Energy, Vols. 1 and 2. Nauka, Novosibirsk, 1991 (in Russian). K.I.Zamaraev, V.N. Parmon and Z.R. Ismagilov, in D. Behrens (ed.), Strategies 2000, 4th World Cong. of Chem. Eng., Karlsruhe (1991), DECHEMA, Frankfurt am Main, 1992, p. 49. K.I. Zamaraev and V.N. Parmon in [6], p. 123; E.N.Savinov, S.S.Saidkhanov, V.N.Parmon and K.I.Zamaraev, React.Kinet.Catal.Lett., 17 (1981) 407; E.N.Savinov, S.S.Saidkhanov and V.N.Parmon, Kinet.Katal., 24 (1983) 68; E.N.Savinov, S.S.Saidkhanov, V.N.Parmon and K.I.Zamaraev, Dokl.Akad.Nauk SSSR, 272 (1983) 916. F. Chanon and M. Chanon, in [4], p. 489. P.C. Ford and A.F. Friedman, in [4], p. 541. H.J. Ledon and M. Bonnet, J. Am Chem. Soc., 103 (1981) 6209. R.F.Khairutdinov, K.I.Zamaraev and V.P.Zhdanov. Electron Tunnelig in Chemistry. Chemical Reactions over Large Distances. in: Comprehensive Chemical Kinetics. R.G.Compton (ed.). v.30, 1989. V.A.Shafirovich, S.V.Lymar, V.N.Parmon, A.E.Shilov and K.I.Zamaraev, in [8], p.18 (in Russian). S.V.Lymar, V.N.Parmon and K.I.Zamaraev, Topics in Current Chemistry, 159 (1991) 1. P.Seta, E.Bienvenue, A.L.Moore, P.Mathis, R.V.Bensasson, P.Liddel, P.J.Pessiki, A.Joy, T.A.Moore, D.Dust, Nature, 316 (1985) 653. K.Domen, K.Kudo, A.Shinozaki, A.Tanaka, K.Maruya and T.Onishi, J.Chem.Soc., Chem.Commun., 356 (1986); A.Kudo, A.Tanaka, K.Domen, K.Maruya, K.Aika and T.Onishi, J.Catal., 111 (1988) 67. D.Meisel and M.S.Matheson, in [4] p.385. N.Serpone, E.Pelizzetti and H.Hidaka, in [3], p.225. A.J.Nozik, in [3], p.39. Yu.A.Gruzdkov, E.N.Savinov and V.N.Parmon, in [8], vol.1, p.138. M.A.Fox, in [4], p.421. P.Pichat and J.M.Herrmann, in [4], p.217. H.AI-Ekabi and N.Serpone, in [4], p.457. A.Sclafani, A.Brucato and L.Pizzuti, in [3], p.533. T.Sakata, in [4], p.311. E.Miller and R.Fox, in [4], p.573. Yu.l.Aristov and V.N.Parmon, in [8], vol.1, p.315. V.N.Parmon and K.I.Zamaraev, in [4], p.565. V.S.Zakharenko, V.N.Pan~on and K.I.Zamaraev, Kinet.Katal., 1996 (in press).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

51

T o w a r d s M o l e c u l a r Design of Solid Catalysts Alfons Baiker Department of Chemical Engineering and Industrial Chemistry, Swiss Federal Institute of Technology, ETH-Zentrum, CH-8092 Ziirich

1. INTRODUCTION Rigorous molecular design, as opposed to empirical design of a catalyst, requires not only some understanding of the reaction mechanism, but also knowledge and control of the crucial structural and chemical properties of the solid catalyst. Major advances in several fields essential for catalysis, such as surface analytical instrumentation [ 1], surface science [2], organometaUic chemistry [3], theoretical techniques [4], solid state chemistry [5], material science [6], and reaction engineering [7], have brought the molecular design of solid catalysts within reach in certain cases. Progress in analytical techniques has made it possible to characterize catalyst surfaces under reaction or near reaction conditions and to reveal the nature of adsorbed surface species and their concentration. Dynamic techniques and isotope tracers allow to distinguish between catalytically important and spectator species. Surface science has greatly contributed to molecular-level understanding of the elementary processes on which catalytic reactions are based. Important relationships between structure and activity have been uncovered and confirmed on industrial catalysts, leading to detailed mechanistic insight of some reactions. Advances in synthetic organic chemistry and organometallic chemistry have broadened the scope and relevance of homogeneous catalysis and opened new opportunities in the design of catalytic processes for the production of fine chemicals. Quantum chemical techniques on both ab initio and semiempirical level have been developed to a valuable tool aiding in understanding molecular interactions relevant to the catalytic process and catalyst design. Progress in solid state chemistry provides new opportunities for synthesizing structurally and chemically tailored materials. All these advances have led to a drastic reduction in the level of empiricism in catalyst design and moved it closer towards the ultimate goal of molecular design. It is not surprising that progress towards this aim has been most prominent with solid materials whose intrinsic properties are amenable to control on the molecular level [8]. Examples are zeolites [9] and molecular sieves [10], heteropoly acids [11] and perovskites [12]. However, significant advances have also been made with supported metals [13] and metal oxide catalysts [14]. Here we shall illustrate and discuss the opportunities provided by two design concepts, which have undergone crucial development in the past years. These are the solution-sol-gel method combined with supercritical drying fo, tailoring mixed oxides and the design of supported metals modified with adsorbed auxiliaries for controlling the chemoand stereoselectivity of metal-catalyzed reactions. 2. SOL-GEL DERIVED MIXED OXIDES Mixing metal oxides on atomic scale provides a unique opportunity for tailoring the chemical and structural properties of oxide catalysts. A versatile method for achieving intimate mixing is the solution-sol-gel method (SSG) [15] which represents a highly controllable preparation route with inherent advantages, such as molecular scale mixing of the constituents, purity of the precursors, homogeneity of the sol-gel product (isotropy), and the

52 use of different wet-chemical preparation tailoring tools. Generally, the SSG method proceeds from a molecular-level precursor solution via a system of colloidally dispersed particles (sol) to a disordered, branched, and interconnected network (gel). Depending on the nature of the precursor, one can also term it inorganic oxidic (or organic) polymerization or network formation. The main wet-chemical reactions involved are hydrolysis (formation of hydroxyl groups - hydroxylation) and (poly)condensation (formation of bridging bonds -oxolation and olation) of the appropriate precursor molecules. Three principal routes to sol-gel systems are followed, depending on whether the precursor is an inorganic salt in aqueous solution, a solid aggregate consisting of colloidal subunits, or a network forming agent in mainly organic solution. The SSG method with network forming agents is most widely used in catalyst preparation due to its superior potential for controlled processing. It is based on the application of various organometallic compounds, among which metal alkoxides are most frequently used. Important SSG preparation parameters used to control and tailor the chemical and textural properties of solids include type of precursor, type of solvent, water content, acid or base content, precursor concentration, and temperature. These parameters affect the structure of the initial gel and, in turn, the properties of the material at all subsequent processing steps. Another important parameter is the aging, i.e. the time between the formation of a gel (solid matrix encapsulating a solvent) and its drying (removal of solvent). Aging is a useful step consolidating the wet-chemical structure by continuous condensation, esterification, dissolution, reprecipitation, depolymerization, and repolymerization. The processes concerned are Ostwald ripening, coalescence-coarsening, sintering and syneresis (network densification by continuous condensation reactions or attraction between agglomerates). A crucial step affecting the chemical and structural properties of SSG-derived solids is the drying. Depending on the drying method applied the final solids are termed xerogels (conventionally dried), cryogels (freeze-dried) or aerogels (supercritically-dried). Conventional evaporative drying, such as heating a gel in an oven, induces capillary pressure associated with the liquid-vapor interface within the pores, causing differential macroscopic and microscopic shrinkage and consequently severe cracking of the tenuous gel structure. Supercritical drying (SCD) is a suitable procedure for reducing the differential capillary stresses which result in such drastic structural rearrangements. Two methods of SCD for eliminating any liquid-vapor interfaces during the solvent removal can be distinguished: the high-temperature method, where capillary stress is avoided by transferring the solvent into the supercritical state (alcoholic solvents Tc > 510 K), and the low-temperature method which is based on extraction of the solvent using supercritical CO2 (Tc = 304 K). An interesting recent example of successful application of the SSG process combined with ensuing supercritical drying is the design of titania-silica mixed oxides for the epoxidation of bulky olefins [16-18]. This example will be used to illustrate the opportunities the combined use of SSG and SCD provide for tailoring the chemical and structural properties of mixed oxides. The epoxidation of olefins has attracted considerable interest due to the versatility of epoxides as intermediates in organic synthesis. The first truly heterogeneous catalyst (TiO2on-SiO2), discovered by Shell researchers [19], is used in industrial scale for propylene epoxidation. It is assumed that the active species of this catalyst are the surface Tialkylperoxo groups, isolated by O-Si(Iv) fragments [20]. The chemical stability of the catalyst and the absence of leaching problems are attributed to the formation of Si-O-Ti bonds between the Ti-precursor and the silica support during calcination. Other silica-supported oxides, such as V205, MoO3 and ZrO2, were found to be inferior to TiO2-on-SiO2, due to either lower activity or chemical stability [21 ]. The discovery of titanium-substituted zeolites with MFI (TS-I) and MEL (TS-2) structures opened a new direction in the development of heterogeneous oxidation catalysts [22-24]. In this titanium silicalites titanium is assumed to be uniformly distributed in the crystalline framework by isomorphous substitution of a part of silicon with titanium [23]. Unfortunately, the use of these catalysts is limited to relatively small reactants which are able to penetrate into the channels of ca. 0.55 nm. Considerable

53 effort has been undertaken in the past years to overcome this restriction by synthesizing largeand ultra large-pore titanium-containing zeolites isomorphous to zeolite 15 [25] or molecular sieve MCM-41 [26]. For bulky reactants these materials are superior to titanium silicalites, but their intrinsic oxidation activity for C6-C10 aliphatic alkenes is considerably lower than that of titanium silicalite [27]. Another problem is that in the oxidation of higher alkenes, complications arise due to the immiscibility of the reactant and the aqueous hydrogen peroxide solution. From the knowledge gathered on the various known titania-silica epoxidation catalysts, two crucial requirements emerge: the importance of Si-O-Ti linkages as active sites, and the mesoporosity providing accessibility of the bulky olefins to the active surface. These properties can be influenced at various stages of the SSG process, i.e. by various preparation parameters, such as reactivity of Ti- and Si alkoxides, concentrations of alkoxides, type of solvent, type and concentration of hydrolysant, temperature, aging, drying and calcination. Systematic studies of the influence of these parameters on the final chemical and structural properties provided the basis for the controlled preparation of mesoporous titania-silica aerogels with a high density of heteronuclear Si-O-Ti connectivity [16]. Generally, the positive charge of Ti in the alkoxide precursor is significantly higher than that of Si in related alkoxides [28], which enhances the sol-gel activity of Ti over that of Si [29]. This property promotes an undesired "core-shell" structure, with titania forming the cores. In addition, the homocondensation rate of HO-Si(OR)3 species is significantly lower than the heterocondensation rate with RO-Ti(OR)3 [30]. A further point, which has to be taken into account, is that the titanium alkoxide catalyzes the condensation of silanol groups [31]. Chemical modification of the titanium alkoxide by chelate ligands, acid-catalyzed hydrolysis (stability of TiOH and SiOH groups in acidic medium against ensuing undesired homocondensation [29]), and the use of the most reactive silicon alkoxide (tetramethoxysilicon(IV)) are favorable to reach high density of the desired heteronuclear SiO-Ti connectivity. The drying procedure is a crucial parameter influencing both the abundance of Si-O-Ti connectivity and the pore structure (pore size). Titania-silica aerogels, possessing both high density of Si-O-Ti connectivities and mesoporous structure, can only be prepared by combining the optimized sol-gel procedure with ensuing low-temperature supercritical drying. High-temperature supercritical drying affords meso- to macroporous aerogels, but Ti segregation leads to the formation of undesired Ti-O-Ti connectivity (anatase domains) and consequently low Ti dispersion. Conventional evaporative drying results in microporous xerogels with a density of Si-O-Ti linkages comparable to that obtained by low-temperature SCD [ 161. Both the relative proportion of Si-O-Ti structural units (determined by FTIR analysis [ 16]), and the pore size are influenced by the Ti concentration as illustrated in Fig. 1 for titania-silica aerogels derived by low-temperature SCD. A titania content of ca. 20 wt. % proved to be most suitable for achieving these properties [ 17,32]. Note that the curve for the activity (initial rate) of the epoxidation of et-isophorone runs nearly parallel to that representing the proportion of Si-O-Ti connectivities, which indicates that these structural units can be associated with the active centers. In contrast to the rate, the selectivity to the epoxide (ca. 100%) does not depend on the Ti-content in the concentration range shown. Similar behavior was observed for the epoxidation of several other olefins, including, cyclododecene, norbomene, limonene, and a-isophorone [17,18]. In these epoxidations 76-100% peroxide conversion was obtained in 1 h with only 1-2 wt% catalyst/olefin ratio. For comparison, TiMCM-41, which has been found to be more active than Ti-[3 and TS-1, showed only 30% peroxide conversion after 5 h in the oxidation of cyclododecene, despite the extreme catalyst/olefin ratio (> 100 wt% ) which was used [26]. Figure 2 provides a direct comparison of the activities and selectivities achieved with various Ti- and Si containing catalysts in the epoxidation of a-isophorone [18]. The titania-silica aerogel with 20 wt% TiO2 has outstanding activity and selectivity. Inspite of the fact that the corresponding xerogel possesses a similar proportion of Si-O-Ti connectivities, its activity is poor due to its

54 microporous structure (mean pore size ca. 0.3 nm as compared to the size of ct-isophorone of 0.75 x 0.58 x 0.48 nm).

Figure 1. Influence of TiO2 content of LTaerogels on: relative proportion of Si-O-Ti connectivities R = [Si-O-Ti]/[Si-O-Si], mean pore diameter, and initial rate (ro) of ct-isophorone epoxidation with t-butyl hydroperoxide at 60~ Data taken from ref. [18].

Figure 2. Comparison of initial rates and selectivities (related to olefin and peroxide) for the epoxidation of ot-isophorone with t-butyl hydroperoxide at 6 0 o c over (a) lowtemperature aerogel (LT) (20 wt% TiO2), (b) titania-on-silica, (c) TS-1 and (d) xerogel.

3. METALS MODIFIED BY ADSORBED AUXILIARIES

The principle aim pursued with modifying supported metal catalysts is to improve their activity, selectivity or life time [13]. The general strategy is to deposit deliberately foreign species on the metal surface, which change either the geometry (ensemble effect), local electronic properties, or bring in additional catalytic functions (multi-functional catalysis). Geometrical changes in the active metal matrix, which can be beneficial with structuresensitive reactions, are achieved by blocking of metal surface sites by foreign species, which are either deposited inert foreign metals or strongly adsorbed species (selective poisoning). Less well understood is the modification of the local electronic properties of metal surfaces, bimetallic alloys represent an exception [13]. A possible way of changing the local electronic properties is the selective anchoring of organometallic complexes, which also allows to insert other catalytic functions into the metal surface [33]. A common feature of all the above concepts, used for modifying metal catalysts, is that the foreign species are either metals or organometallic complexes which are immobilized on the metal surface during the catalyst preparation. Here we shall consider a different concept, which has an interesting potential, particularly in liquid phase reactions used for the production of fine chemicals. The concept is schematically illustrated in Fig. 3. The modification of the metal catalysts is achieved by very small quantities (usually a sub-monolayer) of adsorbed auxiliaries (modifiers), which are either simply added to the reaction mixture (in-situ), or brought onto the catalyst surface in a

55

|

+

Metal

+

Figure 3. Principle of modifying supported metal catalysts by co-adsorbed auxiliaries

special pretreatment step before reaction (ex-situ). The auxiliary is adsorbed on the active metal surface and by interacting with the reactant(s), it can control the chemoand/or stereoselectivity of the catalytic process. Thus knowledge concerning the adsorption (anchoring) of the auxiliary on the metal surface and its interaction with the reactant(s) are of prime importance for designing this class of catalysts. Various types of functions can be introduced to metal-catalysis by the auxiliary, depending on its nature.

3.1 Stereoseleetive modification

Among the various strategies [34] used for designing enantioselective heterogeneous catalysts, the modification of metal surfaces by chiral auxiliaries (modifiers) is an attractive concept. However, only two efficient and technically relevant enantioselective processes based on this principle have been reported so far: the hydrogenation of functionalized 15ketoesters and 2-alkanons with nickel catalysts modified by tartaric acid [35], and the hydrogenation of a-ketoesters on platinum using cinchona alkaloids [36] as chiral modifiers (scheme 1). O

OH

5% Pt/AI203, H2.._

RI~O"R2 O

1

R1

R2

O (R)-ct-hydroxyester

ct-ketoester R R1 = C6H5CH2CH2, CH3 R2 = CH3, CH2CH3

Cinchonidine (X=OH, R=C2H3) lb : 10,11-Dihydrocinchonidine(X=OH, R=C2H5) 1r 10,11-Dihydro-O-methylcinchonidine(X=OCH3, R=C2H5) l a :

1

Scheme 1 Enantiomeric excesses (ee) as high as 85-95% have been achieved with these catalytic systems and considerable effort has been undertaken to gain a molecular understanding of their functioning [34,37,38]. Intrigued by the remarkable properties of the platinum-cinchona system, and by the fact that since its discovery in the late seventies, no new chiral catalysts with similar efficiency have been reported, we started recently searching for new enantiodifferentiating auxiliaries which could be used to substitute the cinchona alkaloids for chirally modifying platinum. It was hoped that structurally much simpler modifiers can be found, which are more amenable to structural tailoring, a basic requirement to broaden the scope of

56 this catalytic system. The knowledge gathered concerning the crucial parameters of the platinum-cinchona system and the mechanism of the chiral hydrogenation of a-ketoesters [37,38] formed the basis for this endeavor. The efficiency of the platinum-chinchona system for the enantioselective hydrogenation of a-ketoesters depends mainly on the structure and concentration of the modifier, the structural properties of the supported platinum, and the solvent used. Although all these properties have to be optimized to achieve high optical yields, the structure of the modifier is most crucial for enantioselection. Already very small quantities of modifier (modifier : Ptsurf ~ 1) are sufficient to induce both enantio-differentiation and rate acceleration. Systematic studies on the relationship between modifier structure and enantio-differentiation provided information on the structural requirements of modifiers [39]. Changing the absolute configuration at C-8 and C-9 of cinchonidine (scheme 1), i.e. substituting cinchonidine by its near-enantiomer cinchonine, alters the chirality of the product lactate. Alkylation of the quinuclidine nitrogen results in a complete loss of enantio-differentiation, which indicates that this center plays a crucial role in the mechanism of enantioselection. Partial hydrogenation of the quinoline ring causes a drop in ee to below 50%. The selectivity is hardly influenced by O-methylation, whereas replacing the OH by hydrogen or using the acylated derivative results in a decrease in ee of more than 20%. Interestingly, protonation of the quinuclidine nitrogen increases the ee by 5-15%. The three structural elements which are crucial for the functioning of the cinchona alkaloids as chiral modifiers are: (i) the anchoring part, represented by the flat aromatic ring system (quinoline), which is adsorbed on the platinum surface via the ~t-bonding system, (ii) the stereogenic region, embracing C-9 and C-8, which detem~nes the chirality of the product; and, (iii) the tertiary quinuclidine nitrogen, which is directly involved in the interaction with the reactant ct-ketoester. Theoretical studies aimed at rationalizing the interaction between the chiral modifier and the pyruvate have been undertaken using quantum chemistry techniques, at both ab initio and semi-empirical levels, and molecular mechanics. The studies were based on the experimental observation that the quinuclidine nitrogen is the main interaction center between cinchonidine and the reactant pyruvate. This center can either act as a nucleophile or after protonation (protic solvent) as an electrophile. In a first step, NH3 and NH4 + have been used as models of this reaction center, and the optimal structures and complexation energies of the pyruvate with NH3 and NH4 § respectively, were calculated [40]. The pyruvate--NH4 + complex was found to be much more stable (by 25 kcal/mol) due to favorable electrostatic interaction, indicating that in acidic solvents the protonated cinchonidine will interact with the pyruvate. The possible conformations of the pyruvate--protonated cinchonidine system were investigated using molecular mechanics [41]. Figure 4 shows the side and top views of the most stable complex formed on interaction of protonated cinchonidine and methyl pyruvate, which upon hydrogenation would yield R-(+)-methyl lactate and S-(-) methyl lactate, respectively. The complexes are accommodated on a platinum (111) surface in order to illustrate the space requirements of the adsorbed complexes. Note that the precursor to R-(+)methyl lactate can be adsorbed in a planar 7t-bonding mode on the Pt surface via the aromatic quinoline ring, without hindering the interaction of the carbonyl groups of methyl pyruvate with the Pt surface. This adsorption mode is sterically hindered for the complex suggested to be the precursor to S-(-)-methyl lactate. The opposite behavior is found when the complexes formed upon interaction of protonated cinchonine (the near-enantiomer of cinchonidine) with methyl pyruvate are optimized [41 ]. The precursor complex resulting in S-(-)-methyl lactate upon hydrogenation can be adsorbed without significant steric hindrance, while the one affording R-(+)-methyl lactate is sterically hindered. Thus calculations showed that - in agreement with the experimental observations - a change in the chirality of the stereogenic region (C-8, C-9) of the cinchona alkaloid results in a corresponding change in the chirality of the product lactate. In the transition complex the substrate is bound to the modifier via N-H-O interaction of the protonated quinuclidine nitrogen and the oxygen of the ct-carbonyl of pyruvate. The complex resembles a half-hydrogenated state of the reactant [42].

57

R

$

Figure 4. Side and top views of the energetically most favorable complexes formed between protonated cinchonidine and methyl pyruvate which would yield (R)-methyl lactate (left) and (S)-methyl lactate (right), respectively, upon hydrogenation. The complexes have been accomodated on a space filling model of platinum (111) surface in order to illustrate the space requirements of the adsorbed complexes. For the sake of clarity, in the side views the carbon atoms of the reactant are marked with a white square and the oxygen atoms with an o. Data taken from ref. [41 ]. The molecular modelling approach, taking into account the pyruvate -- cinchona alkaloid interaction and the steric constraints imposed by the adsorption on the platinum surface, leads to a reasonable explanation for the enantio-differentiation of this system. Although the prediction of the complex formed between the methyl pyruvate and the cinchona modifiers have been made for an ideal case (solvent effects and a quantum description of the interaction with the platinum surface atoms were not considered), this approach proved to be very helpful in the search of new modifiers. The search strategy, which included a systematic reduction of the cinchona alkaloid structure to the essential functional parts and validation of the steric constraints imposed to the interaction complex between modifier and methyl pyruvate by means of molecular modelling, indicated that simple chiral aminoalcohols should be promising substitutes for cinchona alkaloid modifiers. Using the Sharpless symmetric dihydroxylation as a key step, a series of enantiomerically pure 2-hydroxy-2-aryl-ethylamines

58 were synthesized and tested as modifiers in the hydrogenation of ethyl pyruvate over Pt/alumina [43,44]. The napthalene derivative (2-(1-pyrrolidinyl)-l-(1-naphthyl)ethanol) (PNE, Fig. 5) proved to be a remarkably efficient modifier inducing at low hydrogen pressures (1-10 bar) enantiomeric excesses as high as 75%, which is comparable to that obtained with cinchonidine under these conditions. In contrast to cinchonidine, which is more effective at higher pressures (70-100 bar), the enantioselectivity of PNE decreases at higher hydrogen pressures due to partial hydrogenation of the napthalene ring and concomitant loss of the ~t-bonding. The importance of the extended aromatic ring system for suitable anchoring is illustrated by the fact that corresponding benzene and pyridine derivatives are ineffective (ee < 5%). The structurally related diol (pyrrolidinyl ring in PNE substituted by OH group) does also not provide enantioselection, indicating the relevance of the N-H-O interaction. .o,

HO

X =CH, N ----ee<5 %

PNE (X=CH --* ee = 75 %, X = N --- ee = 67 %)

O

H~ H3C

OEt O

NEA; ee = 82 %

+2

R1 R2

-H20

a:b

OEt

> 97:3

a (R1 = H, R 2 = CH 3) b (R1 = CH3, R2 = H)

Figure 5. Novel efficient modifiers for the hydrogenation of ct-ketoesters over Pt/alumina An attractive alternative to these novel aminoalcohol type modifiers is the use of l-(1naphthyl)ethylamine (NEA, Fig. 5) and derivatives thereof as chiral modifiers [45-47]. Trace quantities of (R)-or (S)-l-(l-naphthyl)ethylamine induce up to 82% ee in the hydrogenation of ethyl pyruvate over Pt/alumina. Note that naphthylethylamine is only a precursor of the actual modifier, which is formed in situ by reductive alkylation of NEA with the reactant ethyl pyruvate. This transformation (Fig. 5), which proceeds via imine formation and subsequent reduction of the C=N bond, is highly diastereoselective (d.e. >95%). Reductive alkylation of NEA with different aldehydes or ketones provides easy access to a variety of related modifiers [47]. The enantioselection occurring with the modifiers derived from NEA could be rationalized with the same strategy of molecular modelling as demonstrated for the Ptcinchona system. Not so long ago, the general opinion was that high enantioselectivity can only be achieved with natural, structurally unique, complex modifiers as the cinchona alkaloids. Our results obtained with simple chiral aminoalcohols and amines demonstrate the contrary. With enantiomeric excesses exceeding 80%, commercially available naphthylethylamine is the most effective chiral modifier for low-pressure hydrogenation of ethyl pyruvate reported to

59 date. Thus we can be confident that further research will lead to other effective modifiers extending the scope of enantioselective heterogeneous hydrogenation. 3.2 Chemoselective modification Liquid-phase hydrogenation has become a very established method in the production of fine chemicals. In the past decades most of the classical methods using chemical reducing agents have been substituted by catalytic hydrogenation methods using either non-modified or modified transition metal catalysts [48]. In contrast, progress in the catalytic liquid-phase oxidation with molecular oxygen [49,50] has been comparatively slow since its early discovery by D6bereiner [51]. The main reason for this is the often difficult control of selectivity and catalyst deactivation [50]. The latter is reflected by the extremely high catalyst/ reactant ratios applied in these oxidations, which render them frequently uneconomic. A considerable effort seems to be necessary to make this principally cheap and environmentally friendly method competitive to other oxidation methods applied in organic synthesis. Recently we have applied the concept of modifying the metal surface with adsorbed auxiliaries in the platinum-catalyzed oxidation of L-sorbose (scheme 2) to 2-keto-L-gulonic acid (2-KLG), intermediates in vitamin C synthesis [52]. This direct oxidation could offer a catalytic alternative to the presently used route for the oxidation of the C-I hydroxyl group of Lsorbose to the corresponding carboxylic acid, which involves three steps to achieve good selectivity due to the necessary protection (and subsequent hydrolytic deprotection) of the other four reactive hydroxyl groups. O H II I

OHH I I

HOH2C'-'C --(3 ""-"C--"C-CH2OH I

!

OH H

L-sorbose

I

OH

02

r suppo.

O H II I

OH I

H I

HO2C-'C -'C --G --C-CH2OH I

I

OH H

I

OH

2-KLG

Scheme 2 The direct oxidation of L-sorbose in aqueous phase with molecular oxygen over supported Pt and Pd catalysts is rather inefficient due to low selectivity paired with significant catalyst deactivation [53]. Modifying platinum metal catalysts by blocking some of the active surface metal atoms by deposition of inactive foreign metals, leading to a reduction of the ensemble size, has been successfully applied in designing catalysts for alcohol oxidation [54,55]. However, this strategy was not successful for L-sorbose oxidation. Upon modification with Bi or Pb, the initial rate increases, but the overall performance of the catalysts is unsatisfactory due to oxidation and subsequent dissolution of the promotors (modifiers) under reaction conditions [56]. An alternative approach to increase the oxidation rate is the use of alkaline solutions, because bases enhance the reactivity of L-sorbose and weaken the adsorption strength of 2KLG. Unfortunately, the rate enhancement at higher pH is accompanied by a drop in selectivity due to the poor stability of 2-KLG in alkaline solutions. To circumvent this problem, we have modified the platinum catalysts by adsorbed tertiary amines and carried out the oxidation in neutral aqueous solution [57]. This allowed to enhance the rate without increasing the pH of the bulk liquid, which leads to detrimental product decomposition. Small quantities of amines (molar ratio of amine : sorbose = I: 1700, and amine : Pts = 0.1) are sufficient for modification. Using amines of pKa ~ I0 for modification, resulted in a considerable rate enhancement (up to a factor of 4.6) with only a moderate loss of selectivity to 2-KLG. The rate enhancement caused by the adsorbed amines is mainly determined by their basicity (pKa). In contrast, the selectivity of the oxidation was found to depend strongly on the structure of the amine.

60 As concerns the nature of base catalysis in the L-sorbose ---- 2-KLG transformation, it is most likely that the oxidation of the aldehyde intermediate is accelerated via a rapid hydration to the corresponding geminal glycol [58]. The formation of a carboxyl group from the glycol intermediate is an oxidative dehydrogenation step, which is usually much faster on platinum metals than the direct oxygen insertion to the carbonyl group of the aldehyde [50,51 ]. In order to improve the selectivity, we have systematically changed the structure of the amine modifier and rationalized the interaction complex between modifier and L-sorbose using molecular modelling. The aim was to find an amine, strongly adsorbing on platinum and forming an interaction complex with L-sorbose, in which the oxidation of the C-1 hydroxyl group is sterically favored. This is the case with hexamethylentetramine (HMTA), which upon adsorption on Pt still possesses three accessible nucleophilic nitrogen atoms, enabling interaction with an OH group of sorbose, as illustrated in Fig. 6. Electrochemical model studies [59] revealed that hexamethylentetramine is adsorbed on Pt and not oxidized under reaction conditions. The co-adsorption of L-sorbose and hexamethylentetramine on Pt results in a tilted position of the reactant, in which only C-1 is exposed to oxidative dehydrogenation. This geometric constraints result in a strong improvement of the selectivity to 2-KLG upon modification of a Pt/C catalyst (Fig. 7).

Figure 6. Side view of energetically most favorable complex formed between L-sorbose and hexamethylentetramine (HMTA) auxiliary. Complex is stabilized by N-H-O hydrogen bond interaction.

Figure 7. Effect of HMTA auxiliary. Selectivity to 2-KLG as a function of Lsorbose conversion over 5 wt% Pt/C. Unmodified catalyst (O); catalyst modified with HMTA (O).

4. CONLUSIONS Using the design of titania-silica mixed oxides for epoxidation of bulky olefins, the combined use of the solution-sol-gel method and supercritical drying has been shown to be a potent and versatile tool for the structural and chemical tailoring of mixed oxides. The large variety of controllable sol-gel parameters together with the use of appropriate drying methods provide exceptional control of chemical and structural properties of as-derived mixed oxides. In the epoxidation of bulky olefins, amorphous titania-silica aerogels possess a considerable advantage compared to presently known Ti-substituted zeolites, due to the larger range in which the pore size can be tuned. The excellent epoxidation activity of the aerogels demonstrates that crystallinity is not a requirement for highly active Si-O-Ti connectivities. A powerful concept for improving the chemo- and stereoselectivity of metal catalysts is their modification by co-adsorbed auxiliaries. This concept should be particularly useful for

61 liquid-phase reactions performed at relatively low temperature, as applied in fine chemical catalysis. The auxiliaries, used in very small quantities (auxiliary : reactant ratio < 1:1000), can either be simply added to the reaction solution or brought onto the metal surface in a special pretreatment step. They have to possess a suitable anchoring group and the functions necessary for the desired interaction(s) with the co-adsorbed reactant(s). An interesting conceptual advantage of this method, compared to the well-known immobilization of auxiliaries (promotors), is the possibility of self-organization of the interaction between the mobile auxiliary and the reactant(s). With immobilized auxiliaries the geometric requirements for favorable interaction are difficult to control. The concept with co-adsorbed auxiliaries has been successfully applied for designing catalysts for the enantioselective hydrogenation of a-ketoesters and the platinum-catalyzed oxidation of L-sorbose to 2-keto-Lgulonic acid, and may open routes to several other applications, where chemo- and stereocontrol are demanding. Essential for further progress is a better understanding of the interaction of co-adsorbed surface species, particularly in the liquid phase, where this concept seems to be most promising. REFERENCES

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J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

63

A R e t r o s p e c t i v e V i e w of A d v a n c e s in H e t e r o g e n e o u s C a t a l y s i s : 1956-1996, Science Robert L. Burwell, Jr. Catalysis Laboratory, Northwestern University, Evanston, IL 60201, USA Most of the large number of acronyms added to BET between ICC I in 1956 and ICC 11 in 1996 represent new techniques for analyzing the products of reaction, for characterizing catalysts and for determining the structure of chemisorbed species. Use of these techniques has substantially altered the practice of catalysis. Only the elderly will remember the great improvement in the quality of life which resulted from the fulfillment of Emmett's prophecy, that gas chromatography appeared "destined to provide a rapid, accurate method for the analysis of complicated mixtures of products in a relatively simple, straightforward manner" (ICC 1, paper 65). To GC can be added scanning FTIR, NMR, EPR, microwave spectroscopy, and various MS. Adequate description of many catalysts will require a large number of bits of data since they are usually rather complicated materials rather than simple chemicals. Attempts at this were just beginning by ICC 1, but now, one expects authors to give specific surface areas and some details of the porosity of their catalysts. Automation of the former tedious point by point measurement of the N 2 adsorption isotherm has greatly facilitated this. Use of chemisorption to measure specific site numbers has become much more common, for example, CO for Cr203, and H 2 to measure H/M (usually equated to Ms/M where M is an atom of a supported Group VIII metal). If the metal particles are roughly spherical, the H/M can be converted into an average particle size, but modern high resolution electron microscopy and wide angle x-ray scattering (WAXS) provide particle size distributions rather than mere averages. Further, in favorable cases, the first can provide atomic level resolution and lattice details and the second can diagnose the degree of strain in metallic particles larger than about 2 nm. However, most of these techniques require significant funding. For example, for best results, WAXS needs cyclotron radiation. The author thought himself very fortunate when his department got him a Type K Potentiometer and a Podbielniak distilling column in 1941-42. Among other new techniques for study of catalyst structure, EXAFS diagnoses the atomic environment about catalytic sites of noncrystalline materials and the x-ray absorption line just before the EXAFS wiggles has helped to identify the oxidation number of catalytic sites as have EPS and NMR in appropriate cases. Provision of information on catalyst structure probably represents the major advance in catalysis since ICC 1, but

64 even so, exact details of the surface sites are stilllacking and particles smaller than 1 run can be elusive. Chemisorption on nonmetallic catalysts should provide the n u m b e r of catalytic sites and for comparative purposes a single M s can be taken as a catalytic site on metals. This permits the calculation of turnover frequencies which was a n e w concept in post ICC 1 and which permitted intercomparison of catalyst activities. For the firsttime then, one has been able, for example, quantitatively to discuss support effects in Rh/support catalysts. Except for support effects, structure sensitivity has usually appeared in one of two aspects, variation of rate with surface crystal face or with particle size. In ICC 1 Gwathmey reported in one of the first experiments with single crystal faces that different faces machined from Ni single crystal spheres catalyzed the hydrogenation of ethylene at different rates (ICC 1 paper 5). Many similar results have followed, Farnsworth (ICC 1, paper 15) studied 1 cm 2 nickel and platinum sheets employing UHV techniques. The much augmented rate of the hydrogenation of ethylene restdting from argon ion bombardment was drastically reduced by annealing, but that of H 2 + D 2 ,~ 2HD was unchanged. This may have been the first specific report of structure insensitivity. Despite many subsequent papers general agreement as to the origin of this counter intuitive phenomenon has not yet developed. In studying the structure of chemisorbed species, infrared absorption spectroscopy has been outstanding partictdarly as facilitated by scanning FTIR. Eischen's paper (ICC 1, 67) represented what was probably the first in situ examination of chemisorbed species during a catalytic reaction (IR of CO + O 2 on Ni-NiO). The other vibrational spectroscopies, laser Raman and magic angle spinning NMR, have also been useful. Despite its low resolution, high resolution EELS has been useful in UHV work for assessment of surface cleanliness and for the identification of adsorbed species. In ICC 1 there were only a few references to diffusional limitations, but they may have been present in a number of papers. Despite improved attention, problems may still exist particularly in systems involving transport from the gas to the liquid phase. Absent a demonstration that the rate of a hydrogenation was proportional to the amount of c a t ~ y s t one may suspect that C(H2)(liq.) was not in equilibrium with P(H2)(gas). Acidic, high area silica-alumina had received substantial attention in ICC 1, (52-58). Perhaps the most dramatic change in the subsequent cat~ytic literature was the debut of zeolites. Why acid catalyzed reactions are so much faster on zeolites than on silica-alumina has been extensively discussed but probably not conclusively. One should be able to know the exact structures of catalytic sites in zeolites, but initial hopes that this would do wonders for mechanistic understanding have not been fully realized. Super acids and carbonium ions came into heterogeneous catalysis from homogeneous chemistry and in special cases reaction via carbonium ions seems to occur,

CnH2n+2 + [H +] -~ CnH2n+3+ --) CnH2n+l + + H 2 Carbonium Carbenium

65 The carbenium ion so formed then reacts in the ICC 1 manner except perhaps for not abstracting a hydride ion from another alkane. Although initial views that zeolites in general were super acids have come into question, definite super acids have been found such as calcined H2SO 4 oZr(OH) 4 which catalyze the isomerization of alkanes at low T. Noble metal/acidic zeolite catalysts constitute "dual functional catalysts" whose basic chemistry resembles that of M/SiO 2 ~ A1203. They have been extensively investigated with particular emphasis upon determination of the exact size and location of the metal cluster. Although fusion of isolated noble metal atom clusters of the size of those which could exist in zeolite pores into large cryst~s would be very exothermic, conditions for preparing some metal/zeolite catalysts with clusters cont~ning ca. 6 atoms located in the pores have been developed. How such clusters differ electronic~ly (the quantum size effect) and catalytically from dusters of 200 atoms is still under investigation. Zeolites have led to a new phenomenon in heterogeneous catalysis, shape selectivity. It has two aspects: (a) formation of an otherwise possible product is blocked because it cannot fit into the pores, and (b) formation of the product is blocked not by (a) but because the transition state in the bimolecular process leading to it cannot fit into the pores. For example, (a) is involved in zeolite catalyzed reactions which favor a para-disubstituted benzene over the ortho and meso. The low rate of deactivation observed in some reactions of hydrocarbons on some zeolites has been ascribed to (b) inhibition of bimolecular steps forming coke. Full application of structural findings requires a complete determination of mechanism which in most heterogeneous cat~ytic reactions involves a closed cycle of adsorption-desorption steps sandwiching a sequence of elementary steps among chemisorbed intermediates. The problem is more difficult than in homogeneous catalysis in part because two or more surface sites of somewhat different properties are likely to be involved. Although the zeitgeist of the period favored structural studies rather than mechanism, advances, albeit only partial, were made in the mechanism of particular reactions. The general state of knowledge about elementary steps is relatively primitive, the mechanism of the hydrogenation of ethylene still lacks a consensus and, in particular, the role of slowly reacting surface species in hydrocarbon reactions is still unclear. Two relatively new areas which have flourished during the interval since ICC 1 and have strongly influenced ideas as to possible structures of chemisorbed complexes and possible elementary steps are organometallic chemistry-homogeneous catalysis and surface science (in which there were two papers in ICC 1, papers 5 and 46). Surface science has provided structures of various chemisorbed species on the surfaces of single crystal planes of metals and some reactions of such species. Applicability of these data was restricted by their UHV origin. However, improvement resulted from the development of pressure cells in which catalytic reactions could be run at 1 atm or above and followed by examination of the surface in UHV. The great advantage here over conventional catalytic work was the provision of detailed information about the state of the surface. However, even so reaction could not be followed in situ as

66 is possible with infrared or EPR. Further, species found on the surface in UHV work come without labels identifying them as reaction intermediates. Errors in proposing particular species as reaction intermediates have inevitably occurred, but, of course this is equally the case for species found by IR or EPR and it does not detract from the importance work in this area. ICC 1, paper 61, surveyed reactions homogeneously catalyzed by metal carbonyls and presented one of the earliest discussions of the application of the mechanisms of such reactions to heterogeneous catalysis, for example, of the structure of CO as a ligand to that as a chemisorbed species. Further, a new class of heterogeneous catalysts has appeared, heterogenized homogeneous catalysts. The chemistry of the new catalysts was usually similar to that of their homogeneous predecessors, but some were unparalleled among organometaUic catalysts. Some actinide complexes on alumina which appear to consist of adsorbed L2ThH+ (L is pentamethylcyclopentadienyl), rapidly catalyze isotopic exchange between D 2 and H 2 at-195~ (only Pt having a comparable rate), and between D 2 and ethane at 90~ and also hydrogenations of simpler olefins at -45~ and of benzene at 90~ Since the O.N. of Th is effectively held at +4, oxidative addition and reductive elimination steps cannot be involved in hydrogenation and exchange and evidence is strong for the following intermediates,

D.~...D

..""D

"Th " ~ H g+ A

~-

.... ----

D:::

...-H -....

"'"'""H

"'"'TM'""'" O

H:'"" "...

..;"R

""Th"" C

A (either a transition state or a free energy minimum) reacts to form B. In hydrogenation, R in C is an alkyl group formed by reaction of H with adsorbed olefin and the mechanism is not Horiuti-Polanyi-like. Rather reaction via both A and C are among the few relatively clear examples of Rideal-Eley processes. Charging problems have limited the UHV work largely to metals and absence of organometallic complex analogs to metals has limited applications of work in this area largely to nonmetals. Two zeolites with chemistry not analogous to acidic SiO 2 ~ have received particular attention. The low alumina ZSM-5, free or metal-loaded, catalyzes CH3OH -~ gasoline (on HZSM-5), propane -~ aromatics (on ZSM-5 loaded with gallium), and other reactions too numerous to detail. A material of the same crystal structure but containing about 1% of structural Ti and devoid of A1 (TS-1) has unique oxidative capacities for those hydrocarbons which can enter its relatively narrow pores. At about 50~ and with aqueous H20 2 as the oxidant, it catalyses the epoxidation of alkenes and the hydroxylation of alkanes and arenes. Pt/L-zeolite in which all acidic [H +] including that resulting from reduction of Pt 2+ is replaced by K + catalyzes conversion of hexane to benzene with high selectivity. A key feature is the very low rate of

57 deactivation. Shape selective favoring of the initial transition state is not necessary since a non-microporous PtIMgO-A120 3 is equally effective. For economic reasons the oxidative dimerization of CH 4 at about 750~ was extensively investigated. Li+/MgO was the most studied catalyst, but a wide variety of oxides are effective, none lamentably with enough selectivity to be interesting commercially. The reaction which probably involves oxidation at the surface with liberation of CH 3 to the gas phase where it dimerizes is, like the Rideal-Eley step, an exception to the earlier statement about mechanism. Strong metal support reaction SMSI was a popular subject for study. The type catalyst, Pt/TiO 2 reduced at 300~ behaves much like Pt/SiO 2, but reduction at 500~ largely eliminates its capacity for the chemisorption of H 2 and hydrogenation while inducing activity for the hydrogenation of CO. Ti suboxide formed by reduction encapsulates the Pt particles. In recent years research of possible utility in the production of fine chemicals has increased substantially and in part consequent to government policy. This work has been too variegated to summarize briefly. A flurry of work in the hydrogenation of CO also originated in government policy. It led to the elaboration of our understanding of these reactions, but it is not clear that it led to major developments. There are myriads of possible oxidations for which AG is negative but which are unknown non-catalytically except in some cases, enzymatically. This situation has stimulated a great deal of research in which one success, TS-I was mentioned above. Oxidative dehydrogenation has received extensive attention. In particular, many compounds of Mo and V have been studied because of their practical utility. The Mars and van Krevelen mechanism (1954) has become dominant in interpreting oxidations on altervalent oxides, i.e. the proximate oxidant is a metal ion in a higher oxidation state. Oxygen serves to return the reduced metal ion to its original state. The reaction, CH 4 + 1/202 --~ CO + 2H 2, has been found to occur with surprisingly high selectivity on Rh or Pt monoliths at 1000~ and very high space velocities. Many papers have employed the concept of spillover, most commonly of hydrogen dissociatively adsorbed on supported X onto the support. In type cases like reduction of MoO 3 by H 2 promoted by Pt on the MoO 3 or D 2 exchange with surface OH in Pt/A1203, the occurrence of spillover is clear but perhaps not the details of migration of H over the support. However, not all papers clearly establish the catalytic relevance of spillover even where it occurs. Further several papers have reported that the rate of a hydrogenation reaction on Pt/A120 3 or the like is increased by mixing some support into a batch of catalyst. Whence, H spills over onto the support to effect hydrogenation and the authors all quote an earlier paper with such a result. But the referees failed to point out to the authors that subsequently J. Catal. 24, 482 (1972) demonstrated that the augmented rate resulted from scavenging of poisons by the added support. It is difficult now-a-days to be conversant with all previous catalytic literature, but the level of ignorance appears to have increased substantially (and mea culpa). Too many papers report no assay of poisons in the feeds.

68 Stimulated by the growing importance of enantiomeric purity in drugs and insecticides, enantioselective hydrogenation has received augmented attention, primarily with homogeneous catalysts. However, the optically active alkaloid, cinchonidine, adsorbed on Pt/SiO 2 catalyses the liquid phase hydrogenation of pyruvate, C H 3 C O C O O M e , to methyl lactate, CH3CHOHCOOMe, of substantial optical activity. Presumably, adsorption of cinchonidine on Pt generates an optically active catalyst, but why should the rate be 30X that of the same catalyst without cinchonidine? Consequent to the work of m a n y and employing such techniques as structure variation, isotopic tracers, and stereochemistry, a large number of different adsorbed hydrocarbon fragments have been identified as key intermediates in various reactions of hydrocarbons. Correlation of these species with similar polynuclear organometallic species has been of interest. However, the author feels that mechanistic understanding has lagged behind some other aspects of catalysis. Finally, theory involving collective electrons and the like, characteristic of ICC 1, appears less frequently in the catalytic literature and theories involving local sites appear more oRen. The existence of coordination complex analogs m a y have played some part in this development. More recent theoretical developments relate primarily to chemisorbed species. Catalytic polymerization is surveyed in Heinemann's paper on advances in catalytic technology. Drastic sampling and subjective judgment was required to condense 40 years of research in heterogeneous catalysis into a few thousand words. There might be rather littleoverlap between this paper and that with the same title written by another author. References have been omitted to avoid presenting a sampling of names with the implication that they alone were responsible for the advances.

J.W. Hightower,W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

69

A R e t r o s p e c t i v e V i e w of A d v a n c e s in H e t e r o g e n e o u s Catalysis: 1956-1996 T e c h n o l o g y Heinz Heinemann Materials Sciences Division, Lawrence Berkeley National Laboratory, University of California, Berkeley, CA 94720*

The first International Congress on Catalysis in 1956 occurred a t a time at which catalytic technology and industrial applications of catalysis were rapidly increasing at an almost exponential rate which accelerated through the next 20 years. One would have been hard put to envision the explosion of applied catalytic work, which, to some extent, was stimulated by the worldwide demand for goods following the recovery from the Second World War. It is difficult to do justice to the many developments in the few pages allotted to the author and this article will be limited to what the author considers the most outstanding new catalysts and catalytic processes. One objective of this paper is to show that, contrary to some statements that the rate of catalytic developments has slowed in recent years, there has been a rapid growth of technology, though the areas to which it is being applied have shifted. A major factor in the growth of new technologies, one that is not specific for catalytic developments, but applies equally to all technologies, as well as to theory, is the d e v e l o p m e n t and wide availability of more and more sophisticated computers. In the last few decades, computers have changed means of observation, reduced calculation times and simplified operations, both on the bench and on the industrial scale. The proceedings of the First ICC are divided into four major subjects and the same subjects could be applied to the current meeting. They were: "Chemistry and Physics of Solid Catalysts", "Homogeneous Catalysis and Related Effects", "Surface Chemistry and Its Relation to Catalysis" and "Techniques and Technology of Catalysis". About 20 of the 80 papers presented at the meeting were in the last category, a percentage that has not increased substantially at subsequent Congresses. 12 papers were concerned with homogeneous catalysis, none of which fitted into the category of technology. Homogeneous catalysis in technical applications has grown even faster than heterogeneous catalysis, though starting from a much lower base. While there are many homogeneously catalysed processes in use, each of them is limited to a relatively small volume of products. Present address: LBNLWashington, D.C. Project Office, 1250 Maryland Ave., S.W., Suite 500, Washington, D.C. 20024.

70 In view ot~ the limited space allowed, this paper will be essentially restricted to heterogeneous catalysis. The major initial driving force in the expansion of catalytic processing was the worldwide demand for energy and the availability of relatively cheap petroleum. This led to the development of major new processes in petroleum refining and in the petrochemical industry, as well as to inventions which revolutionized existing technology (Table 1). Table 1 Major new catalytic technology developments between 1956 and 1996 ,

Catalysts

...

Year

Process

1957

Polymerization

Ziegler-Natta catalysts

1962

Steam reforming

NiK2AI203

1964

Catalytic cracking with faujasite zeolites

x and y zeolites

1967

Bimetallic reforming

Pt-Re; Pt-Ir

1968

Selectoforming (shape selectivity)

Erionite

1972

Low pressure CH3OH

Cu-Zn-Al20 3

1974

Acetic acid v/a carbonylation

RhI

1976

Auto emission control for HC and CO

Pt-AI203

1980

Gasoline from methane

ZSM-5 zeolite

1982

NOx control

Pt-Rh for autos V205-TiO2 stack gas

1988

Selective oxidation

TiSiO2

1988

Chiral catalysis

Zeolites, or SO2 cinchonidine on supported Pt

1991

Polymerization

MetaUocenes

In the 1960s and subsequent years, catalytic processing was extended to new areas, particularly to the abatement of environmental concerns (Table 1). Automobile emission control catalysis became a major factor of catalytic processing and catalysts for this purpose today constitute a large portion of catalyst manufacture. Other environmental concerns resulted in stack gas catalytic conversion, primarily to remove NOx and SO2. It was also necessary to improve the selectivity of established processes to minimize toxic emissions and to increase energy efficiency. During the 1980s, a trend toward fine chemicals production and particularly toward pharmaceuticals became noticeable, which resulted in numerous new processes and reactions involving catalysis. There has thus been a

71 change from the very large volume of relatively few low-priced products to a large number of small volume, high-priced products. The importance of industrial catalysis is evidenced by the fact that over 18% of the U.S. GNP is achieved by catalytic processing. Catalyst manufacture has large markets of its own, as shown in Table 2. Table 2 Catalyst markets ($ billion)

Western Europe U.S.A. Worldwide

1989

2000 (projected)

1.3 1.9 5.0

1.9 2.4 6.5

Among the many new process inventions of the last 40 years, three stand out as having ushered in several new technologies. The first is the discovery of the catalytic properties of zeolites, the second is the use of precious metals in emission control and the third is selective polymerization. While catalytic cracking was a subject of major interest at the time of the first Congress in 1956, it was then generally assumed that only amorphous solids, such as clays or SIO2-A1203, possessed the necessary properties for acid-catalyzed reactions and that crystalline materials were of little, if any, interest. This myth was debunked by a paper which was presented at the second ICC in Paris in 1960. In the early 1960s, workers at Mobil demonstrated not only the greater activity but also the greater selectivity of cracking catalysts containing faujasite-type zeolites. Zeolite-containing catalysts dominate catalytic cracking today and their use has resulted in savings of petroleum of ~ 400 million barrels/year for the same amount of gasoline produced. At $17/bbl, this amounts to ~ $7 billion/year. In the mid-1960s, the shape selective properties of small and medium pore size zeolites were discovered at Mobil. The first commercial application was the "Selectoforming" process which selectively cracked straight-chain, low-octane number paraffins in gasoline. This was followed by new or improved petrochemical processes, such as the easier isomerization of xylenes at higher para selectivity, the disproportionation of toluene to benzene and xylene, and many others. Introduction of metals into the pores of zeolites at specific crystal sites made other processes possible. Today there are dozens of applications of zeolite catalyzed processing. While the discovery of the catalytic properties of zeolites was driven by the desire to improve industrial processing, the development of emission control catalysts was necessitated by governmental fiat. The first requirement was for 90+% removal of CO and of hydrocarbons, a goal which could not be met by oxidation with base metal oxides. To achieve the required specifications during automobile operations, it was necessary to develop supported platinum catalysts. Originally the support was alumina in pellet form. Later platinum on cordierite was used in honeycomb form, containing 200-400 square channels per square inch.

72 The original objective was the control of hydrocarbon and CO emissions. The danger of NOx emissions and the need to limit them became increasingly important in the 1980s. This was not possible with Pt on cordierite catalysts (usually containing rare earths, e.g., lanthana) but required the addition of rhodium (in quantities of approximately half of that of Pt) and a very careful control of the air/fuel ratio at 14.7 + 0.1 volumes. As shown in Fig. 1, there is a very narrow window in which both CO and NO removal are above 90%. An exhaust oxygen sensor and electronically controlled oxygen feed maintain this ratio in modern automobiles. The development of automobile exhaust catalysts resulted in these catalysts becoming the largest volume of catalysts produced, replacing catalytic cracking catalysts. Numerous other examples of environmental control by catalysis will be mentioned later. It must also be stated that in the attempt to minimize toxic by-products, not only their destruction but also avoidance of their production, has become important. Catalysis has and will continue to play a major role in improving the selectivity of chemical processes, thus eliminating or reducing undesirable by-products. Another field in which catalysis has played a big role in consumer products is polymerization. Ziegler-Natta polymerization of ethylene or propylene was invented just before the 1956 Congress, but became large-scale commercial in the late 1950s. The aluminum alkyl-titanium chloride catalysts originally employed in the high-pressure synthesis of polyethylene, polypropylene and other plastics were modified a n d / o r replaced by heterogeneous catalysts, such as chromia on silica, and a new low-pressure synthesis was developed. Books can and have been written on catalytic polymerization of various monomers (e.g., styrene, in addition to C2H4 and C3H6), on stereospecific polymers and control of the chain length of polymers. 100

o~

=s O

rt... r O

O

0 14.0

.j

14.5

'i' ...... 15.0

15.5

Simulated A/F Ratio

Figure 1. Principle of 3-way exhaust emission control. Catalyst efficiencies measured in the laboratory with a steady feed stream composition at various simulated air/fuel ratios. Catalyst: 0.042 wt % Pt/0.018 wt % Rh/alumina.

73 In 1991, metallocene-catalyzed polymerization became commercial technology, particularly for polyethylene and polypropylene. These two polyolefins account for 35% of all thermoplastics and elastomers. MetaUocene can also readily polymerize bulky monomers, such as styrene, to make novel polymers with physical properties competitive with nylon, polycarbonates and polyesters. Metallocenes are single site catalysts comprising a metal atom sandwiched between parallel planar cyclopentadenyl groups. The most common metals used for olefin polymerization are zirconium, titanium and hafmium. These systems have extremely high activity (as high as 40,000 kg polyethylene/g, metal/hour), which more than compensates for the relatively high catalyst cost and makes the system competitive with Ziegler-Natta catalysts. There are several features that distinguish metaUocene catalysts from other systems. They can polymerize vinyl monomers regardless of the monomers molecular weight or stress hindrance; they produce very uniform polymers of narrow molecular weight distribution; they can polymerize r with very high stereoregularity to give isotactic or syndistactic polymers. The production of some oxygen-containing monomers (e.g., various alcohols or aldehydes) later to be polymerized to fibers and plastics has been simplified and new chemistry has been opened up by the discovery in the early 1990s of the amazing properties of titanium silicates in crystalline structures as oxidation catalysts, which have been named TS-1, etc. Highly selective oxidations can be carried out using hydrogen peroxide (in the presence of oxygen) as oxidant to produce epoxides, among others (Fig. 2). These oxidations can be accomplished Reactants

~

Products

+ H202

-OH

OH

OH

+ H20 OH

+ H202

R~C= C

* H202

+ H20 O R~C/_.-~ C

OH

+ H20

O

C---~C~c._.C

+ H202

/\ C~C~c._. C

R-CH2-OH

+ H202

R-CHO

R~

+ H202

R/CH-OH

R~

R / C -- 0

+ H20

+ H20

+ H20

Figure 2. H202 oxidations catalyzed by TS-1.

74 with dilute aqueous H202 (e.g., 40 wt % H202) with no loss in selectivity which, in almost all cases, is higher than 80%. The synthesis of propylene oxide from propylene with aqueous H202 has a selectivity of 98%. This area is expected to grow rapidly in the next few years. While the titanium silicates have structural similarity to certain zeolites, they are not acidic zeolites and catalyze entirely different reactions. Another area of oxidation that has shown major growth is dehydrogenation and oxydehydrogenation, particularly in the conversion of butane and butenes for MTBE production. Several processes have been developed using metal oxide or precious metal catalysts but, in general, they are improvements of the pre-1956 Houdry butane dehydrogenation and the Texas Butadiene process. Reactor design and process engineering have greatly contributed to the improvements. The oxidation of butanes to compounds, such as maleic anhydride, in the presence of vanadium or molybdenum oxides has gained importance since 1956, again mostly due to design improvements. A new catalytic route to acetaldehyde from ethylene was introduced as the Wacker process. In a partially homogeneous reaction, the ethylene is oxidized with PdCI2 in the presence of water. Reduced Pd metal is oxidized back to PdCI2 with cupric chloride and the resulting cuprous chloride is oxidized with HCI and oxygen. Several other important commercial processes need to be mentioned. They are (not necessarily in the order of importance): the low pressure methanol process, using a copper-containing catalyst which was introduced in 1972; the production of acetic acid from methanol over RhI catalysts, which has cornered the market; the methanol-to-gasoline processes (MTG) over ZSM-5 zeolite, which opened a new route to gasoline from syngas; and ammoxidation of propene over mixed-oxide catalysts. In 1962, catalytic steam reforming for the production of synthesis gas a n d / o r hydrogen over nickel potassium alumina catalysts was commercialized. At times, it is difficult to distinguish between revolutionary new breakthroughs and evolutionary improvements of technology resulting in a replacement of previously used catalysts or processes. An example of this is the introduction in 1967 by Chevron of bimetallic reforming catalysts. Reforming with Pt-alumina catalysts was an important development at the time of the first ICC and many variants of catalyst composition and process operating conditions were introduced in the 1950s and early 1960s. The new bimetallic catalysts, mostly containing rhenium, in addition to platinum and alumina, greatly increased the stability and thus life of the catalysts and also permitted lower pressure operation, increasing the yield of reformate. An additional step was taken by Exxon with the introduction of Pt-Ir catalysts. In general, bimetallic reforming catalysts have largely replaced the Pt-AI203 catalysts. The new reforming catalysts were much more sulfur sensitive than previous ones and that led to hydrodesulfurization catalysts which would lower the sulfur level of reforming feeds to the part per billion level. Hydrocracking, though well-known in 1956, was undergoing major improvements in the 1960s and 1970s by using zeolite (faujasite)-based acidic components, along with tungsten or molybdenum oxides or sulfides as

75 hydrogenation components. Dependence on heavier petroleum has formed greater dependence on hydrocracking processes. Hydrodesulfurization has become vastly more important in the last 40 years. The great sensitivity of metal catalysts to sulfur has necessitated reducing sulfur levels in the feed to reforming processes to the parts per billion level, and the presence of aromatic sulfur and nitrogen compounds in heavy oils requires severe hydrotreating. The catalysts used are upgraded versions of oxides or sulfides of molybdenum, tungsten and cobalt. In the area of pollution control, the removal of NOx from stationary sources effluents, such as power plant stack gases, has been accomplished by use of titaniavanadia catalysts, which promote the reduction of NOx with NH 3 to produce nitrogen and water. During the last 20 years, emphasis on catalytic steps in the synthesis of pharmaceuticals and of agrochemicals has sharply increased. Major progress in novel catalytic reactions has, to some extent, shifted from hydrocarbon conversions and fuels to environmental control and to organic (and, to a lesser extent, inorganic) chemical synthesis. Several fine chemicals are now produced by catalytic reactions (often using zeolites) with greater selectivity and therefore better meet environmentally acceptable standards. Important among biochemicals is "chirality", the "handedness" or optical rotation of the molecule, which determines its bioactivity. One of the chiral isomers can be very effective, while the other may be ineffective or even harmful. A typical example is the case of thalidomide, which is a sedative for pregnant women in its R-form, but is a potent teratogen in its Sform, causing children to be born with deformities (Fig. 3). The U.S. market for chiral compounds is expected to grow from $0.5 billion in 1990 to $3 billion in 2000. Any short review is bound to be incomplete and the selection of successful processes and catalysts expresses only the opinion of the author. There is, however, no doubt about the magnitude of important developments of the last 40 years, with at least ten breakthrough accomplishments, and there is every indication that the trend of the last 40 years will continue. H

H 0 . ~ 0

0

0

0

(R) Sedative

0

(S) Teratogen

Figure 3. Physiological effects of chirality: thalidomide.

This Page Intentionally Left Blank

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

77

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

An example of novel o x y n i t r i d e s or "A1PONs"

basic

catalysts:

the

aluminophosphate

A. Massinon a, E. Gu~guen a, R. Conanec b, R. Marchand b, Y. Laurent b and P. Grange a a Unit~ de Catalyse et Chimie des Mat~riaux Divis~s, Universit~ Catholique de Louvain, Place Croix du Sud 2/17, 1348 Louvain-la-Neuve, Belgium* b L a b o r a t o i r e de Chimie des Mat~riaux, URA 1496 CNRS "Verres C~ramiques', Universit~ de Rennes I, 35042 Rennes Cedex, France

et

New a l u m i n o p h o s p h a t e oxynitrides solid basic catalysts have been synthesised by activation under ammonia of an A1PO4 precursor. When the nitrogen content increases, XPS points out two types of nitrogen phosphorus bonding. The conversions in Knoevenagel condensation are related to the surface nitrogen content. Platinum supported on aluminophosphate oxynitride is an active catalyst for isobutane dehydrogenation. 1. INTRODUCTION The development of solid basic catalysts is nowadays a subject of increasing interest. Modified oxides [1], zeolites [1], hydrotalcites [2] and alkaline-substituted sepiolites [3] have shown interesting activities as solid basic catalysts. Phosphates, in particular aluminophosphates, are known to be good acid catalysts. A new family of aluminophosphate oxynitrides, called "A1PONs", has been prepared by nitridation of reactive aluminophosphate powders [4]. The presence of nitrogen in the three-dimensional network of the oxynitrides results in the creation of basic sites on their surface [5]. This paper reports surface characterisation of a series of aluminophosphate oxynitrides with variable nitrogen contents and their catalytic evaluations in the Knoevenagel condensation and isobutane dehydrogenation. 2. EXPER/MlgNTAL 2.1. Synthesis of oxide, oxynitrides a n d impregnation of Pt and Sn To prepare a high surface area amorphous phosphate precursor A1PO4, the citrate method was used [6]. To reach an Al/P ratio fixed at 1, 0.667 mole of Al(NO3)3.9H20 (Merck) and 0.667 mole of (NH4)H2PO4 (Merck) were dissolved * We acknowledge the financial support of the "R6gion Wallonne", Belgium, for this COST pro~am.

78 in distilled w a t e r u n d e r stirring. Both solutions were mixed t o g e t h e r at room t e m p e r a t u r e . After 1 hour stirring, an excess of citric acid (Merck) was added. The resulting aqueous solution was f u r t h e r stirred overnight. W a t e r was t h e n evaporated under reduced pressure and the obtained white gel was dried for 10 hours at 378 K in a vacuum oven (50 mbar). The gel swelled and looked like a "meringue". The sample was then calcined for 16 hours at 823 K. N i t r i d a t i o n of the oxide precursor was performed u n d e r p u r e a m m o n i a flow. Different aluminophosphate oxynitrides "A1PONs" with variable nitrogen contents were obtained by modifying the time and/or the t e m p e r a t u r e of nitridation (Table 1). A 1.5 wt.% Pt/A1PON catalyst was prepared by i m p r e g n a t i o n of p l a t i n u m on the A1PON support. The Pt/A1PON s a m p l e was o b t a i n e d by i n c i p i e n t wetness i m p r e g n a t i o n with a methanolic solution of H 2 P t C 1 6 . 6 H 2 0 (Merck). The excess of methanol was evaporated u n d e r an argon flow a n d the s a m p l e was then dried at 383 K overnight. The sample was decomposed u n d e r N2 flow w i t h a t e m p e r a t u r e r a m p of 2.5 K.min "1 up to 773 K and held at t h a t t e m p e r a t u r e for two hours, after which the s a m p l e was r e d u c e d in p u r e hydrogen at the same t e m p e r a t u r e during two hours. The addition of tin was accomplished by impregnation with a methanolic solution of Sn(CH3)4 (Merck) on the Pt/A1PON sample, using the same drying and reduction procedures as for Pt/A1PON. 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 f i o n The specific surface areas of the samples were m e a s u r e d by the single point BET method (p/p0=0.3). The total a m o u n t of n i t r o g e n (bulk n i t r o g e n of n i t r i d e - t y p e a n d h y d r o g e n a t e d N H x (x=l to 4) surface species) was d e t e r m i n e d by Grekov titration [7]. The principle of this chemical analysis of nitrogen is based on the reaction of the nitride ions N 3" with a strong base and the f o r m a t i o n of a m m o n i a t h a t is then titrated. In the traditional Kjeldahl method the alkaline a t t a c k occurs in solution. But in some cases, with r e f r a c t o r y nitrides, the products are not totally attacked under these conditions. This has been solved by using the same principle but by h e a t i n g the product at high t e m p e r a t u r e (673 K) with melted potassium hydroxide. The ammonia was dissolved in w a t e r and titrated with sulphuric acid. In the aim to estimate the a m o u n t of surface N H x species (p)S P - H, P- N(H, NH3, PO" . .NH4 + ) adsorbed at the surface the classical method of Kjeldahl was used. X-ray Photoelectron Spectroscopy analysis of the samples was performed with a Surface Science I n s t r u m e n t s spectrometer (SSI 100) with a resolution (FWHM Au 4f7/2) of 1.0 eV. The X-ray beam was a m o n o c h r o m a t i s e d A1Ka r a d i a t i o n (1486.6 eV). A charge n e u t r a l i s e r (flood gun) was a d j u s t e d at an energy of 6 eV. As the C l s spectra of these compounds were very complex, the binding energies were referenced to the binding energy of O l s , considered experimentally to be at 531.8 eV [8].

79 2~3. Catalytic evaluations

2~.1. Knoevenagel condensation F o u r mmoles of malononitrile and b e n z a l d e h y d e were introduced in a batch stirred t a n k reactor at 323 K with toluene as solvent (30 ml). Then 0.05 g of a l u m i n o p h o s p h a t e oxynitride was added. Samples were analysed by gas c h r o m a t o g r a p h y ( I n t e r s m a t Delsi DI200) using a capillary column (CPSi18CB25 m). Care was t a k e n to avoid mass or h e a t transfer limitations. Before the reaction no specific catalyst p r e t r e a t m e n t was done. 2~.2. Isobutane dehydrogenation Catalytic evaluations of isobutane dehydrogenation were carried out in a c o n v e n t i o n a l c o n t i n u o u s flow m i c r o r e a c t o r o p e r a t i n g at a p p r o x i m a t e l y a t m o s p h e r i c p r e s s u r e and using i s o b u t a n e in helium (Air Liquide 0.95% isobutane N25 in 99.05% helium N50) and hydrogen in helium (Air Liquide 1% H2 N50 in 99% He N50). The molar ratio between isobutane and hydrogen was m a i n t a i n e d in each case at 1/6. The reaction products were analysed using an on-line P a c k a r d (model 428) gas c h r o m a t o g r a p h equipped w i t h a flame ionisation detector, with helium as c a r r i e r gas. A 60m x 0.32ram (RSL 160 Alltech) column was used for the s e p a r a t i o n of the various compounds. The space velocity is given as the ratio between the weight flow of isobutane per hour and the weight of the catalyst. Total isobutane conversion is defined as the p e r c e n t a g e of i s o b u t a n e t r a n s f o r m e d into all products. The selectivity to isobutene is defined as the amount of isobutane converted into isobutene divided by the total isobutane conversion. Initial conversion is obtained by extrapolating the curve conversion versus time-on-stream to time zero. 3. R E S U L T S AND DISCUSSION

3.1. Characteristics of samples The characteristics of the studied oxynitrides are reported in Table 1. Table 1 Composition, nitridation t e m p e r a t u r e and time, surface area and nitrogen content of the "A1PONs" (~omposition 'Nitridation Nitridation Surface area Total N Surface N t e m p e r a t u r e (K) time (h) (m2.g-1) (wt.%) (wt.%) A1PO3.64N0.24 1073 3 275 2.8 2.7 A1PO3.55N0.30 1073 8 275 3.6 1.3 A1PO3.10N0.60 1073 40 235 7.2 1.0 A1PO2.67N0.89 1073 65 230 11 1.1 A1PO1.96N1.35 1073 120 215 17.5 2.4 A1PO1.71N1.53 1073 200 195 20 2.6

3.2. C h a n g e of s t r u c t u r e d u e to t h e n i t r i d a t i o n With XPS a good fitting of the N l s spectra is found with one peak of 85% Gaussian and 15% Lorentzian character.

80 It allows to observe a shift on the binding energies values. Indeed a break is evidenced between the behaviour of the samples with nitrogen content lower than 7% and the samples with higher nitrogen content (Figure 1). In the literature, Marchand et al [9] have shown that two N ls peaks exist in the phosphate glasses: a peak at 397.8 eV corresponding to P=N-P bonds and P

a peak at 399.3 eV corresponding to p>N- P bonds. The same trend is observed in our catalysts that suggests the presence of these two kinds of nitrogen coordination in the "A1PON". On the other hand the N l s binding e n e r g i e s of the a l u m i n o p h o s p h a t e O 398.8 oxynitrides "A1PONs" agree with the Z 398.6 N l s binding energies of bulk phosphate o oxynitride ceramic m a t e r i a l PON [10], >., 398.4 ea~ situated at 317.8 and 399.3 eV, r a t h e r 398.2 t h a n the N l s b i n d i n g e n e r g i e s of "A1ON" family [11,12], found between 398.0 396.3 and 396.6 eV. M o r e o v e r a = 397.8 progressive decreasing t r e n d of the I I I binding energies of P2p (from 134.1 to 397.6 0 5 10 15 20 133.5 eV) is observed when the atomic n i t r o g e n percent age increases. This Total nitrogen content (% bulk) behaviour can be explained by the better Figure 1. Variation of N l s binding nucleophilic c h a r a c t e r of n i t r o g e n c o m p a r e d to oxygen, r e d u c i n g the energy. positive charge a r o u n d p h o s p h o r u s atoms. Both observations indicate that, during nitridation, nitrogen seems to substitute more easily the oxygen atoms present in an environment of phosphorus r at her than in the environment of aluminium. >

399.0

cD

.,..

3dl. K n o e v e n a g e l condensation To evaluate properties of basic catalysts, the Knoevenagel condensation over aluminophosphate oxynitrides was investigated [13]. In this reaction usually catalysed by amines, the solid catalysts function by abstraction of a proton from an acid methylene group, which is followed by nucleophilic attack on the carbonyl by the r e s ul t ant carbanion, re-protonation of oxygen and e l i m i n a t i o n of water. The c o n d e n s a t i o n bet w een b e n z a l d e h y d e and malononitrile is presented below. To check if this liquid-phase reaction is not controlled by diffusion, the reaction is repeated with different weights of catalysts. A linear correlation is found between the initial rate of reaction and the weight of catalyst, indicating that the rate is not controlled by external or internal diffusion.

81

C +

\

NC

N

H

Base i

"

2

i

1

X

N

CN

H

/N +

H20

The i n t r i n s i c conversion (%.m "2) of m a l o n o n i t r i l e and b e n z a l d e h y d e v e r s u s time is shown in Figure 2. Commercial MgO (40 m2.g "1) used in the s a m e conditions as the "A1PONs" (i.e. w i t h o u t p r e t r e a t m e n t ) gives a low conversion [14]. W i t h o u t p r e t r e a t m e n t MgO is not an e x t r e m e l y basic compound but these results show t h a t "A1PONs" are more active t h a n MgO at the chosen conditions and such a c h a r a c t e r could be useful for i n d u s t r i a l applications.

10 O 9~

8

~

6

A

A

8

A

d

6 ~

0

~4 9~

,10

&

2:

"A

m

2 ~

2 D

0

0

100

200

O |

300

Time (rain) Figure 2. Conversion of malononitrile vs time in Knoevenagel condensation (O 2.8%N, 9 3.5%N, ~ 7.2%N, m l l % N , A 17.5%N, A 20%N).

10

0 ~ 20

Nitrogen content (%) Figure 3. Intrinsic conversion at 300 rain and surface N content vs total nitrogen content.

Figure 2 shows t h a t the conversion does not depend directly on the total n i t r o g e n content of the c a t a l y s t as it was seen before for the "A1PONs" synthesised by the sol-gel method [5]. W h e n the N H x surface groups (Table 1) are only considered, a good correlation with the malononitrile condensation is obtained (Figure 3). A c t u a l l y N H x surface species are not the only active species in the Knoevenagel reaction. Indeed the catalysts with 2.8, 17.5 and 20 wt.% bulk nitrogen have the same NHx surface species content but their activities are

82 very different. The 2.8 wt.% bulk nitrogen sample can be a s s u m e d w i t h o u t nitride bulk nitrogen because its total nitrogen content and its surface nitrogen content (2.7 wt.%) are equivalent. On the other h a n d the 17.5 and 20 wt.% bulk nitrogen content samples contain the highest quantity of nitride bulk nitrogen. The higher the nitride nitrogen content, the higher the activity. The s t r e n g t h of the basic sites is not yet evaluated. Corma et al. [3,15] d e m o n s t r a t e d t h a t using methylenic compounds with different acidities, a basicity range can be determined for their samples. However the condensation is carried out without solvent. Using toluene as solvent the reactions between b e n z a l d e h y d e a n d m a l o n o n i t r i l e ( C N - C H 2 - C N - pKa = 11.2 [ 1 6 ] ) o r e t h y l c y a n o a c e t a t e ( C 2 H 5 C O 2 C H 2 C N - pKa < 9 [16]) are studied on the "A1PONs". N e v e r t h e l e s s the conclusions are not so evident in toluene. In toluene the condensation of malononitrile is easier t h a n the condensation of ethylcyanoacetate whereas the opposite behaviour was expected. Indeed if the first step (abstraction of the acid proton) is limiting, the abstraction of acidic p r o t o n of e t h y l c y a n o a c e t a t e should be e a s i e r t h a n the a b s t r a c t i o n of malononitrile one due to the lower pKa. To check if this behaviour does not depend on our catalysts, both condensations are done using liquid bases: pyrrolidine (pKa = 11.3 [17]) and n-nonylamine (pKa = 10.6 [17]). The results found with both liquid bases present the same trend but experimental data do not lead to conspicuous results because a steric h i n d r a n c e with the ester function or the better charge stabilisation of ethylcyanoacetate can occur and the second step could become limiting ( c o n d e n s a t i o n b e t w e e n the two reactants). So both facts lead to a lower conversion with e t h y l c y a n o a c e t a t e because the conversion does not depend on the pKa of the reactifs. On the other h a n d , s o l v e n t effects could play a role d u r i n g the r e a c t i o n . F u r t h e r experiments are needed to solve this doubt. 3.4~ I s o b u t a n e d e h y ~ e n a t i o n The oxynitride "A1PON" used as support was synthesised as the others. It is well represented by the A1PO3.16N0.56 formula, its nitrogen content being 6.7 wt.%. This sample has a surface area of 310 m2.g -1. 3.4.1 Influence of t h e noble metals s u p p o r t e d o n a l u m i n o p h o s p h a t e oxynitride T h r e e catalysts with different compositions were used in this study. Catalyst A corresponds to the oxynitride precursor. Catalysts B and C contain 1.5 wt.% of platinum, however catalyst C contains 0.91 wt.% tin as well corresponding to an atomic ratio Pt/Sn of 1. For q u a n t i t a t i v e comparisons between catalysts, care was t a k e n to ensure t h a t the kinetic d a t a were not influenced by mass or heat transfer. Figure 4 shows the evolution of the initial conversion versus t e m p e r a t u r e at a space velocity of 0.03 h -1. The equilibrium conversion of i s o b u t a n e to i s o b u t e n e is 100% in our conditions. An increase of the conversion with t e m p e r a t u r e up to 773-823 K is observed. When metals were added, we also noted a large increase in isobutane dehydrogenation. Table 2 gives initial isobutane conversions, isobutene selectivities and yields of the reaction at 823 K for the three tested samples.

83

100 A

80

9

o

.0,,q

~

60

= o

40 20

I00 ~ ' - ~

&

me

80 ~

9

dD

60 40

T

A

I=

0.~ - $ ~ ~--o--o--o--c 373 473 573 673 773

o

20"

o

O

0

873

40

60 Time (h)

Temperature (K) Figure 4. Isobutane initial conversion vs temperature (C) A1PON, I Pt/A1PON, A Pt-Sn/A1PON

20

Figure 5. Isobutane conversion and isobutene selectivity vs time for Pt-Sn/A1PON (O conversion, 9 selectivity).

Table 2 Initial isobutane conversions, selectivities and ~elds of the reaction at T=823 K catalyst A B C Isobutane conversion (%) 2.7 70.5 84.5 Isobutene selectivity (%) 50.5 44.8 38.4 Yield of the reaction (%) 13.6 31.6 32.4 At first, we show that adding metals to the oxynitride affects significantly the conversion while keeping isobutene selectivity at a high level. The byproducts of the reaction are due to isomerisation, hydrogenolysis and coking reactions. The addition of tin to the Pt/A1PON catalyst increases the conversion. Platinum-tin alloy formation, as already suggested in literature [18], or improvement of the metal dispersion cannot yet be proposed. Figure 5 shows the evolution of conversion and selectivity versus time for catalyst C, which presents the best yield. This catalyst shows a significant deactivation probably due to the coke deposit. 3.4.2 I n f l u e n c e o f t h e s p a c e v e l o c i t y WHSV

Two space velocities, i.e. 0.03 and 0.3 h -1, have been used in the evaluation of catalytic activities of catalysts B and C at 823 K. Figure 6 shows a decrease in activity of the catalyst B when space velocity increases. The accessible sites are saturated at the lowest space velocity. This explains thus the lower conversion levels at a higher space velocity. However, for catalyst C, the evolution of the conversion, which is also depicted in Figure 6, is almost identical for both space velocities. This result could be explained by a better dispersion of the platinum due to the presence of tin.

84 lOO

~

so

~

60

~

40

~

9o

20

o

0

9I n ~

10

20

9

30

40 50 Time (h)

F i g u r e 6. I n f l u e n c e of WHSV on conversion at 823 K (O Pt-Sn/A1PON at 0.03 h -1, 9 Pt-Sn]A1PON at 0.3 h -1, Q P t / A 1 P O N at 0.03 h -1, i Pt/A1PON at 0.3 h "1)

In the presence of tin, the n u m b e r of active p l a t i n u m sites s e e m s to be superior compared to c a t a l y s t B a n d t h u s an increase in the space velocity by a factor of t e n does not seem to s a t u r a t e all t h e sites. T h e s e r e s u l t s show the importance of the role played by tin since p l a t i n u m loading was the same in both cases. It is r e a s o n a b l e to think t h a t in the case of catalyst B, due to the m e t h o d of d e p o s i t i o n , some aggregates of p l a t i n u m are formed on the surface of the c a t a l y s t . In t h e presence of tin, a part of the aggregates could d i s a p p e a r and some P t / S n alloy p a r t i c l e s , b e t t e r d i s p e r s e d at t h e surface, could be formed.

3.4~q. C o m p a r i s o n w i t h MgO F i g u r e 7 compares in function of time the activities of c a t a l y s t B a n d platinum s u p p o r t e d on magnesium oxide MgO, a well-known basic i n d u s t r i a l catalyst. 4O 30 20

10~I~ 0

0

oo oooooooooooo .......

1=====o~p

10

o ~

9

20 30 Time (h)

The MgO support has been synthesised by a c o p r e c i p i t a t i o n r e a c t i o n a n d presents a surface a r e a of 300 m2.g -1 t h a t is almost identical to the surface a r e a of A1PON (310 m2.g-1). P t / M g O h a s b e e n p r e p a r e d by t h e s a m e procedure as for Pt/A1PON. We could see t h a t the activity of Pt/MgO is very low compared to the activity of the Pt/A1PON under the same conditions.

F i g u r e 7. C o m p a r i s o n b e t w e e n O Pt/A1PON and 9 Pt/MgO. 4. C O N C L U S I O N A novel basic support and catalyst have been p r e p a r e d by activation of a l u m i n i u m p h o s p h a t e with ammonia. Fine control of time and t e m p e r a t u r e allows to adjust the O/N ratio of these oxynitride solids and thus to t u n e the acid-base p r o p e r t i e s . The a l u m i n o p h o s p h a t e o x y n i t r i d e s a r e active in K n o e v e n a g e l condensation, but a basicity r a n g e can not yet d e t e r m i n e d . Supporting Pt or Pt/Sn on A1PONs allows to prepare catalysts t h a t are highly active and selective in dehydrogenation reactions.

85 R~~CF~ o

o

3. 4. o

o

o

o

o

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

H. Pines and W.M. Stalick, Base catalyzed reactions of hydrocarbons and related compounds, Academic Press N.Y. 1977. F. Cavani, F. Trifiro and A. Vaccari, Catal. Today, 2 (1991) 11. A. Corma and R.M. Martin-Aranda, J. Catal., 130 (1991) 130. R. Conanec, R. Marchand and Y. Laurent, High Temp. Chem. Process, 1 (1992) 157. P. Grange, Ph. Bastians, R. Conanec, R. Marchand, Y. Laurent, L.M. Gandia, M. Montes, J. Fernandez and J.A. Odriozola, Preparation of Catalysts VI, Elsevier, Amsterdam, p.381, 1994. Ph. Courty, H. Ajot, Ch. Marcilly and B. Delmon, Powder Technology, 7 (1973) 21. F.F. Grekov, J.Guyader, R. Marchand and J. Lang, Rev. Chim. Min., 15 (1978) 341. A. Massinon, J.A. Odriozola, Ph. Bastians, R. Conanec, R. Marchand, Y. Laurent and P. Grange, Appl. Catal., in press. R. Marchand, D. Agliz, L. Boukbir and A.Qu~merais, J. Non-Cryst. Solids, 103 (1988) 35. R. Marchand, Y. Laurent and A. Qu~merais, Rivista della Staz. Sper. Vetro, 5 (1990) 101. J.J. Benitez, M.A. Centeno, J.A. Odriozola, B. Viot, P. Verdier and Y. Laurent, J. Mater. Chem., submitted. H.M. Liao, R.N.S. Sodhy and T.W. Coyle, J. Vac. Sci. Technol. A, 11 (1993) 2681. B.P. Mundy and M.G. Ellerd, Name Reactions and Reagents in Organic Synthesis, Wiley-Interscience, New York, 1988. P. Grange, Ph. Bastians, R. Conanec, R. Marchand and Y. Laurent, Appl. Catal. A: 114 (1994) L191. A. Corma, V. Forn~s, R.M. Martin-Aranda and F. Rey, J. Catal.,134 (1992) 58. R.G. Pearson and R.L. Dillon, J. Am. Chem. Soc., 75 (1952) 2439 Handbook of Chemistry and Physics, Editors R.C. Weast and M.J. Astle, CRC Press, Florida, 61st Edition 1980-1981, D-161 R.D. Cortright and J.A. Dumesic, J. Catal., 148 (1994) 771.

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J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) l lth International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

87

R e a c t i o n K i n e t i c s of t h e H y d r o d e n i t r o g e n a t i o n of D e c a h y d r o q u i n o l i n e o v e r NiMo(P)/A1203 C a t a l y s t s M. Jian and R. Prms Laboratory for Technical Chemistry, Swiss Federal Institute of Technology (ETH), CH-8092 Zurich, Switzerland

ABSTRACT The reaction mechanism and the kinetics of hydrodemtrogenation (HDN) of decahydroqumoline (DHQ) over NiMo(P)/A1203 catalysts were studied in the presence and absence of H2S. Cis- and trans-propylcyclohexylamine were identified as the most important reaction intermediates; their reactivity was found to vary under different reaction conditions. The kinetic constants in the HDN network of DHQ were calculated assuming a Langmuir-Hmshelwood mechanism. It was found that, in the phosphorus-containing catalysts, the adsorption of DHQ was enhanced and the rate constant of the first C-N bond cleavage decreased. As a consequence, the overall HDN reaction rate of DHQ decreased when phosphorus was added to a NiMo/A120,~ catalyst. The effect of H2S was the opposite of that of phosphorus: it increased the rate constants and decreased the adsorption constants. The inhibiting effect of DHQ on its own conversion was different from that on the hydrogenation of cyclohexene, proving that different catalytic sites are involved in these reactions.

1. I N T R O D U C T I O N Hydrodenitrogenation(HDN) is an important process in petroleum refining. It removes nitrogen from oil distillates, so that less NOx pollutes the air when oil is burned and poisoning of the subsequent refining catalysts is reduced when the oil is processed further. Although HDN has been studied intensively and different reaction mechanisms, catalytic active sites, and functions of the catalytic components have been proposed, there are still many questions to be answered in order to better understand the reaction and the catalyst (1-4). Three kinds of reactions are revolved in the HDN process (5): C-N bond cleavage, hydrogenation of the aromatic ring, and hydrogenation of the nitrogencontaining aromatic heterocyclic ring. C-N bond cleavage is usually the final and

88 most important reaction step m the HDN process, since the nitrogen atom can only be removed after the C-N bond is broken. However, in most kinetic and mechanistic studies, lumped Langmuix-Hinshelwood kinetics were used, and it was assumed that all HDN reaction steps (hydrogenation of aromatics and olefins, elimination of ammonia, etc.) take place on the same catalytic site (1). As a consequence, the functions of the catalytic components of the catalyst in the ldnetic network could not be completely unravelled; m a n y assumptions had to be made to simulate the ldnetic network, and the resulting kinetic constants differed from one author to the other (6, 7). In the widely used industrial NiMo-P/A120,~ HDN catalyst, the promotional effect of phosphorus has been demonstrated by its industrial performance. In model HDN compound studies, however, phosphorus promoted the HDN of quinolme (8), but decreased the HDN of decahydroqumoline and piperidme (9). A negative effect of phosphorus has also been observed on hydrogenation (10). The question, therefore, arises as to the role of phosphorus in an HDN catalyst. In the present study, the HDN of decahydroquinoline (DHQ) was studied over NiMo(P)/A120.~ catalysts in the presence and absence of H2S. The reaction took place at 593 K and 3.0 MPa, thus allowing us to observe the most important reaction intermediate, propylcyclohexylamme, and to calculate the kinetic constants from the experimental results. Rate and adsorption constants for the different reaction steps were determined by separate and by combined HDN studies of DHQ and cyclohexene.

2. E X P E R I M E N T A L The NiMo/A120:, catalysts (3 wt% Ni, 8 wt% Mo and 0 or 2 wt% P) were prepared by the incipient wetness impregnation method as described in ref. (10). The HDN reactions were carried out in a continuous-flow microreactor. A sample of 0.1 g catalyst diluted with 9.5 g SiC was first sulphided in situ with a mixture of 10% H,~S and H2 at 643 K and 1.5 MPa for 4 h. After sulphidation the pressure was increased to 3.0 MPa, and a solution of the reactant in n-octane was fed to the reactor by means of a high pressure pump. The initial reactant concentration (Ao) was adjusted by changing the reactant concentration. Dimethyldisulphide was added to the solution to generate H2S m the reaction stream (P(H2S)=6.53 kPa). The influence of the NH.~ product was determined by co-feeding pentylamine as a source of NH3 (P(NH3)=l.59 kPa). Reaction products were analyzed by on-line gas chromatography with a Shimadzu GC-14A gas chromatograph equipped with a 50 m CP Sil-5 fused silica capillary column and a flame ionization detector. Reaction intermediates were identified by GC-MS. Samples were taken after 50 h on stream when the activity of the catalyst was stable, with n-nonane and n-dodecane as internal standards. Space time was defined as ~ = e-Vc,d v ~ , where ~ is the void fraction of the

89

catalyst bed (e was assumed to be 0.4), Vat is the catalyst volume, and v$~ is the volume flow rate of the gas phase reactant.

3. R E S U L T S AND DISCUSSION The HDN reaction network of qumolme is shown in Fig. 1 (7, 8). A critical reaction step in this network is the breaking of the C-N bond through the consecutive THQ I ~ O P A ~ H C or D H Q ~ H C reaction (HC= hydrocarbons). Since equilibria can easily be established between Q, THQ5, THQ 1, and DHQ, all three main types of reactions in the HDN process (C-N bond cleavage, aromatics hydroge-nation, and aromatic heterocycle hydrogenation) are usually involved m the HDN of quinoline-type compounds. It was found that the rate constants of some reaction steps (such as T H Q 5 ~ D H Q and DHQ~PCHE) were of the same magnitude, so that a simple rate-limiting step treatment is not always possible (7, 11). Disagreement exists as to whether the HDN reaction proceeds mainly via THQ 1-->OPA~HC or via DHQ--->HC (1, 12-14). In almost all k~netic studies, several reactions were grouped together; thus the calculated ldnetic constants did not give a clear indication of reaction mechanism and rate-limiting steps and even less so of the catalytic sites and functions of the catalytic components. One of the most important reaction intermediates PCHA, has been observed only rarely (6). Q

THQ-1

OPA

It

It

It

THQ-5

DHQ

PCHA

PB

i

s,, o~

PCHE

,,-...., 1l PCH

Figure 1. HDN reaction network of quinolme-type compounds. Q=quinolme, THQ5=5,6,7,8-tetrahydroquinoline, DHQ=decahydroqumoline, THQ l=l,2,3,4tetrahydroqumi~ne OPA=ortho-propylaniline, PCHA=2-propylcyclohexylamine, PCHE=propylcyclohexene, PCH=propylcyclohexane, PB=propylbenzene. 3.1. T h e H D N r e a c t i o n n e t w o r k o f d e c a h y d r o q u i n o l i n e DHQ is an important intermediate in the HDN network of quinolme; only C-N bond cleavage is needed to remove the nitrogen atom from DHQ. However, since dehydrogenation of DHQ proceeds qmte fast, THQ5, THQ1, and Q are usually present as well, thus making the study of the HDN of DHQ difficult.

90 Table 1 Product compositions (%) in the HDN of DHQ at 593 K and 3.0 MPa (~= 0.20 sec) catalyst

H,_,S PCH PCHE

PB

PCHA DHQ T H Q 5 0 P A

Q

THQ1 others

NiMo

yes

1.81

2.20

0.02

2.89

89.1

2.85

0.03

0.15

0.26

0.72

NiMoP

yes

1.33

2.09

0.02

2.34

89.0

3.38

0.04

0.10

0.34

1.47

NiMo

no

0.57

0.78

0.07

0.16

92.6

4.43

0.04

0.18

0.35

0.86

NiMoP

no

0.81

1.21

0.13

0.16

91.7

4.38

0.08

0.25

0.42

0.82

In the present study our reaction system and sensitive analytic technique allowed us to perform the HDN reaction of DHQ under such reaction conditions t h a t only small amounts of Q, THQ-1, and OPA were formed (Table 1). This indicates that dehydrogenation of the carbocyclic ring of DHQ was slow and could be neglected. Therefore, the reaction network can be simplffied as in Fig. 2. THQ-5

DHQ

PCHA

KA

KB

PCHE

PCH

KC

Figure 2. Simplified HDN reaction network of decahydroqumoline The mass balance calculation showed that the performance of the reactor was reliable, and there were no diffusion limitations. The operating conditions (low temperature and high space velocity) allowed us to identify both c- and t-PCHA and to confirm t h a t they are the reaction intermediates between c- and t-DHQ and PCHE (Fig. 2) (gas chromatography gave three peaks which, according to the mass spectrum, all belong to PCHE). The mass spectrum of t-PCHA consists of a strong peak at m/e=141 (molecular ion) and small peaks at m / e = l l l and 56, while the mass spectrum of c-PCHA consists of a middle-sized peak at m/e=141, a strong peak at m/e=56, and small peaks at m/e=l 11, 98, 70, and 43. A further analysis of the steric isomers of PCHA (Table 2) shows that the reaction of c-DHQ to c-PCHA was faster than t h a t of t-DHQ to t-PCHA, that the isomerization of c- and t-DHQ was fast, and the isomeration of c- and t-PCHA was somewhat slower. The presence of PCHE in the HDN product (Table 1) indicates that at least part of the HDN reaction of PCHA proceeds through elimination of ammonia r a t h e r than by direct hydrogenolysis. Nevertheless, the direct product of the elimination reaction (allylcyclohexylamine) was not observed. This must be due to its strong adsorption and fast hydrogenation to PCHA.

91

Table 2 Ratios of trans-/cis-isomers at different space times over a NiMo/A1,_,O:3catalyst * space time (ms)

0

38

57

112

203

340

512

t-/c-PCHA

/

2.3

3.0

3.5

3.9

4.5

4.7

0.3

5.3

6.2

6.5

6.7

6.8

6.8

t-/c-DHQ

* 593 K, 3.0 MPa, H,_,S/H2 - 3.0x10 3 mol/mol. The concentration of PB m the reaction products (Table 1) is too high to be accounted for by dehydrogenation of PCHE, especially m the absence of H2S, since at thermodynamic equilibrium the PCH/PB ratio should be greater than 50 (15) under our experimental conditions. Furthermore, no toluene was observed m the simultaneous reaction of methylcyclohexene and DHQ under the same reaction conditions. Therefore, there must be another reaction path to account for the formation of PB m the HDN products of DHQ. This second reaction path can only be the reaction of D H Q ~ T H Q I ~ O P A ~ HC. It has been demonstrated that a relatively high concentration of PB is present in the HDN of OPA due to the direct hydrogenolysis of the C(spe)-N bond of OPA (16). 3.2. I n f l u e n c e o f s p a c e t i m e a n d i n i t i a l c o n c e n t r a t i o n The conversion of DHQ to C-N bond cleavage products as a function of space time (Fig. 3A) demonstrates t h a t PCHA is a p r i m a r y HDN product and that PCH is a secondary HDN product. ) 1 PCH J P C H E

A

0.20

0.15 PCHA x"

o- r ,x~ (..-

~" 0.05 O.,F~ 0.0

._,

r 0.2

. •

, 0.4

space time, s

~ PB 016

o . o o .,.

0.0

0:1

0:2 0:3 0:4 space time, s

0:5

0.6

Figure 3. Effect of space time on the HDN of DHQ at 593 K, 3.0 MPa and P(H,,S) = 6.53 kPa. A: product composition, B" first-order kinetic fitting. The ~ "" t plot (XDHQ -- PCHA 4- PCHE + PCH + PB) could be fitted with a straight line (Fig. 3B), but, under our reaction conditions (small x), this does not distinguish between first and zero order. Another way to distinguish

92 between first and zero order reactions is to measure the effect of the initial concentration on the rate. At low conversion (XDHQ < 0.1), the concentration of DHQ in the reaction stream can be regarded as constant, that is A,~,K~ = Ao. Therefore, assuming t h a t a Langmuir-Hinshelwood mechanism applies, t h a t all the reactants, intermediates, and products adsorb on one and the same site, and t h a t the adsorption of hydrocarbons, NH3, and HeS can be neglected, we obtain dA dt

k,KAA I+ZK~I

k,KAA I+K AAo

where A stands for the concentration of DHQ (Ao is the imtial concentration), k, is the rate constant, and KA is the adsorption constant of DHQ (cf. Fig. 2). The slope of the - I n ( 1 - X D H Q ) "~ t plot is equal to k~KA/(I+K.~Ao). By varying Ao, k, and KA can be calculated from the slope" against Ao plot. The plots are shown in Fig. 4, and the resulting kinetic constants are given in Table 3. The adsorption constants KA show that, in the present concentration range (Ao > 2 kPa), the order of the reaction is between zero and one. 25

25

2O

20, v =-

15

NiMoPIAI203

15

o v +

+

"" 0

|

,

,

,

2

4

6

8

Ao, kPa

5

10

10

Ao, kPa

Figure 4. Graphs used in the determination of the kinetic constants of the HDN of DHQ at 593 K and 3.0 MPa. I1: with H2S,o: no H2S. As we have already seen from the HDN product yields (cf. Table 1 for the compositions after 0.2 s), phosphorus has a negative effect on the HDN of DHQ in the presence of H2S but a positive effect in the absence of H2S. The kinetic constants (Table 3) show that introducing phosphorus to a NiMo/A120:3 catalyst increases its adsorptivity toward DHQ but decreases the rate constant of the first C-N bond cleavage in the presence of H2S. The effective rate constant k,K.~/(I+K~) of the first C-N bond cleavage (rate limiting step) in the presence of H,_,S is lower for the NiMoP/A1,_,O~ catalyst, which accounts for the negative effect of phosphorus. Likewise, the effective rate constant is higher for the NiMoP/A120~ catalyst in the absence of HeS.

93 Table 3. Kinetic constants of the HDN of DHQ at 593 K, 3.0 MPa catalyst

HeS

kl

KA

k2KB

k.3Kc

NiMo/A1203

yes

2.6

0.4

25

23

NiMoP/A120:3

yes

1.8

0.5

26

34

NiMo/A1903

no

0.4

1.4

360

/

NiMoP/A1203

no

0.6

2.0

540

130

The reaction p a t h D H Q ~ T H Q I ~ O P A ~ H C must be taken into account to explain the promotional effect of phosphorus in the absence of H2S. A strong promotional effect of phosphorus has been observed for the HDN of OPA over NiMo/A1,.,O~ catalysts, which could be explained by the larger adsorption constant of OPA on the P-containing catalyst. The HDN activity of OPA was even higher in the absence of H~S (16). If we assume that all the PB produced in the HDN of DHQ under our present reaction conditions was formed through the reaction of OPA, and that the PCH/PB ratio is the same for the reaction path through OPA in the HDN of DHQ and in the HDN of pure OPA, then the HC/PB ratios from the HDN of OPA can be applied to the HDN of DHQ under the same reaction conditions; the relative contributions of the two reaction pathways can be estimated. The results show that about 40% of the HDN reaction of DHQ takes place through the reaction path D H Q ~ T H Q I ~ O P A ~ H C in the absence of H,.,S but less than 10% in the presence of H,_,S. The very low concentration of PCHA (compared with that in the presence of H,_,S) also indicates that the rate limiting reaction steps might have changed in the absence of H2S.

3.3. A d s o r p t i o n c o n s t a n t s from h y d r o g e n a t i o n of CHE The inhibiting effect of DHQ and its NH~ product was studied on the final step in the network of Fig. 2, the alkene hydrogenation. To avoid confusion with the PCHE olefm formed from DHQ, cyclohexene (CHE) was used as the reactant, and pentylamine (PA) was used as the source of NH~. When the hydrogenation of CHE is performed in the presence of NH~, we have dC dt

kcKcC I+KcC+KNN

kcKcC I+KNN

(C stands for CHE and N for NH~)

Assuming that the adsorption of NH.~ and the hydrogenation of CHE take place on the same catalytic site, and that the adsorption of NH3 is much stronger than that of CHE, we obtain, with N - No = constant,

94

-In(1 - x c ) =

kcKc 9t l+KNN o

where kc is the rate constant and Kc the adsorption constant of cyclohexene. kcKc can be calculated from the first order hydrogenation of CHE alone, while KN can be calculated from the slope of the -ln(1-XCHE)~t plot of the hydrogenation of CHE in the presence of NHa.

.

NiMo/AI203

NiMoPIAI203

/

/

6

6

/

/ /

/

,V/

~//

!

v

2 0

0.0

"~" 2

0:2

024 space time, s

"

!

o 0.o

026

,/

-

,

-

-

i

O.2 0.4 space time, s

9

-

i

0.6

Figure 5. Hydrogenation of CHE at 593 K, 3.0 MPa. 9in the presence of H,_,S, ...... "in the absence of H,_,S; i - C H E alone, I : CHE+PA ( P A - 1.59 kPa) At small space times, the C5 hydrocarbons could not account for all the PA converted, although no PA was found in the reactor outlet; a similar observation was reported by La Vopa and Satterfield (17). But at higher space time the C5 hydrocarbons do account for more t h a n 90% of the PA converted. Thus only the data at very high space time were used to calculate KN (Fig. 5). The adsorption constants of DHQ were obtained in the same way. The m a s s balance of DHQ and CHE was always good. The resulting adsorption constants are given in Table 4. Table 4. Adsorption constants (kPa-~) obtained from the hydrogenation of CHE with H2S

no H2S

KDHQ

Kr,m

NiMo/A120.~

2.7

1.0

7.0

3.5

NiMoP/A1,_,Oa

4.1

1.7

11.0

4.8

catalyst

KDHQ

Kr,m

Table 4 shows t h a t the P-containing NLMoP/A1203 catalyst favours the adsorption of NH~ and of DHQ. The adsorption of NH~ is strongly enhanced m the absence of H2S for both catalysts. The effect of phosphorus, which is applied

95 in the form of phosphate, is certainly not restricted to an increase or decrease in the number of sites, since then only the rate constants, and not the adsorption constants, should have changed. Apparently, phosphorus changes the chemical properties of the catalytic sites. A comparison of the results in Tables 3 and 4 shows t h a t the adsorption constants differ substantially when different reactions are inhibited by the same molecule. The adsorption constants KDt~ are much smaller when they are determined from the HDN of DHQ itself than from the inhibition effect of DHQ on CHE. This confirms that different catalytic sites are needed for these chemically different reaction steps m the HDN process. This is also supported by the observation t h a t H2S promotes the C-N bond cleavage reaction of DHQ and piperidine, while it inhibits the hydrogenation of alkenes (10, 12). Therefore, one and the same catalytic site can not be responsible for these different reaction steps in the HDN kinetic studies. It also means that care should be exercised with the choice of the molecule and reaction when determining the adsorption constants from the inhibiting properties of compounds on certain reactions (17), because different catalytic sites might be involved for different reactions (HDN, HDS, hydrogenation, etc.). The presence of different catalytic sites for the first C-N bond breaking of DHQ and the hydrogenation of CHE is confirmed by the -ln(1-x~liQ) versus -In(1-x(nt~:) plot. In the simultaneous reactions of A and B, in which A and B are both adsorbed on the same catalytic site and follow a Langmuir-Hmshelwood mechanism, we have dA kAKAA dt = 1+ ZK,I

-

and

dB kBK~B - d--t-= 1+ ZK~I '

and thus

kAKA ln(l

In(l- x A) = kBK~

x B)

A plot of ln(1-xA) ~ In(1-xs) should result in a straight line if the assumption is true. Figure 6 shows the result of simultaneous reactions of DHQ and CHE. The curvature of the ln(1-XDHQ) ~ ln(1-XCHE) plot confirms that the adsorption sites for CHE and DHQ over the NiMo(P)/AI,,O~ catalysts are not the same.

1.0 0.8 W

-,-0.6 ,x0

Figure 6. Simultaneous reaction of DHQ and CHE at 623 K, 3.0 MPa, and H2S/I-I2 = 3.0x10 3 (tool/tool) over the NiMoP/Al~O3 catalyst.

~. 0.4 0.2 0.0 . . . . . . . . . . 0.0

0.1

0.2 0.3 -In(1 "•

0.4

0.5

95 4. C O N C L U S I O N S The present results show that the separate steps in an HDN reaction network can not be lumped together into one kinetic equation. The intermediate reactions may take place on different catalytic sites which differ m their ability to bind reactants, intermediates, and products. Phosphorus was found to modify the rate constants as well as the adsorption constants of the HDN reaction steps, indicating th at it changes both the number and nature of the active sites of NiMo/AI203 catalysts. The crucial reaction intermediate PCHA in the HDN network of quinolmetype compounds has been clearly observed. Formation of cis-PCHA was faster than that of trans-PCHA, but isomerization was relatively rapid. The presence of H2S m the reaction stream favours the cleavage of the first C-N bond in DHQ, but slows down the C-N bond cleavage in PCHA. The presence of H2S decreases the adsorption constants of DHQ and NH.~. It is concluded that ~-40% of the HDN reaction of DHQ takes place through the reaction path of D H Q ~ T H Q I ~ O P A ~ H C at 593 K and 3.0 MPa in the absence of H2S, while less than 10% takes place m the presence of H,_,S.

REFERENCES

1. M. J. Girgis and B. C. Gates, Ind. Eng. Chem. Res., 30 (1991) 2021. 2. H. Schulz, M. Schon and N. M. Rahman, Stud_ Surf. Sci. Catal., 27 (1986) 201. 3. R. Phns, V. H. J. de Beer and G./~ Somorjai, Catal. Rev. Sci. Eng., 31 (1989) 1. 4. T. C. Ho, Catal. Rev. Sci. Eng., 30 (1988) 117. 5. G. Perot, Catal. Today, 10 (1991) 447. 6. S. S. Shah, K. N. Mathur, J. R. Katzer, H. Kwart and A. B. Stiles, Prepr. Am. Chem. Soc., Div. Pet. Chem., 22 (1977) 919. 7. C. N. Satteriield and S. H. Yang, Ind. Eng. Chem. Proc. Des. Dev. 23 (1984) 11. 8. S. Eijsbouts, J. N. M. van Gestel, J. A. R. van Veen, V. H. J. de Beer and R. Prms, J. Catal., 131 (1991) 412. 9. M. Jian and R. Prms, Catal. Letters, 35 (1995) 193. 10. M. Jian, J. L. Rico Cerda and R. Prins, Bull. Soc. Chim. Belg., 104 (1995) 225. 11. J. L. Rico Cerda and R. Prms, Bull. Soc. China. Belg., 100 (1991) 815. 12. C. N. Satterfield and S. Gulteldn, Ind. Eng. Chem. Proc. Des. Dev. 20 (1981) 62. 13. S. H. Yang and C. N. Satterfield, Ind. Eng. Chem. Proc. Des. Dev. 23 (1984) 20. 14. K. S. Lee, H. Abe, J. A. Reimer and A. T. Bell, J. Catal., 139 (1993) 34. 15. J. F. Cocchetto and C. N. Satterfield, Ind. Eng. Chem. Proc. Des. Dev., 20 (1981) 49. 16. M. Jian and R. Prms, Catal. Today, in press. 17. V. La Vopa and C. N. Satterfield, J. Catal., 110 (1988) 375.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

97

Effect of SpiUover Hydrogen on Amorphous Hydrocracking Catalysts A.M. Stumbo*, P. Grange and B. Delmon Universit6 Catholique de Louvain - Unit6 de Catalyse et Chimie des Mat6riaux Divis6s, Place Croix du Sud, 2/17 - 1348 Louvain-la-Neuve, Belgium

Abstract The cracking of diphenylmethane on mixtures of sulfided CoMo/SiO 2 and amorphous silica-alumina particles was studied. The products were benzene and toluene. The addition of CoMo/SiO 2 to silica-alumina strongly increases the cracking rate and the OH-OD exchange, and diminishes the amount of coke formed. This is interpreted by a spillover of dissociated H 2 (or D 2) onto the silica-alumina, with spillover hydrogen forming Br6nsted sites and reacting with coke precursors, and spillover deuterium exchanging with the hydroxyls.

1. I N T R O D U C T I O N This work is a contribution to the understanding of the effect of spillover hydrogen in a type of catalyst of considerable industrial importance, namely that composed of transition metal sulfides and amorphous acidic solids. This is typically the case of sulfided CoMo supported on silica-alumina used for mild hydrocracking. Hydrocracking catalysts possess two functions, they consist of a hydrogenationdehydrogenation component (noble metals or transition metal sulfides) and an acidic support (silica-aluminas or zeolites). In the mechanism generally accepted to explain the hydrocracking reaction (the so-called "ideal hydrocracking"), an alkane is initially dehydrogenated to an olefin on the metal phase and then adsorbed on an acidic site, where it is converted to an alkylcarbenium ion. The latter, often after a rearrangement to a more stable form, is cracked via a 13-scission mechanism, forming a lighter olefin and an ion that are hydrogenated to the corresponding paraffins [1 ]. Spillover is a phenomenon that involves the formation of an active species on one phase and the migration of this species onto another phase which is not able to form this species itself. It has been demonstrated that spillover hydrogen (Hsp) can create protonic acidic sites on many kinds of solids, enhancing their activity for a variety of acid catalyzed reactions. This is the case of benzene cracking on silica activated with Pt/A120 3 [2], butane isomerization over Pt/SOn2--ZrO2 [3], cracking of n-hexane [4] and n-heptane [5] on erionite and toluene disproportionation on Fe-HY [6]. Nevertheless, many studies concerning hydrocracking catalysts do not allude to a possible influence of spillover, because they are not * Acknowledges the financial support from the CNPq (Conselho Nacional de Desenvolvimento Cientffico e Tecnol6gico - Brazil) and the Federal Science Policy Office of Belgium (IPA Program).

98 conceived in a way that allows the detection nor the investigation of that phenomenon, since they employ model molecules that need the simultaneous presence of both functions mentioned above to react. In order to investigate the effect of Hsp on amorphous silica-aluminas, we selected diphenylmethane (DPM) as model molecule. The cracking by the "ideal hydrocracking" mechanism is not possible, because DPM does not offer the possibility of a 13-scission. Since the mechanism by which this molecule reacts does not seem to clearly fit into a precise category, we shall refer to it throughout this work using the general term "cracking". Hattori et al. [7], studying DPM cracking over a large variety of catalysts, found three possible routes for this reaction: (i) via carbocations intermediates, formed on acidic sites, producing mainly benzene and toluene and also diphenylethane as a polymerization by-product; (ii)through hydrogenation of both aromatic rings, followed by the cleavage of the C-C bond, producing cyclohexane and methylcyclohexane, the latter undergoing further cracking to produce methane; (iii) via simple cleavage of C-C bond, forming benzene and toluene. The products of the reaction on silica-aluminas, Mo/SiO 2 and Ni-Mo/A1203 catalysts matched those of route (i). Route (ii) is typical of reduced Fe and Ni supported on silica, alumina, titania or zirconia and route (iii) of the same catalysts in the sulfided form. These observations were confirmed by Shimada et al. [8,9]. When catalysts like those used in this work are employed, the previous hydrogenation of the aromatic ring is not necessary to the cracking of DPM and carbocation formation on acidic sites is the main reaction pathway. This reaction does not correspond to a bifunctional mechanism. DPM is thus a suitable tool to study specifically a possible influence of spillover hydrogen on the acidity of silica-aluminas. Our method was to physically isolate the two functions of the catalyst, i.e., the acidic function and the functions necessary for hydrogen activation (hydrogenation-dehydrogenation and, possibly, production of spillover hydrogen). For that, we used mechanical mixtures of amorphous silica-aluminas of different compositions with a CoMo/SiO 2 catalyst. Silica, a support with only a very weak acidity, has been chosen to avoid that the acidity of the support of the donor contribute significantly to total acidity and thus cause interferences with the reactions of the acidic phase. XPS measurements were used to check whether the phases remained unchanged after mixing and sulfiding. The effect of Hsp will be revealed by the comparison between the activities of the pure phases and those measured for mechanical mixtures of both components in different proportions. Since spillover phenomena have been most directly sensed through the use of IR in OHOD exchange [10] (in addition, in the case of reactions of solids, to phase modification), we used this technique to correlate with the catalytic results. One of the expected results of the action of Hsp is the enhancement of the number of Br6nsted sites. FTIR analysis of adsorbed pyridine was then used to determine the relative amounts of the various kinds of acidic sites present. Isotopic exchange (OH-OD) experiments, followed by FTIR measurements, were used to obtain direct evidence of the spillover phenomena. This technique has already been successfully used for this purpose in other systems like Pt mixed or supported on silica, alumina or zeolites [10]. Conner et al. [11] and Roland et al. [12], employed FFIR to follow the deuterium spillover in systems where the source and the acceptor of Hsp were physically distinct phases, separated by a distance of several millimeters. In both cases, a gradient of deuterium concentration as a function of the distance to the source was observed and the zone where deuterium was detected extended with time. If spillover phenomena had not been involved, a gradientless exchange should have been observed.

99 2. EXPERIMENTAL 2.1. Preparation of the mechanical mixtures Three commercial silica-aluminas were used as acidic phases: we call them SA6 (6.5 wt.% A1203 , 500 mE.g "l, average pore diameter: 80 .~), SA12 (12 wt.% A1203, 500 mE.g-1, 80 ~,) and SA60 (60wt.% A1203, 500mE.g -1, 66 .~,). The precursor of the Hsp generator, a CoMo/SiO 2 catalyst (14 wt.% MoO 3 and 3% CoO, 220m2.g ~ 115 .~), was prepared by successive impregnation. Silica (Kali-Chimie AF-125, 270m2.g -1, 115 .~) was first impregnated with an aqueous solution of cobalt acetate (Merck, ultra pure) and subsequently with an aqueous solution of ammonium heptamolybdate (Merck, ultra pure). After each impregnation step the sample was dried overnight at 393 K and calcined at 673 K for 2 hours, under a stream of air (Air Liquide, S). Mechanical mixtures were prepared according to the following procedure, in order to insure intimate mixing and good mutual contact. The pure phases were grounded and sieved to obtain particles of sizes under 40 ~tm. The powders, mixed in the desired proportions, were suspended in n-pentane (15 ml/g solid), placed in an ultrasonic bath for 5 rain and then vigorously mechanically stirred (Ultra Turrax T-50, 3000 rpm) for 10 rain. The n-pentane was evaporated at room temperature, under a flux of Ar and continuous magnetic stimng. After drying at 393 K overnight, the powder was pressed (10 ton.cm-2), grounded and sieved to obtain particles between 0.315 and 0.5 mm. The pure phases were submitted to the same treatment. The samples will be identified by their relative weight content of silica-alumina, named R m, defined as:

Rm _ -

wt. % SiO 2 - A1203 xl00 wt. % CoMo / SiO 2 + wt. % SiO 2 - A1203

(1)

All samples were sulfided in situ prior to catalytic tests and characterizations. A flow (100 ml.min -l) of argon (Air Liquide, N46) was first established, the temperature raised to 423 K, at 10 K.min -1, and maintained at this value for 30 min. The gas was then changed to a mixture of 15% (vol.) HES (Air Liquide, N28) in H 2 (Air Liquide, N30), at the same flow rate. The temperature was raised to 673 K, at l0 K.min -1, and kept at that level for 2 hours. 2.2. X-Ray Photoelectron Spectroscopy (XPS) measurements A Surface Science Instruments SSX-100 spectrometer (model 206), equipped with an aluminum anode whose radiation was monochromatized (AIKct, 1486.6 eV) and focalized, was used. The positive charge developed at the surface of the samples was compensated with a charge neutralizer adjusted at an energy of 8 eV. The sulfided samples were pressed and transferred to the spectrometer protected by a meniscus of iso-octane, to avoid oxidation by atmospheric oxygen. The lines corresponding to C ls, OEs, SiEs, SEp, A12p, MO3d, SEs and Co2p3/2 were analyzed. Their binding energies were determined taking the position of the C is line, corresponding to carbon in a C-C or C-H environment, as 284.8 eV. The intensities were estimated by calculating the area of each peak. Apparent atomic concentrations, that take into account the amount and the dispersion of the elements on the surface, were calculated using the sensibility factors determined experimentally by Weng et al. [13] for our spectrometer. The results were compared to the

100 theoretical values calculated considering that a simple "dilution" of the pure phase containing the metals takes place, according to the following expression: oorotic ,

=

(2)

2.3. Catalytic activity tests The cracking of diphenylmethane (DPM) was carried out in a continuous-flow tubular reactor. The liquid feed contained 29.5 wt.% of DPM (Fluka, >99%), 70% of n-dodecane (Aldrich, >99%; solvent) and 0.5% of benzothiophene (Aldrich, 95%; source of H2S, to keep the catalyst sulfided during the reaction). The temperature was 673 K and the total pressure 50 bar. The liquid feed flow rate was 16.5 ml.h -1 and the H 2 flow rate 24 1.h-1 (STP). The catalytic bed consisted of 1.0 g of catalyst diluted with enough carborundum (Prolabo, 0.34 mm) to reach a final volume of 4 cm 3. The effluent of the reactor was condensed at high pressure. Liquid samples were taken at regular intervals and analyzed by gas chromatography, using an Intersmat IGC 120 FL, equipped with a flame ionization detector and a capillary column (Alltech CP-Sil-8CB). The results of the catalytic tests were expressed as DPM total conversion. These experimental results were compared to theoretical values (Ct), calculated considering, as an approximation, a zero-order reaction and the absence of interactions of any kind between both phases, according to the following expression:

ct - Rm 100 XCRm=l + E1- ~Rm ] x CRm=0

(3)

where CRm=I and CRm=0 are, respectively, the experimental conversions corresponding to pure silica-alumina and pure CoMo/SiO 2. After 24 h of reaction, the catalytic bed was retrieved and sieved to separate the catalyst from the diluent. The used catalyst particles were placed in a Soxhlet apparatus, washed with n-hexane for 8 hours and then dried overnight at 393 K. Their carbon content was determined by automatic titration of the CO 2 formed by burning the washed sample, in a Str6hlein Coulomat 702 apparatus. 2.4. FTIR of adsorbed pyridine The samples were ground and pressed (2 ton.cm -2, for 15 s) in the form of 13 mm diameter wafers, weighing between 3 to 5 mg. They were placed in a specially designed cell, that allowed the heating of the sample under vacuum or controlled atmosphere. IR spectra could be taken through N aCI windows. The samples were submitted to the sulfidation procedure described above, followed by 2 h of heating at 673 K, under vacuum (about 2x10 -3 Pa). After cooling under vacuum, pyridine was adsorbed at room temperature for 30 minutes. The samples were then outgassed in three steps of 1 h: the first one at room temperature and the others at 423 K and 523 K. Spectra were taken before pyridine adsorption and after each outgassing step, with a FTIR spectrometer Bruker IFS-88 (spectral resolution set at 1 cm -1). Each spectrum represented the average of at least 50 scans.

101 The amount of Brrnsted sites was evaluated by measuring the surface of the characteristic band at 1540 cm -1. Corrections have been made to take into account the differences in weight and surface of the wafers. After each test, the wafer was weighted and its cross section was measured with a planimeter. The results were corrected to represent those of a "standard wafer" (A c) of 5 mg and 25 units of area, according to the following expression: 25 5 A c = A e x ....... x Sw mw

(4)

where A e is the experimental integral absorbance of the band considered, S w is the cross section of the wafer (in arbitrary units) and m w is the weight of the wafer (mg).

2.5. Isotopic exchange (H-D) experiments Wafers were prepared, sulfided and evacuated (2 h at 673 K) as described above. The temperature was then set at 423 K and 80 kPa of purified deuterium (Air Liquide, N28) was admitted into the cell. The purification procedure consisted of passing the gas through a moisture filter (Chrompack Gas Clean 7971), an oxygen filter (Chrompack Gas Clean 7970) and a liquid nitrogen trap. Several spectra were taken at regular intervals, using the analysis conditions mentioned above. Before each measurement, the samples were cooled to room temperature. Preliminary experiments had shown that no exchange took place at that temperature. The amount of deuterium exchanged was measured by the total area of the OD bands situated between 2800 and 2100 cm-1. These results were also corrected according to Equation 4.

3. RESULTS 3.1. XPS measurements No new peaks were observed in the mechanical mixtures. The binding energies of all elements were the same in the pure phases and in the mixtures [14]. Figure 1 shows the apparent atomic percentages of molybdenum and cobalt, as given by XPS, on the surface of the sulfided pure phases and mechanical mixtures. In both cases, the experimental results are close to the theoretical values calculated according to Equation 2. ,,

O

,i,

.....

,

O

0.5

"i.... I.. i

0

50 Rm

" "iii

lOO

0-..l...

0

I

0

50

""m

100

Rm

Figure 1. Molybdenum (left) and cobalt (fight) apparent contents (atomic %), determined by XPS, in SA6 (m), SA12 ( . ) and SA60 (o) series, compared to the theoretical values calculated by Equation 2 (dashed lines).

102

3.2. Catalytic activity tests The main products of diphenylmethane (DPM) cracking were benzene and toluene. Very small amounts of polymerized by-products have been found (< 0.5%), but no cyclohexane or partially hydrogenated compounds like cyclohexylphenylmethane were detected. Figure 2 shows the conversions obtained with the three series studied, as a function of the mechanical mixtures composition, one hour after the beginning of the reaction and at the steady-state. Each series presents a maximum of activity, but at a different composition. SA6 series has a maximum between R m values of 50 and 75, whereas SA12 series has a maximum around R m = 50, and SA60 series near R m = 75. The dashed lines on the figures represent the sum of the individual contributions of the pure phases, calculated according to Equation 3. A very important synergetic effect is observed in all series, i.e., the activity of the mixtures is considerably higher than the calculated values (increase by 200% to 750%). 100 Pure silica-aluminas are strongly deactivated, losing about 80% of their activity before reaching the steady-state. The loss in 50 pure CoMo/SiO 2 catalyst is much less pronounced (about 15%). Mechanical mixtures represent an intermediate case; they 0 lose between 35% and 50% of their activity. Table 1 shows the experimental carbon 100 contents of the used samples, compared to the theoretical values obtained by adding the O "~ 50 k.. contributions of the individual pure phases, ~D taking into account their proportions in the O mixtures. All mechanical mixtures present r,.) 0 experimental carbon contents considerably lower than these calculated values. 100

3.3. FTIR of adsorbed pyridine 50

0

50

100

Rm

Figure 2. DPM conversion as a function of the mechanical mixtures composition, after 1 h (..) of reaction and at the steadystate (e), compared to the theoretical values calculated by Equation 3 (dashed lines).

Figure 3 shows the amount of Brrnsted sites, as measured by the surface of the characteristic IR peak at 1540cm -1 after outgassing at 523 K, as a function of the composition of the mechanical mixtures. The dashed lines represent the addition of the contribution of the pure phases, calculated as in Equation3. An enhancement of the amount of Br/Snsted sites on the mixtures, when compared to the theoretical values, is observed. This effect is not very clear in SA6 series, but it is more evident in SA12 and SA60 series. The reproducibility of the experiments has been checked; the variation between different wafers of the same sample was always inferior to 10%.

103

Table 1 Carbon content (wt.%) after catalytic test SA12 Series SA6 Series Theor. Experim. Rm Experim. Theor. 0.45 0 0.45 25 11.3 7.3 50 75 16.4 26.3 22.1 100 34.9 -

70

0.6

SA60 Series Experim. Theor. 0.45 2.3 6.7 4.5 8.4

12.9 19.0

25.2

-

3.4. I s o t o p i c e x c h a n g e e x p e r i m e n t s

All the samples showed two kinds of OD bands: a relatively sharp peak 35 0.3 around 2760 cm ~ and one broad band at lower frequencies, whose limits were always within approximately 2700 and 00 2100 cm-1. The former is assigned to isolated OD species and the latter to OD 60 0.8:5 species in interaction [ 15]. "=" 1SA12 II ] The pure silica-aluminas exchanged O the lowest amounts of deuterium among = 30 0.4 all the samples. Their OD bands were not very intense and increased very :0 9 O0 slowly. The pure CoMo/SiO 2 catalyst, where a potential source of spillover 130 0.6 hydrogen is present, exchanged a higher amount of deuterium than the pure silica-aluminas. The three series of 65 0.3 mechanical mixtures had bands located exactly at the same positions as those of the pure phases. However, their 0-0 intensities were considerably higher. 0 50 100 After 20 hours of exchange, pure silicaRm aluminas and CoMo/SiO 2 exchanged virtually no more deuterium. In the case of the mechanical mixtures, the surface Figure 3. OD bands surface (m) and Br6nsted of the OD bands was still growing after sites band surface (o) as a function of sample that period, but at a much lower rate. In composition, compared to the theoretical values spite of this slow growth, the values calculated as in Equation 3 (dashed lines). obtained at that point were chosen to compare quantitatively the amounts of deuterium exchanged. The reproducibility of the experiments has been checked; the variation between different wafers of the same sample was always inferior to 5%. Figure 3 shows the total surface of the OD bands, after 20 h of exchange, plotted against the composition of the mechanical mixtures. The dashed lines represent the sum of the individual contributions of

104 the isolated phases, calculated as in Equation 3. Each series presents a maximum of exchange, located approximately at the same compositions as those of DPM cracking activity. An important synergetic effect is observed. In all cases, the experimental results are considerably higher (increase by 200% to 1000%) than the theoretical values.

4. DISCUSSION XPS results show that no new peaks are formed or significant energy shifts occur, which indicates that the identity of the two kinds of particles contained in the catalyst are preserved. A significant sintering or redispersion of Co or Mo during the preparation and/or sulfidation of the mechanical mixtures can also be excluded, since the surface compositions correspond to the amounts expected due to the simple "dilution" of the CoMo/SiO 2 phase with the pure silica-alumina. These mechanical mixtures are therefore suitable to study the possible effect of spillover hydrogen. Any change observed in the catalytic behavior or physico-chemical properties of the surface when comparing the mixtures with the individual components must therefore be attributed to some interaction between the unmodified phases. The products obtained from DPM cracking in the present work agree with the results from the literature, mentioned in the Introduction, which indicate that the reaction proceeds via carbocation formation on acidic sites. This implies that the decomposition of DPM does not need the successive intervention of two catalytic sites, like in the "ideal hydrocracking" mechanism. Only acidic sites are sufficient to carry out the reaction. The improved activity of the mixtures when compared to the pure phases must therefore be explained differently. The ability of the unsupported or supported transition metal sulfides to adsorb and dissociate molecular hydrogen from the gas phase is well known [16-19]. It is therefore natural to suppose that the mobile hydrogen species thus formed (Ho, formed by the homolytic scission of H 2, or H + and H-, formed by heterolytic scission) could migrate (spillover) from the sulfides to the surface of the silica-alumina, where they could create new active sites and also contribute directly to the cracking of DPM. The first step of that reaction would be similar to that of a dealkylation [20]. Initially, DPM is adsorbed on a BrtJnsted acidic site and protonated on an aromatic carbon to give the corresponding carbonium ion. The latter can crack to give benzene in the gas phase plus a benzyl carbonium ion adsorbed on the surface. This cation can then react with two hydrogen atoms or with a hydride and desorb, giving toluene in the gas phase (R § + H- ---) RH or R § + 2H ---) RH + H§ A proton regenerates the Brtinsted site. By-products with higher molecular weight can be produced by transalkylation reactions [20]. This mechanism could explain the synergy observed between the two phases. This interpretation of the experimental data is supported by the differences observed in the deactivation patterns and carbon contents after test, since one notorious effect of Hsp is the capacity to diminish the deactivation caused by coke deposition on the active sites [21,22]. This is supposed to be due to a reaction with the coke precursors, very likely a hydrogenolysis. In pure silica-aluminas, where no source of spillover is present, no special protection against deactivation should be observed. Indeed, the silica-aluminas lose most of their activity (about 80%) before reaching the steady-state and present the highest carbon contents after catalytic test. On the other hand, in the case of the mechanical mixtures, where spillover hydrogen is continuously produced by the CoMo/SiO 2 phase and can migrate to the silica-alumina surface, the predicted protection effect is noticed. The relative losses of activity are much lower

105

(between 35% and 50%). Carbon contents are also lower than the simple addition in adequate proportions of the individual contributions of the components of the mixtures. In pure CoMo/SiO 2, the source of Hsp is in the immediate vicinity of the few weak active sites that SiO 2 contains. The distance that activated hydrogen should migrate to reach the coke precursors is very small; in consequence, the protection offered by this mobile species can be most effective. We observe accordingly that this sample presents the lowest relative loss of activity and the lowest carbon content. The results obtained by FTIR measurements of adsorbed pyridine also support the explanation proposed to the observed synergy effect. They indicate an enhancement of the amount of Brtinsted acidic sites of the mechanical mixtures, when compared to the sum of the individual contributions of the pure phases. In the SA6 series the effect cannot be shown unequivocally, but it was clearly detected in SA12 and SA60 series. The action of spillover hydrogen is the most likely explanation to the creation of these acidic sites. In this context, it is striking that the compositions where the amount of these sites was maximum correspond to those where the catalytic activity is the highest (R m = 50 in SA12 and R m = 75 in SA60). The final proof is brought by the isotopic exchange experiments. The interpretation of deuterium exchange results needs some care, since species like HDO or D20 may be responsible for a non-spillover exchange mechanism [23]. The authors proposing this mechanism also indicate that these species could be formed by the reaction between oxygen traces and deuterium, in the presence of metals. One possibility, in the present work, is therefore that HDO and/or D20, that could be present in the deuterium as impurities, would be responsible for the observed exchange. This is not likely, because the gas had been extremely carefully purified. Besides, no IR bands corresponding to these hypothetical oxygenated species (1445 and 1218 cm -1 for HDO and D20, respectively [15]) have been detected on the pure phases or mechanical mixtures. The explanation based on the role of HDO or D20 should be rejected on the ground of experimental evidences. If we accepted this reasoning, we should have observed a significant exchange in the pure silica-aluminas, where there was no Hsp source. This was not the case. Practically no deuterium was detected on their surfaces. We must therefore conclude that there is no or only a negligible direct exchange through the gas phase. Consequently, we must conclude that our results represent a direct evidence for the existence of spillover phenomena in our mixtures. D 2 is adsorbed and dissociated on the metal sulfides. The species created in this way can then migrate (spillover) to the surface of the silica (primary spillover) and of the silica-alumina (secondary spillover), making the exchange with the hydroxyls of the latter possible. The shape of the curves in Figure 3, representing the amount of deuterium exchanged after 20 h, is similar to those of the catalytic activity at the steady-state in Figure 2. This strongly suggests that the magnitude of the synergy in catalytic activity is related to the generation of spillover hydrogen. The maxima found in both cases (catalytic tests and isotopic exchange) correspond to a situation where there is an ideal balance between the activity of the donor (supported metal sulfides) to produce Hsp and the capability of the acceptor (silica-alumina) to exchange with Hsp. In the range of R m values below that optimum composition, the growth of the amount of deuterium exchanged (or of catalytic activity) reflects the increase in the number of active sites that can be generated, due to the increasing silica-alumina content. When this content is superior to that optimum value, the donor is not able any more to produce enough Hsp to fully activate the acidic phase. This explains the decrease in the number of deuteroxyls (and in catalytic activity) observed beyond that ideal composition.

106 5. CONCLUSIONS A very important synergy effect in the cracking of diphenylmethane is observed when particles of a sulfided CoMo/SiO 2 catalyst are mixed with particles of an amorphous silicaalumina. This can be attributed to the action of spillover hydrogen, which can create new active sites and partially protect them from deactivation by coke deposition. Isotopic exchange experiments, showing that the presence of CoMo/SiO 2 brings about the exchange of deuterium with the hydroxyls of the silica-alumina, constitute a direct evidence of this phenomenon.

REFERENCES

1. J.A. Martens, P.A. Jacobs and J. Weitkamp, Appl. Catal., 20 (1986) 283. 2. M. Lacroix, G.M. Pajonk and S.J. Teichner, Bull. Soc. Chim. Fr., 7-8 (1981) 265. 3. H. Hattori, in "New Aspects of Spillover Effect in Catalysis" (T.Inui et al., eds.), p. 69, Elsevier, Amsterdam, 1993. 4. F. Roessner, U. Roland and T. Braunschweig, J. Chem. Soc. Far. Trans., 91 (1995) 1539. 5. I. Nakamura, K. Aimoto and K. Fujimoto, AIChE Symp. Ser., 85 (1989) 15. 6. I. Nakamura, R. Iwamoto and A. I-ino, in "New Aspects of Spillover Effect in Catalysis", (T.Inui et al., eds.), p. 77, Elsevier, Amsterdam, 1993. 7. H. Hattori, K. Yamashita, T. Tanabe and K. Tanabe, Proc. 9th Int. Congr. Catal., p. 27. 8. H. Shimada, M. Kurita, T. Sato, Y. Yoshimura, T. Hirata, T. Konakahara, K. Sato and A. Nishijima, Chem. Lett., (1984) 1861. 9. H. Shimada, T. Sato, Y. Yoshimura, J.Hiraishi and A.Nishijima, J. Catal., 110 (1988) 275. 10. W.C. Conner, Jr., G.M. Pajonk and S.J. Teichner, Adv. Catal., 34 (1986) 1. 11. W.C. Conner, J.F. Cevallos-Candau, N. Shah and V. Haensel, in "Spillover of Adsorbed Species" (G.M. Pajonk et al., eds.), p. 31, Elsevier, Amsterdam, 1983. 12. U. Roland, R. Salzer and S. Stelle, in "Zeolites and Related Microporous Materials: State of the Art 1994" (J. Weitkamp et al. ,eds.), p. 1231, Elsevier, Amsterdam, 1994. 13. L.T. Weng, G. Vereecke, M.J. Genet, P. Bertrand and W.E.E. Stone, Surf. Interf. Anal., 20 (1993) 179. 14. A.M. Stumbo, P. Grange and B. Delmon, to be published. 15. E. Baumgarten and E. Denecke, J. Catal., 95 (1985) 296. 16. M. Karroua, H. Matralis, P. Grange and B. Delmon, J. Catal., 139 (1993) 371. 17. S. Giraldo, P. Grange and B.Delmon, in "New Aspects of Spillover Effect in Catalysis" (T. Inui et al., eds.), p. 345, Elsevier, Amsterdam, 1993. 18. X. Chu and L.D. Schmidt, J. Catal., 144 (1993) 77. 19. N.M. Rodriguez and R.T.K. Baker, J. Catal. 140 (1993) 287. 20. P.A. Jacobs, in "Carboniogenic Activity of Zeolites", Elsevier, Amsterdam, 1977, p. 113-120. 21. G.M. Pajonk, in "2nd Conference on Spillover" (K.-H. Steinberg, ed.), p. 1, Leipzig, 1989. 22.J. Kapicka, N.I. Jaeger and G. Schulz-Ekloff, Appl. Catal., 84 (1992) 47. 23. D. Bianchi, D. Maret, G.M. Pajonk and S.J. Teichner, in "Spillover of Adsorbed Species" (G.M. Pajonk et al., eds.), p. 45, Elsevier, Amsterdam, 1983.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) l l t h I n t e r n a t i o n a l C o n g r e s s on Catalysis - 40th Anniversar 3,

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

107

H y d r o d e s u l f u r i z a t i o n of B e n z o t h i o p h e n e Catalyzed by M o l y b d e n u m Sulfide Cluster E n c a p s u l a t e d into Zeolites M. Taniguchi, a S. Yasuda, ' Y. Ishii, b T. Murata, b M. Hidai b and T. Tatsumi ~ aEngineering Research Institute, Faculty of Engineering, The University of Tokyo, Yayoi, Tokyo 113, Japan bDepartment of Chemistry and Biotechnology, Faculty of Engineering, The University of Tokyo, Hongo, Tokyo 113, Japan

A cationic molybdenum sulfide cluster [Mo354(H20)9] 4§ with incomplete cubane-type structure and a cationic nickel-molybdenum mixed sulfide cluster [Mo3NiS4CI(H20)9]3§ with complete cubane-type structure were introduced into zeolites NaY, HUSY and KL by ion exchange. Stoichiometry of the ion exchange was well established by elemental analyses. The UV-visible spectra and EXAFS analysis data exhibited that the structure of the molybdenum cluster remained virtually intact after ion exchange. MoNi/NaY catalyst prepared using the molybdenum-nickel sulfide cluster was found to be active and selective for benzothiophene hydrodesulfurization.

1. INTRODUCTION The great practical significance of the processes of hydrodesulfurization (HDS) of petroleum feedstocks has developed a keen research interest in the structure and synthesis method of the active component of HDS catalysts. All HDS processes have been performed on catalysts of the same type, based on molybdenum (or tungsten) sulfide supported on high surface area 7-aluminas with cobalt or nickel added as a promoter. In recent years much attention has been focused on the deep HDS of light oil, since the reduction in sulfur concentration in fight oil is effective in reducing particulate and NOx in the exhaust gas of Diesel engines. The main problem in the deep HDS of fight oil is the conversion of alkyldibenzothiophenes, e.g., 4,6-dimethyldibenzothiophene, which is one of the most difficult compounds to desulfurize among dibenzothiophene derivatives [1]. One approach to this problem is to rearrange and/or to eliminate alkyl groups by using catalysts with enhanced acidity. Zeolite-supported molybdenum with nickel or cobalt as a promoter is promising in this regard. Molybdenum/zeolite catalysts prepared by impregnating zeolites with ammonium heptamolybdate solution generally give rise to poor dispersion of molybdenum [2]. In contrast, ion exchange would be an ideal method for loading active metal species onto supports. Few cationic forms are available as simple salts of molybdenum of high oxidation

108

state, however. Furthermore, most of them exist only in strongly acidic solutions where many zeolites are unstable and where exchangeable cations must compete with protons. Therefore few studies have succeeded in introducing cationic molybdenum compounds into zeolites by ion exchange [3-5]. Molybdenum/Y-zeolites were prepared by aqueous ion exchange with Mo2(ethylenediamine)4Ch [3] and MoO2CI: [4]. Lunsford et al. also reported solid-state exchange of HY and ultrastable form of HY (HUSY) with MoC15 [5]. A molybdenum sulfide cluster, [Mo3S4(H20)9] 4§ (1) with incomplete cubane-type structure was recently prepared [6-8]. It is stable in water and air and easily forms mixed metal clusters with cubane-type Mo3MS4 cores (M = Fe, Co, Ni, Pd, etc.) [9]. Mixed metal clusters with a Mo3PdS4 core have been derived from 1, showing intriguing reactivities at the Pd site toward alkenes, CO, isonitriles, and alkynes [10,11]. A molybdenum-nickel cluster with a Mo3NiS4 core has also been found to uptake CO stoichiometrically to give a new cluster where one CO molecule combines with the nickel atom [12]. Thus these clusters are of considerable interest in connection with potential application to a variety of catalytic reactions. It occurred to us that the molybdenum and mixed metal clusters could be incorporated into zeolites by ion exchange and promote unique catalytic reactions. Preparation of molybdenum/zeolites using the molybdenum sulfide cluster cation as a precursor by ion exchange is considered to have several advantages. First, molybdenum species could be loaded with high dispersion. Secondly, molybdenum is loaded on zeolites as sulfide, not as oxide; presulfiding is unnecessary before HDS reactions. To our best knowledge, there is no such report that succeeded in introducing molybdenum species into zeolites in its sulfide state and applying the resulting catalyst to HDS reactions. Thirdly, by loading molybdenum-nickel cluster with a Mo3NiS, core into zeolites, molybdenum and nickel could be introduced in the homogeneously well-mixed state. This paper describes the successful incorporation of molybdenum and molybdenum-nickel clusters into zeolites with 12-membered ring by aqueous ion exchange and application of the resulting materials to HDS reaction of benzothiophene. Stoichiometry of the ion exchange was examined by elemental analysis. UV-visible spectroscopy and EXAFS measurements were carried out to investigate the structure of molybdenum species loaded on zeolites.

2. EXPERIMENTAL

2.1. Catalyst preparation The chloride salt of cluster 1 was synthesized by the reported method [8]. As supports, NaY (Nikka Seiko, SK-40; Si/AI = 2.3), HUSY (Tosoh, HSZ-330HUA; Si/A1 = 3.1) and KL (Tosoh, TSZ-500KOA; Si/AI = 3.1) zeolites with 12-membered ring were used. To a suspension of zeolite (4.58 g) in water (91.6 g) was added dropwise a 0.01 M aqueous solution of the chloride salt of 1 (86.9 ml) with vigorous stirring at 313 K. In the case of NaY the color of zeolite changed from white to brown and the solution was colorless. For HUSY and KL, brown-colored species remained in the solution after ion exchange. The metal cluster-containing zeolite was separated by filtration and washed with distilled water, followed by drying at 313 K under reduced pressure. Resulting materials are referred to as Mo/NaY, Mo/HUSY and Mo/KL, respectively. The chloride salt of cluster [Mo3NiS4CI(H20)9]3+ (2) was also synthesized according to the reported method [9] and

109 similarly treated with NaY to afford MoNi/NaY. MoNi/A1203 was also prepared by impregnating A1203 (JRC-ALO-4) with the solution of chloride salt of 2. 2.2. Characterization Molybdenum and sodium concentrations in the filtrate were determined by ICP on a Nippon Jarrell-Ash ICAP-575 spectrophotometer in order to estimate molybdenum loadings after ion exchange and to confirm the stoichiometry of ion exchange. Sulfur contents were determined by the LECO method. Chlorine contents of the catalysts were determined by ion chromatography. Powder X-ray diffraction patterns of the catalysts were collected on a Rigaku RINT 2400 X-ray diffractometer. Crystallinity of the zeolites were determined by the average of the relative intensities of selected seven intense peaks, using starting material NaY, HUSY and KL as the reference. UV-visible spectra were recorded on a Hitachi U-4000 spectrophotometer. Mo K-edge X-ray absorption spectra were collected on a R.igaku R-EXAFS 2100S (30 kV, 280 mA) instrument using a Ge(400) crystal monochromator at room temperature in air atmosphere. Four samples were applied to the EXAFS measurement: (a) 5% physical mixture of cluster 1 and NaY, (b) Mo/NaY, (c) sample (b) treated in flowing helium at 373 K for 0.5 h and (d) sample (b) treated in flowing helium at 573 K for 0.5 h. The sample (a) was set on a sample holder with 0.3 mm thickness, and the samples (b)-(d) were pressed into self-supporting wafers. The method of EXAFS data analysis utilized REX, an EXAFS data analysis software provided by Rigaku. The sample (a) was used as a standard for Mo-S, Mo-O and Mo-Mo interactions. Crystallographic data (N and R of Mo-S, Mo-O and Mo-Mo, where N is the coordination number and R the radial distance from the absorber to the backscatter atom) of the cluster 1 [13] were used in order to determine the EXAFS parameters o and AE0 (where o is the Debye-Waller factor and AEo the inner potential correction of the edge position) which were used for the samples (b)-(d). Once a close fit was obtained for the parameters N and R with k3-weighting, the fit was optimized by allowing a few of the parameters to be fitted while the rest of the parameters were held fixed. 2.3. HDS reaction The catalytic activities of the zeolite- and alumina-supported Mo and Mo-Ni catalysts for benzothiophene HDS were measured in a flow system incorporating a microreactor made of stainless steel tube containing 0.2 g of catalyst. The catalysts were pretreated in flowing helium by ramping the temperature from room temperature to 573 K at a rate of 3.3 K/min and then in flowing hydrogen for 0.5 h at 573 K. When presulfided, the catalysts were treated in flowing 5% H2S/H2 mixed gas for 2 h at 573 K following hydrogen pretreatment. Benzothiophene in decane (0.5% or 5% as sulfur) was supplied to the reactor at a rate of 4.5 ml/h. The reaction conditions were 573 K, 3.0 MPa, H2/benzothiophene = 379 or 3790 (mol/mol) and W/F = 38.2 or 382 g-cat.h/mol-benzothiophene.

110 3. RESULTS AND DISCUSSION

3.1. Chemical composition The results of elemental analysis of the molybdenum and molybdneum-nickel/zeolite catalysts are shown in Table 1. After ion exchange of NaY with the solution of 1, almost all molybdenum was found to be transferred from the solution to NaY; the resulting Mo/NaY contained 5.0 wt% of molybdenum, which corresponds to the presence of 0.41 Mo3S4 cores per a supercage of the FAU type structure. The size of cluster 1 was estimated to be 0.72 nm, smaller than the diameter of the aperture (0.74 nm) [14] of supercages of the FAU type structure. The C1/Mo ratio of Mo/NaY was 0.11, suggesting that most of the cluster 1 acted as a tetravalent cation in the ion exchange. Fled sodium ion per loaded Mo cluster in the filtrate was 4.6 times (ideally 4 times) as much as the loaded cluster 1. The slight excess of lost sodium over the supported cluster may be due to the acidity of the cluster salt. The S/Mo ratio was maintained after ion exchange. In the preparation of Mo/HUSY, the cluster 1 amounting to 2.5 wt% (as molybdenum metal) of HUSY was added to the suspension of HUSY; 92% of the molybdenum was loaded onto HUSY. The CI/Mo ratio of Mo/HUSY was found to be 0.34, suggesting that in ion exchange the cluster 1 acted as a trivalent cation on the average. These findings indicate that the protons in HUSY are less exchangeable by the cluster cation than the Na cations in NaY. In the preparation of Mo/KL, the addition of 1 was stopped when pH of the solution was lowered to about 3. The resulting Mo/KL contained only 2.1 wt% (74% of added molybdenum clusters) of molybdenum. Chlorine was absent, which indicates that the cluster 1 acted as a tetravalent cation. The LTL structure is characterized by a monodimensional system of channels, whose diameter (0.70 nm) [14] is close to the size of the cluster 1. It is conceivable that, once the cluster 1 was incorporated into an LTL main channel and present at a site near the external surface, the sites deep in the channel are no more accessible to another cluster. The cluster 2 was also introduced into NaY by ion exchange. After the ion exchange, 99% of molybdenum and nickel were loaded on NaY; the Ni/Mo ratio did not change. The CI/Mo ratio suggests that excess chlorine was present on the catalyst.

Table 1 Elemental analysis of Mo and Mo-Ni / zeolite catalysts Catalyst

Si/A1 ratio

Loaded % loaded % loaded Na CI/Mo Mo (wt%) Mo a Ni" released b ratio

Mo / NaY

2.3

5.0

Mo / HUSY

3.1

2.3

Mo / KL

3.1

2.1

MoNi / NaY

2.3

5.0

99

100

S/Mo ratio

-

4.6

0.11

92

-

-

0.34

n.d.c

74

-

-

0.00

n.d.C

4.2

0.96

n.d.c

99

a Fraction of Mo or Ni loaded based on Mo or Ni added to the solution. b Mol Na fled into the solution per mol cluster loaded. c Not determined.

1.3

111

Table 2 Crystallinity of Mo and Mo-Ni / zeolite catalysts Catalyst

Si/A1 ratio

pH after ion exchange

Crystallinity (%)

Mo / NaY Mo / HUSY Mo / KL

2.3 3.1 3.1

4.6 4.2 3.2

75 85 93

MoNi / NaY

2.3

4.7

72

3.2. Crystallinity of zeolites It is reported that introduction of molybdenum cations into zeolites often gives rise to destruction of the zeolite skeleton because of their acidity [5,15]. The crystaUinity of the ion exchanged zeolites are shown in Table 2. Although the solutions during ion exchange were slightly acidic (pH = 3.2 - 4.7), the X-ray powder diffraction patterns were only slightly attenuated, indicating the zeolites kept their crystallinity without significant destruction. The decrease in crystallinity was not related to the pH after ion exchange but to the Si/AI ratio of the zeolites.

(a)

O

< (

I

200

300

I

1

1

I

400 500 600 700 Wavelength (nm)

1

800

900

Figure 1. UV-visible spectra of (a) physical mixture of NaY and the chloride salt of 1, (b) Mo/NaY, (c) Mo/HUSY and (d) Mo/KL.

112 3.3. Structure of the clusters UV-visible spectra of the cluster 1 and molybdenum/zeolite catalysts are shown in Figure 1. The cluster 1 showed bands at 300, 390 and ca. 650 nm. Similar bands were observed for the spectrum of each molybdenum/zeolite catalyst, suggesting that the structure of cluster 1 was practically unchanged after ion exchange. In order to obtain more structural information about the molybdenum species in Mo/NaY, EXAFS measurements of the cluster 1 and Mo/NaY were carried out. The Fourier transforms of the EXAFS data are shown in Figure 2. Structural parameters (Table 3) showed no change of the Mo-O, Mo-S and Mo-Mo distances, suggesting that there is no significant structural difference between the cluster 1 and the molybdenum compound in the Mo/NaY. From these EXAFS parameters and the UV-visible spectra, it is considered the structure of cluster 1 remained virtually intact after ion exchange.

14

14 (a)

12

(b)

12 .~10

~10

...a

'~

8

6 a~ 4

u2 4

0

0.1

0.2

0.3 0.4 R (nm)

0.5

0

0.6

0.1

0.2

0.3 0.4 R (nm)

0.5

0.6

14

14

(c)

12

(d)

12

~10

.~10

'2

'~ e~

8

8

6

6

a~ 4

4

0

0 0

0.1

0.2

0.3 R (nm)

0.4

0.5

0.6

0

0. I

0.2

0.3

0.4

0.5

0.6

R (nm)

Figure 2. Fourier transforms of EXAFS data of (a) the chloride salt of 1, (b) Mo/NaY, (c) sample (b) treated in flowing He at 373 K and (d) sample (b) treated in flowing He at 573 K.

113

Table 3 Structural parameters from EXAFS of the Mo/NaY catalysts Mo-S Sample Cluster I

Treatment -

N~

Mo-Mo R (nm) b

(3.0) ~ (0.230) ~

Na

Mo-O R (nm) b

(2.0) c (0.274) c

Na

R (nm) b

(3.0) ~ (0.218) ~

3.5

0.232

2.0

0.276

3.9

0.217

Mo / NaY

He, 373 K

3.1

0.232

1.7

0.277

3.3

0.218

Mo / NaY

He, 573 K

1.2

0.238

0.69

0.276

1.3

0.214

Mo / NaY

-

"Coordination number. b Radial distance from the absorber to the backscatter atom. ~ Taken from ref. 13 and treated as fixed parameters.

After treatment at 373 K in helium flow, the coordination numbers of Mo-O, Mo-S and Mo-Mo were not largely different from those in the cluster 1 and the distances of Mo-O, Mo-S and Mo-Mo were hardly changed. After treatment at 573 K in He flow, however, the color of NaY changed from brown to black and the curve fitting results of the EXAFS data exhibited lower coordination numbers of all interactions than those of the cluster 1. The decrease in the coordination number seems to be not due to the decrease in sulfur amount in the catalyst but to the disordering of each interaction since the S/Mo ratio hardly decreased after thermal treatment. These results show that the structure of the cluster 1 loaded on NaY was maintained at 373 K, but lost at 573 K.

3.4. HDS reactions The results of hydrodesulfurization of benzothiophene (0.5 wt% as sulfur) in decane are shown in Table 4. Over Mo/NaY and Mo/HUSY, selectivity for ethylbenzene was low and main products were dihydrobenzothiophene and alkylbenzothiophenes. The latter seems to be produced by the reaction of unreacted benzothiophene with alkenes generated from decane cracking on acid sites of the zeolite. For Mo/I-/USY, decane cracking and following alkylation of benzothiophene took place appreciably. Mo/KL showed high HDS activity compared to Mo/NaY in spite of its lower molybdenum loading. This is consistent with our supposition that molybdenum clusters were loaded on near external surface of KL, as described above. Since the KL zeolite has no strong acidity, decane cracking and following alkylation of benzothiophene proceeded only to a small extent. Involvement of Ni resulted in great enhancement of HDS activity; MoNi/NaY proved to be a highly active and selective catalyst for HDS of benzothiophene to ethylbenzene, and other products were scarcely observed. No change in the activity was observed during 4 hours. Table 5 shows HDS product distributions over several catalysts prepared by using the molybdenum-nickel cluster 2. Sulfur content in decane was adjusted to 5.0 wt% in these experiments. MoNi/NaY was found to be more active than MoNi/Al203. It is to be noted that during the high temperature pretreatment the original cluster structure would have been changed. However, the high activity of the MoNi/NaY catalyst for benzothiophene HDS is probably due to the formation of active sites derived from this particular mixed metal cluster,

114 Table 4 Effect of nickel and acidity of zeolites on the benzothiophene hydrodesulfurization ~ Selectivity (%) Catalyst

Conversion (%)

EB b

DHB"P

Others

Mo/NaY

53

8.3

76

16

Mo / HUSY Mo/KL

39 75

6.1 38

66 52

28 9.4

MoNi / NaY

97

88

1.2

5.7

a Reaction conditions: 573 K, 3.0 MPa, W/F = 382 g-cat.h/mol-benzothiophene, S content 0.5 wt%. b Ethylbenzene. Dihydrobenzothiophene.

Table 5 Catalytic performance of MoNi catalysts for benzothiophene hydrodesulfurization ~ Selectivity (%) Catalyst

Conversion (%)

EB b

DHBT:

Others

MoNi / NaY MoNi / NaY (presulfided)

36 19

51 23

39 55

11 4.7

MoNi / A1203 MoNi / A1203 (presulfided)

23 30

42 39

52 45

5.1 16

a Reaction conditions: 573 K, 3.0 MPa, W/F = 38.2 g-cat.h/mol-benzothiophene, S content 5.0 wt%. b Ethylbenzene. c Dihydrobenzothiophene.

together with high dispersion. The pore structure of the zeolite support may also play a role in the stabilization of active species reminiscent of the cluster structure. Presulfiding effect was quite different between two supports; although MoNi/AI203 was activated by presulfiding, MoNi/NaY was slightly deactivated by presulfiding, suggesting that the active species in MoNi/NaY is different from that in MoNi/A1203. The S/Mo ratio of the cluster 1 is 1.3. It is suspected that the increase in the HDS activity of MoNi/AI203 by sulfiding resulted from the increase in the sulfur content, in agreement with the generally accepted belief that the active component of A1203-supported conventional HDS catalysts is MoS2 [16], where the S/Mo ratio is 2. Study on applicability of MoNi/zeolite catalysts to the HDS reactions of other sulfur compounds is ongoing in our laboratory.

115 4. CONCLUSIONS A cationic molybdenum sulfide cluster 1 with incomplete cubane-type structure and a cationic molybdenum-nickel sulfide cluster 2 with complete cubane-type structure were employed as new precursors for Mo/zeolite and MoNi/zeolite catalysts. Mo/NaY, Mo/HUSY, Mo/KL and MoNi/NaY were prepared by aqueous ion exchange with the cluster 1 and 2. Stoichiometry of the ion exchange was well established by elemental analysis. Though pH of the solution during ion exchange was lowered, crystallinities of the zeolites were preserved. The UV-visible spectra and EXAFS analysis data exhibited that the structure of the cluster 1 remained virtually intact after ion exchange. In the HDS of benzothiophene Mo/HUSY and Mo/KL showed higher activity than Mo/NaY. The use of molybdenum-nickel mixed cluster resulted in a great enhancement of HDS activity. By presulfiding, MoNi/AlzO3 was activated, whereas MoNi/NaY was not. These results suggest that the active species in MoNi/NaY is different from that in MoNi/AI203.

ACKNOWLEDGMENT We thank Professor K. Asakura, Faculty of Science, The University of Tokyo, for his support with the EXAFS analyses.

REFERENCES

1. A. Ishihara and T. Kabe, Ind. Eng. Chem. Res., 32 (1993) 753. 2. e.g., M. Laniecki and W. Zmierczak, Zeolites, 11 (1991) 18. 3. M.B. Ward, K. Mizuno and J.H. Lunsford, J. Mol. Catal., 27 (1984) 1. 4. M. Huang and R. F. Howe, J. Catal., 108 (1987) 283. 5. P.E. Dai and J.H. Lunsford, J. Catal., 64 (1980) 173. 6. F.A. Cotton, Z. Dori, R. Llusar and W. Schwotzer, J. Am. Chem. Soc., 107 (1985) 6734. 7. M. Martinez, B.L. Ooi and A.G. Sykes, J. Am. Chem. Soc., 109 (1987) 4615. 8. T. Shibahara, M. Yamasaki, G. Sakane, K. Minami, T. Yabuki and A. Ichimura, Inorg. Chem., 31 (1992) 640. 9. P.W. Dimmock, G.J. Lamprecht and A.G. Sykes, J. Chem. Soc., Dalton Trans., (1991) 955 and references therein. 10. T. Murata, H. Gao, Y. Mizobe, F. Nakano, S. Motomura, T. Tanase, S. Yano and M. Hidai, J. Am. Chem. Soc., 114 (1992) 8287. 11. T. Murata, Y. Mizobe, H. Gao, Y. Ishii, T. Wakabayashi, F. Nakano, T. Tanase, S. Yano, M. Hidai, I. Echizen, H. Nanikawa and S. Motomura, J. Am. Chem. Soc., 116 (1994) 3389. 12. T. Shibahara, S. Mochida and G. Sakane, Chem. Lett., (1993) 89. 13. H. Akashi, T. Shibahara and H. Kuroya, Polyhedron, 9 (1990) 1671. 14. W.M. Meier and D.H. Olson, Atlas of Zeolite Structure Types (Third edition), Butterworth-Heinemann, Stoneham, 1992.

116 15. e.g., R. Cid, F.J.G. Llambias, J.L.G. Fierro, A.L. Aguro and J. Villasenor, J. Catal., 89 (1984) 478. 16. A.N. Startsev, Catal. Rev.-Sci. Eng., 37 (1995) 353.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

117

Role o f a d s o r b e d h y d r o g e n species on r u t h e n i u m and m o l y b d e n u m sulfides. C h a r a c t e r i z a t i o n b y inelastic n e u t r o n scattering, t h e r m o a n a l y s i s m e t h o d s and m o d e l reactions. M. Lacroix a, H. Jobic a, C. Dumonteil a, P. Afanasiev a, M. Breysse a and S. Kasztelanb. aInstitut de Recherches sur la Catalyse, 2 avenue Albert Einstein, 69626 Villeurbanne cedex France bInstitut Frangais de P6trole, l&4 avenue de Bois Pr6au, BP 311, 92506 Rueil-Malmaison, France.

Abstract

The interaction of hydrogen over unsupported MoS 2 and RuS 2 has been investigated as a function of the sulfur to metal ratio. On these solids the presence of sulfur deficient sites is required to generate an activity and to allow hydrogen chemisorption. The nature of the adsorbed species differs depending on the catalyst under investigation. On RuS2, two types of hydrogen were evidenced by thermoflash desorption and inelastic neutron scattering: one was assigned to hydrogen adsorbed on surface sulfur anions while the other one is retained on coordinatively unsaturated ruthenium cations. By contrast, only SH groups were detected on MoS2. ESR measurements have shown that a fraction of chemisorbed hydrogen induces a modification of the concentration of paramagnetic Mo(V) and Mo(III) species. Thus both solids behave differently towards an hydrogen atmosphere. RuS2 has a pseudometallic comportment whereas for MoS 2 redox or acid base properties are involved.

1. INTRODUCTION Sulfide catalysts find extensive use as hydrotreating catalysts. Hydrotreating is a general term that includes hydrogenation and hydrogenolysis reactions i.e. hydrodesulfurisation, and hydrodenitrogenation. As these reactions require hydrogen, the interaction of hydrogen with the solid is important in order to understand their catalytic properties. Current industrial catalysts are supported systems containing basically a group VI sulfide like MoS 2 or WS 2 promoted by cobalt or nickel sulfides. The properties of these multicomponent solids are complex because of the observed synergy between groups VI and VIII sulfides and of the interaction with the support which could affect the hydrogen reactivity. From a basic point of view it appears sensible to simplify the industrial catalysts by reducing them into their component parts.

118 Polycrystalline molybdenum sulfide has been frequently used as model catalyst. This solid crystallizes in the hexagonal system in which the molybdenum ions are located at the center of a trigonal prismatic unit formed by six surrounding S2" anions. Assembling these moieties leads to a lamellar structure with the basal planes presenting a continuous layer of sulfur. These [hk0] planes were found poorly reactive towards hydrogen while the edges of the crystal may easily undergo partial desulfurisation generating sulfur vacancies which are seen as the basis of the catalytic and adsorbing properties [1, 2]. The interaction of hydrogen with this anisotropic surface has been the subject of many investigations. The review of Moyes and the paper of Komatsu and Hall [3, 4] have shown that the amount of adsorbed hydrogen changes drastically from one study to another. For instance, using the same preparation method (decomposition (NH4)2MoS4) the values of H/Mo vary between 0.012 and 0.37. This important variation is probably due to various catalysts pretreatment prior to adsorption measurements leading to different number of vacancies. This hypothesis was supported by Jalowiecki et al [5, 6] who reported that a fully sulfur saturated slab was inactive and did not chemisorb hydrogen while both the activity and the hydrogen adsorption capacity increased and went through a maximum as far as sulfur was removed from the catalyst upon reduction. This showed that unsaturation is a necessary prerequisite for giving rise to a reactive surface. Moreover, the observed results suggested that both properties required an optimal concentration of vacancies and sulFar species. Therefore, it has been envisaged that hydrogen is heterolytically adsorbed on a Mo-S pair leading to the formation of Mo-H and SH groups [5-7]. Quantum chemical calculations have also suggested that heterolytic adsorption is the favored route for hydrogen adsorption on MoS2 [8]. From this theoretical study it was also found that sorbed hydrogen is able to move from the edge sites to the basal plane and the resulting vacant site could be replenished by further heterolytic adsorption. However, characterization of the adsorbed species by inelastic neutron scattering (INS) and 1H NMR have detected the presence of SH groups while the signal ascribed to Mo-H was never observed even for a hydrogen saturated slab [4, 9, 10]. Recently, the interaction of hydrogen with a model RuS2 catalyst was studied using thermoanalysis methods, IH NMR and INS [ 1113]. The systematic applications of these techniques for the characterization of the adsorbed species present on different reduced states proved the existence of both Ru-H and SH entities whose relative population drastically depended on the sulfur to metal ratio. Coupling these results with catalytic activity measurements, it was shown that the H2-D2 exchange and the 1butene hydrogenation rates were directly related to the concentration of Ru-H species suggesting that a hydride type hydrogen is the active species involved in hydrogenation. Taking into account these experimental data, the following questions could be addressed: i) Is the presence of hydride type hydrogen essential for explaining the activity of a given sulfide catalyst .9, ii) Why was the Mo-H bond never detected in MoS2? In order to get insight into these fundamental questions, a similar methodology as previously utilized for RuS2 was undertaken and results obtained on both systems were systematically compared.

2. EXPERIMENTAL Unsupported ruthenium sulfide was prepared by precipitation at room temperature from an aqueous solution of RuCI3 by pure H2S and by further sulfidation under an H2S flow at

119 673K for 2h. The solid was then cooled to room temperature under the same sulfiding atmosphere, flushed with an oxygen free inert gas and stored in sealed bottles. Physicochemical characterizations confirmed that the obtained solid has the expected pyrite structure and elemental analysis indicated a S/Ru ratio equal to 2.25. Microcrystalline molybdenum sulfide was prepared by thermal decomposition of 0NIHa)2MoS4 (ATTM). The later was synthesized using the method described by Dieman et al [14]. UV-Vis spectra of aqueous solutions of ATTM have shown that this precursor does not contain any absorption bands corresponding to oxothiomolybdates. MoSx was obtained by heating the ATTM crystals at 673K for 4h in flowing mixtures of 15% HzS in H2. The x value determined gravimetrically on a microbalance (Sartorius 4433) by oxidizing the sulfided sample in 02 at 773K to form MoO 3 was found to be close to 2.27. The desulfurization of the catalysts was carried out in a dynamic microreactor at 1 atmosphere hydrogen pressure and at different temperatures. The setup included a specific UV-photodetector equipped with a 10.2 eV light source which allowed the detection of about 1 ppm of H2S. The catalyst was first flushed with nitrogen at room temperature and then contacted with an hydrogen flow (50 cm3/min). The reactor temperature was then linearly increased to the desired value and left at this temperature for 2h. The amount of H2S released from the solid was quantified after calibrating the detector with a known concentration of H2S diluted in argon. The degree of reduction was defined as the ratio between the amount of H2S eliminated during the reduction process and the total sulfur content. The catalytic properties of the reduced samples have been determined at 273K or 323K using the H2-D2 exchange reaction. This reaction which involved hydrogen activation was chosen because it proceeds at temperatures lower than those required for solid reduction. In a typical run, the solid was reduced at a given temperature and then cooled down to the reaction temperature. The reactor was flushed with nitrogen and then isolated to allow the stabilization of the reactant flows. The catalyst was then submitted to an equimolar mixture of H 2 and D 2 diluted in argon. The H E and D E partial pressures were 76 torr. The variation of the H2-D2 composition was analyzed by means of a mass spectrometer (FISONS Instruments) equipped with a quadrupole analyser working in a Faraday mode. A silica capillary tube heated at 453K continuously bled off a small fraction of the gas phase into the spectrometer. Conversions were calculated either with respect to the decrease of the D2 or of the H 2 signal. For both solids the conversion was kept lower than 20% by adjusting the contact time. It has been verified that under these experimental conditions no H2S, HDS nor D2S was released during a catalytic run indicating that sulfur to metal ratios were stable during the reaction course. The amount of hydrogen retained by the solids after a given pretreatment was determined by thermodesorption by heating abruptly the solid from room temperature up to 573K. Preliminary experiments had shown that this temperature is high enough to desorb the hydrogen present on both solids. Species leaving the catalyst surface were detected by a gas chromatograph equipped with a TCD detector. It was checked that on both solids no H2S was removed during these desorption experiments whatever the degree of reduction of the solids and several adsorption-desorption cycles could be done without changing the adsorption properties of the reduced catalysts. Using this technique two different sets of experiments were performed. The amount of hydrogen retained by the solids during the reduction step was first determined, then the desorbed solids were again submitted to a hydrogen flow at room

120 temperature in order to quantify the hydrogen retained in similar experimental conditions than those used to determine their catalytic properties. The neutron spectra were obtained at the ISIS spallation neutron source, at the Rutherford Appleton Laboratory, U.K., using the spectrometer TFXA. It is a time-of-flight instrument with an inverse geometry and a time-focusing analyser; it gives reasonable counting rates and good energy transfer resolution (AE/E ~, 2%) over a wide range of energy transfers. The precision on the frequencies is _+ 10 cm "1. All the spectra were recorded at 20 K, the evacuation and the loading of the samples being performed out of the cryostat. ESR characterization was performed in situ in order to avoid any contact of the pretreated solids with air. Spectra, recorded as the first derivative of the absorption, were obtained at room temperature or 77K using a Varian E9 spectrometer working in the X band. The g values were measured relative to a DPPH reference (g = 2.0036). The sample tubes were filled with the solid to a height greater than the depth of the resonant cavity and the number of paramagnetic species was calculated by double integration of the recorded spectra normalized to that of Varian Strong Pitch sample (g = 2.0028, 3.1015 spins, cm-l).

3. RESULTS Figure 1 reports the evolution of the degree of reduction of both solids as a function of the temperature of reduction. RuS 2 can be completely reduced at 773K while only 8 % of the sulfur content of M o S E is released at this temperature. For RuS E previous physicochemical characterizations have evidenced that the pyrite phase is stable up to 623K [13]. At higher temperatures, the solid sinters with the concomitant formation of a metallic phase. For MoSE, only a slight decrease in surface area (from 55 to 40 mE/g) has been observed above 573K. 1.0

=

.=o

-

25

RuS9

05

0.8-

-20

RuS2

0.6 o

4

15

2

0.4 ,

I0

1

o.2

0

0.0 273 373 473 573 673 773 Reduction temperature(K)

~6

5 9-

i

1

I

0

1

273 373 473 573 673 773 Reduction t e x t u r e

(K)

Figure 1. Evolution of the degree of reduction Figure 2. HE-D2 activity ofMoS 2 (at 323K) as a fianction of the reduction temperature, and of RuS 2 (at 273K) versus the reduction temperature.

r

121 The changes in the catalytic properties as a function of the temperature of reduction are reported in Figure 2. For both solids, it is clear that as far as the stability of the solids is preserved an increase of the degree of reduction brings about an increase of the H2-D2 conversion. These results confirmed that the presence of coordinatively unsaturated sites is required for hydrogen activation. It should be underlined that in order to determine accurate activity measurements the reaction temperature was not the same for both solids. Taking into account the apparent activation energies (-- 12 kcal/mol) the activity of RuS2 is about two orders of magnitude higher than that of M o S 2. Figures 3 and 5 give examples of hydrogen TPD profiles observed on both catalysts reduced at 473K. The more intense patterns concern the desorption of the hydrogen fixed on the solids during the reduction step (curve a) while the weaker peaks (curve b) are related to the hydrogen retained after readsorption at room temperature. Hydrogen adsorption is thus an activated process. 50

25

- 573

/ ii/--' a?\

- 473 ~

,~ 4 0 -

-20

~.

,~= 30

-15

~

20-

/i : b \ -

0

1

1

I

100 200 300 400 500 "time (s)

-

10

273

:g o

373~

r

Z: I

-

100 -~ 0.0

-5 I

1

1

1

0.2

0.4

0.6

0.8

~!

~0 1.0

Degree of reduction

Figure3. Hydrogen TPD patterns observed on Figure 4. Amount of adsorbed H l and a reduced RuS2 catalyst, catalytic activity (at 273K) versus the degree of reduction of RuS2. On RuS 2 the TPD profiles clearly evidence the existence of at least two different adsorbed species denoted by H x and H 2. Previous work has shown that their relative amount vary with the reduction state of the solid [11-13]. The intensity of the high temperature peak H 2 continuously decreased with desulfurisation while the concentration of H l goes through a maximum (see Figure 4). Interestingly, the variation of both the H 1 concentration and the catalytic activity with the degree of reduction follow a similar trend suggesting that H ~ is the active species involved in this model reaction. By contrast, the smooth profile observed on MoS2 (Figure 5) suggests that only one adsorbed species exists on such a solid. As shown in Figure 6 the exchange activity reasonably correlates with the overall amount of hydrogen retained by the solid during the reduction or after a readsorption at room temperature.

122

500 - 573

rJ2

400 ~ c~ L.

O

- 4 7 3 "=

E =-

300 -

~E - 373 ~

'* .~

200-

t_..

C~ ed2

I

0

1

1

I

100 200 300 400 500 Time (s)

273

0-'1

0.00

~o E

-2

~

-1

:-br

c.;

100-

e',l

-3

!

I

I

0.02

0.04

t-q

0.06

Degree of reduction

Figure 5. Hydrogen TPD patterns observed on Figure 6. Amount of hydrogen adsorbed and a reduced MoS2 catalyst, catalytic activity (at 323K) versus the degree of reduction of MoS2. The INS spectrum recorded at 25K of a reduced and hydrogen desorbed RuS 2 catalyst showed in Figure 7 (spectrum a) exhibits only a weak signal in the range 650-720 cm 1 which can be assigned to bending modes of residual SH. The cell was then taken out of the cryostat, and hydrogen was slowly introduced at room temperature until a pressure of 0.5 bar was reached in the cell. The sample was left overnight to equilibrate and the cell was placed back into the cryostat. In these conditions the amount of adsorbed hydrogen present in the cell was about 150 cm 3. Spectrum b shows that chemisorption gives rise to four peaks found at 540, 648, 719 and 821 cm ~. The intensity enlargement of the peaks at 648 and 719 cm 1 clearly demonstrates that the interaction of hydrogen with the catalyst surface increases the SH group density. The two other bands at 540 and 823 cm ~ cannot be ascribed to new SH groups because these frequencies are well out of the range of usually observed SH bendings. As these vibrations correspond fairly well with metal-hydrogen bending modes already observed on transition metal hydrido-carbonyl complexes, they were assigned to Ru-H species. Figure 8 reports the INS spectra observed on two different reduced states of MoS2. The first one, sample A, was heated under hydrogen flow up to 473 K. The second one, sample B, was reduced at a higher temperature: 673 K. For these experiments about 25g of MoS2 were utilized in order to have a similar or even higher amount of adsorbed hydrogen in the neutron beam with respect to RuS2. The INS spectrum of sample A is similar to the one previously reported by Moyes et al. for the as prepared catalyst [ 15]. The main features can be observed in Figure 8 spectrum a: the peak at 660 cm ~ is assigned to SH bending modes, overtones of these modes are found at 1320 and 2000 cm -~. However, after hydrogen adsorption, some differences can be noticed with respect to the spectra published by the same group [ 16]. In ref. 16, the intensity below 400 cm I was similar in the as prepared catalyst and in the spectrum obtained after adsorption of 1 bar of hydrogen, whereas the peak at 650 cm -~, the SH bending modes, became stronger after hydrogen adsorption. In the present work, the intensity of the SH bands does not increase as the hydrogen pressure is raised, but a peak near 120 cm -~, previously assigned to molecular hydrogen [16], is already observed at 0.5 bar (spectruna b)

123 whereas it appeared only at 10 bars in ref. 16. Furthermore, this peak is splitted into two components at 104 and 117 cm l , indicating a stronger interaction of molecular hydrogen with the catalyst. From the INS intensities, it is found that the amount of SH groups on MoS 2 and RuS2, after reduction at similar temperatures, = 500 K, is comparable on both sulfides. On MoS2, when the temperature of reduction is increased, the number of SH groups decreases (spectrmn c), in agreement with the TPD measurements. Unlike the previous results obtained with RuS2, no clear indication of the presence of Mo-H species can be evidenced by INS. 0

500

1000

1500

I

I

1

2000 cm

-l

500 1000 1500 2000 cm ~ I

I

I

I

~ ,...~

~

b

I

0

50

I

1

100

.,.,~

1

J

150 200

I

l

I

1

50 100 150 200 250 300

250

Energy transfer E l-E2 (meV)

Energy transfer E l-E2 (meV) Figure 7. INS spectra observed on RuS2. a) degassed solid, b) degassed solid + 0.5 bar of H 2.

Figure 8. INS spectra observed on MoS 2. a) sample A, b) sample A + 0.5 bar of H 2, c) sample B + 0.5 bar of H 2. 10

773K

_

@

673K 8-

573K r

~

523K

.,m

*" o,~

473K

7

--

~

6-

298_____~K J

3000

3400 Field (Gauss)

3800

273 373 473 573 673 773 Reduction temperature (K)

Figure 9. ESR signal for various reduced Figure 10. Number of paramagnetic species as solids a function of the reduction temperature.

124 The evolution of the ESR spectra of various reduced and non degassed MoS 2 samples is presented in Figure 9. The ESR spectrum of the freshly decomposed ATTM exhibits an axial signal with gl - 2.04 - 2.06 and a g2 value at 1.98. This signal coincides fairly well with that previously reported for Mo(V) species in a sulfur environment [ 17]. Reduction of the solid up to 573K only affects the signal intensity without any noticeable modification of the g values. At higher reduction temperatures the shape of the ESR is significantly modified by the appearance of a new feature at g --- 2.00. As shown in Figure 10, the presence of this new signal also induces a discontinuity in the evolution of the number of paramagnetic species.

MoS 2 reduced at 573 K

MoS 2

reduced at 773 K

o

before H 2 desorption after H 2 desorption exposed to the air I

1

J

t

I

I

3000

3400

3800

3000

3400

3800

Field (Gauss)

Field (Gauss)

Figure 11. Observed ESR signals after reduction, hydrogen desorption and air exposure. ESR spectra were also recorded after hydrogen desorption. The effect of such catalyst pretreatments on the ESR signal of a solid reduced at 573K and at 773K is illustrated in Figure 11. For solids reduced at TOY 573K, the elimination of the adsorbed hydrogen provokes an increase of the Mo(V) signal. However the resulting number of Mo(V) species remains lower than over the non reduced solid suggesting that both sulfur removal and the presence of adsorbed hydrogen contribute to a decrease of the concentration of Mo(V). As far as hydrogen chemisorption is concerned, these results could be interpreted assuming that a fraction of Mo(V) is transformed into Mo(IV) according to the following overall reaction : Mo(V) + 89H 2

r

.Mo(IV) + 89H §

(1)

As a matter of fact, Mo(IV) has a d 2 electronic configuration and therefore does not give rise to an ESR signal. An opposite effect was observed for solids reduced above 573K since hydrogen removal decreases the signal intensity mostly for the high field species. To account for these results reaction (1) could simply be rewritten by replacing Mo(V) and Mo(IV) respectively by Mo(IV) and Mo(III). While direct evidence of the presence of Mo(III) species was obtained in sulfided catalysts, the observed signal at g -- 2.00 is similar to the one

125 observed by Casewit and Rakowski DuBois on Mo(III) complexes coordinated with sulfur ligands [18]. Furthermore, air exposure of high temperature reduced samples yield to a noticeable change of signal intensity and peak shape while on samples reduced at Tr 573K neither the signal nor the peak shape are affected. These results are in agreement with the expected sensitivity of Mo(III) species towards mild oxidation.

4. DISCUSSION AND CONCLUSION The interaction of dihydrogen with the surface of two model transition metal sulfides have been investigated for various sulfur to metal ratio. Both solids exhibit some similarities as well as great differences. Whatever the solid, some superficial sulfur atoms must be removed to generate an activity and to allow hydrogen chemisorption. However the nature of the hydrogen species interacting with the reduced catalysts differs depending on the nature of the solid under investigation. On RuS 2 two hydrogen species can be detected by thermodesorption and their relative amount depend drastically on the S/Ru ratio. The concentration of the strongly adsorbed species reaches a maximum for a low degree of reduction (-- 0.15) and decreases down to an undetectable level when the surface of the catalyst becomes almost sulfur depleted (-- 0.4). By contrast both the activity and the amount of weakly bonded species reach a maximum for this surface composition. INS characterization has demonstrated that the interaction of hydrogen with a reduced solid increases the SH groups density and leads to the appearance of two Ru-H bending modes at 540 and 823 cm 1. It has been observed that the intensity of both the SH and the high frequency Ru-H species are not significantly affected upon evacuation while the 540 -1 cm band intensity decreases. Therefore this bending mode could be ascribed to the weakly bonded hydrogen species which is probably connected with the catalytic properties. Accordingly, two different adsorption mechanisms may account for the observed data. At low solid reduction an heterolytic dissociation of dihydrogen on a Ru-S pair leads concomitantly to the formation of a Ru-H and SH groups while at high surface reduction an homolytic dissociation on some reduced Ru centers would lead to weakly bonded Ru-H species. Results obtained on MoS2 are quite different. By contrast with RuS2, TPD experiments showed that only one hydrogen species is present on MoS2 independently of its degree of reduction. INS spectra have evidenced the formation of SH groups as well as the presence of molecular hydrogen interacting with the catalyst surface. Unlike the previous results obtained with RuS2, no Mo-H species can be detected. ESR measurements have shown that MoS2+• obtained by decomposition of ammonium thiomolybdate possesses unpaired electrons having a (+V) oxidation state. At moderate temperatures, sulfur removal decreases the concentration of such paramagnetic defects as well as hydrogen adsorption does. The desorption of hydrogen partly restored the Mo(V) signal indicating that an electronic transfer occurs during an adsorption-desorption cycle. At high temperature, a similar adsorption process may also involve paramagnetic Mo(III) species. However the number of such paramagnetic sites never exceeds 0.5 % of the total molybdenum content present on the initial overstoichiometric MoS227. Consequently only a small fraction of the adsorbed hydrogen is activated in such a way and other forms of hydrogen have to be taken into account to assess the amount detected by TPD. For an initial S/Mo ratio equal to 2.27, about 10 % of the sulfur anions should be in a

126 (-I) oxidation state in order to maintain the electrostatic neutrality of the slabs. Therefore this sulfur excess could be either SH groups, bridged anions as well as $22 as suggested by Polz et al. Among these different species, the homolytic adsorption of hydrogen on $22- leading to two SH groups could occur without changing the oxidation state of molybdenum. The obtained results evidence the differences between RUSE and MoS2. The former solid has a pseudometallic behavior whereas for MoS2 redox or acid base properties are involved. This demonstrates the difficulties encountered when the interaction of the reactants and the catalyst are described at a molecular level.

5.REFERENCES

.

.

5.

10. 11. 12. 13.

14. 15. 16. 17. 18. 19.

R.R. Chianelli, Int. Rev. Phys. Chem., 2 (1982) 127. R.R. Chianelli, A.F. Ruppert, S.K. Behal, A. Wold and R. Kershaw, J. Catal., 92 (1985) 56. R.B. Moyes, in: Hydrogen effects in catalysis (Z. Paal and P.G. Menon, eds), Dekker, New York and Basel (1988), pp. 583-607. T. Komatsu and W.K. Hall, J. Phys. Chem., 95 (1991) 9966. L. Jalowiecki, A. Aboulaz, S. Kasztelan, J. Grimblot and J.P. Bonnelle, J. Catal., 120 (1989) 108. A. Wambeke, L. Jalowiecki, S. Kasztelan, J. Grimblot and J.P. Bonnelle, J. Catal., 109 (1988) 320. X.S. Li, Q. Xin, X.X. Guo, P. Grange and B. Delmon, J. Catal., 137 (1992) 185. A.B. Anderson, Z.Y. A1-Saigh and W. K. Hall, J. Phys. Chem., 92 (1988) 803. C. Sampson, J.M. Thomas, S. Vasudevan and C.J. Wright, Bull. Soc. Chim. Belg., 90 (1981) 1215. P. Sundberg, R.B. Moyes and J. Tomkinson, Bull. Soc. Chim. Belg., 100 (1991) 967. M. Lacroix, S. Yuan, M. Breysse, C. Dor6mieux-Morin and J. Fraissard, J. Catal., 138 (1992) 409. H. Jobic, G. Clugnet, M. Lacroix, S. Yuan, C. Mirodatos and M. Breysse, J. Amer. Chem. Soc., 115 (1993) 3654. M. Lacroix, C. Mirodatos, M. Breysse, T. D6camp and S. Yuan, Proc. 10th Int. Cong. Catal. (L. Guczi, F. Solymosi and P. T6t6nyi, eds.) Elsevier, Budapest, 1993, pp. 597609. E. Dieman and A. Mialler, Coord. Chem. Rev., 10 (1973) 79. P. Sundberg, R.B. Moyes and J. Tomkinson, Bull. Soc. Chim. Belg., 100 (1991) 967. P.N. Jones, E. Kn6zinger, W. Langel, R.B. Moyes and J. Tomkinson, Surface Sci., 207 (1988) 159. B.G. Silbemagel, T.A. Pecoraro and R.R. Chianelli, J. Catal., 78 (1982) 380. C.J. Casewit and M. Rakowski DuBois, Inorg. Chem., 25 (1986) 74. J. Polz, H. Zeilinger, B. Mi~ller and H. Kn6zinger, J. Catal., 120 (1989) 22.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

127

Characterization of a Zeolite Membrane for Catalytic Membrane Reactor Application Anne Giroir-Fendler, J6r6me Peureux, Henri Mozzanega and Jean-Alain Dalmon Institut de Recherches sur la Catalyse, 2 av. Albert Einstein, 69626 Villeurbanne Cedex France

Abstract This paper describes the morphological and transport properties of a composite zeolite (silicalite) - alumina membrane. Some advantages obtained in combining the membrane with a conventional fixed-bed catalyst are also reported.

1. INTRODUCTION One of the most studied applications of Catalytic Membrane Reactors (CMRs) is the dehydrogenation of alkanes. For this reaction, in conventional reactors and under classical conditions, the conversion is controlled by thermodynamics and high temperatures are required leading to a rapid catalyst deactivation and expensive operative costs. In a CMR, the selective removal of hydrogen from the reaction zone through a permselective membrane will favour the conversion and then allow higher olefin yields when compared to conventional (nonmembrane) reactors [ 1-3 ]. Following this principle, when using a membrane, lower temperatures can be used leading to a longer catalyst lifetime and energy and cost saving. However such a membrane should be stable at high temperature (ca 500~ highly permselective (loss of reactant should be of course avoided) and permeable enough (the permeation rate should be in the same range as the reaction rate). Dense Pd-based membranes have been first used for CMRs applications [4]. They are indeed highly selective for H 2 permeation but are expensive, sensitive to ageing and poisoning and are strongly limited by their low permeabilities. Classical commercial ceramic porous materials, as those obtained via sol-gel processes, generally have adequate permeabilities but could present some drawbacks. They indeed have a limited thermal stability and are generally not permselective enough: their pores are in the mesoporous range and maximum separation factors correspond to Knudsen diffusion mechanisms. To ensure a better separation, molecular sieving will act much better. This size exclusion effect will require an ultramicroporous (i.e. pore size D < 0.7 nm) membrane. Such materials should be of course not only defect-free, but also present a very narrow pore size distribution. Indeed if it is not the case, the large (less separative and even non separative, if Poiseuille flow occurs) pores will play a major role in the transmembrane flux (Poiseuille and Knudsen fluxes vary as D 2 and D respectively). The presence of large pores will therefore cancel any sieving effect. A zeolite membrane, where the pores originate from the structure, presents only one type of (ultramicro)pore and therefore seems to be a good candidate for CMRs application. Moreover the structural origin of the pores should induce a much better thermal stability of the

128 membrane when compared to sol-gel systems where only the texture originates the pores and controls size evolutions. The preparation of defect-free zeolite membranes is the subject of a recent and intensive research for gas separation and CMRs applications [5]. Beside their use in equilibrium-restricted reactions, CMRs have been also proposed for very different applications [6], like selective oxidation and oxidative dehydrogenation of hydrocarbons; they may also act as active contactor in gas or gas-liquid reactions. This paper describes the morphological and transport properties of a composite zeolite (silicalite) - alumina membrane. Results in CMRs applications are also briefly given.

2. EXPERIMENTAL 2.1 Preparation of the zeolite membrane Commercial porous ceramic tubes (SCT/US Filter Membralox T1-70 [7]) were used in this study as support for the zeolite material. They are made (Figure 1) of three consecutive layers of macroporous a-A1203 with average pore sizes decreasing from the external to the internal layer. A thin toplayer made of mesoporous y-Al20 3 was also present in some samples. For gas permeability, gas separation and catalytic measurements the tubes were first sealed at both ends with an enamel layer before zeolite synthesis. Tubes with porous lengths up to 20 cm were used in this study.

Figure 1: Schematic of a cross-section of a commercial SCT tube used as support. Layers 1, 2 and 3 are made of ~-AI203 and have respective thicknesses of (Bm): 1500, 40, 20 and average pore sizes of (~tm): 12, 0.9, 0.2. Layer 4 (optional) is made of ~/-AI203 and has a thickness of 34 pm and average pore size of 4.5 nm.

The silicalite-alumina membrane was prepared after adding a solution containing the silicalite precursor (i.e. silica + template) to the above-mentioned porous tube (hereafter called support) and a specific hydrothermal treatment performed [8]; under the chosen conditions no material is formed in the absence of the porous support. The tube is then calcined at 673 K for removing the template. 2.2 Characterization of the zeolite membrane SEM an& SEM-EDX analyses have been used in order to observe how and where the new material forms on the alumina support. XRD and 29Si MASNMR studies have been performed for its identification. Porous characteristics of the composite material have been explored using N 2 adsorption-desorption experiments (Micromeritics ASAP 2000M). Transport properties have been studied before and after Si deposition using a rig similar to the one for catalytic testings (Figure 2). Pure gas permeabilities (H2, He, N2, normal and isobutane) were studied by measuring the flux passing though the membrane as a function of temperature and pressure for a constant transmembrane differential pressure (no sweep gas).

129

Figure 2. Rig for gas transport and catalytic measurements Gas separation performances (H2/n-butane , n-hexane/2-2 dimethylbutane) have been measured using a sweep gas (countercurrent mode) in order to increase the permeation driving force (no differemial pressure was used); permeate and retentate compositions (see Figure 2) were analysed using on line gas chromatography. Catalytic testings have been performed using the same rig and a conventional fixed-bed placed in the inner volume of the tubular membrane. The catalyst for isobutane dehydrogenation [9] was a Pt-based solid and sweep gas was used as indicated in Fig. 2. For propane oxidative dehydrogenation a V-Mg-O mixed oxide [10] was used and the membrane separates oxygen and propane (the hydrocarbon being introduced in the inner part of the reactor).

3. RESULTS

3.1 Morphology studies The robe has been first characterised just after the hydrothermal treatment step (i.e. before the calcination). The N 2 isotherm is typical of macroporous materials (Figure 3, curve 1) and the tube is gas-tight.

V(cm31g)

1 0

0:.2

0.)4

P/Po

016

j 0'.8

1

Figure 3. N 2 isotherms before (curve I) and after (curve 2) template removal

130 All the other characterization studies have been performed after the calcination step. XRD ext?eriments have shown that the material formed during the synthesis has the MFI structure. 29Si and 27A1 MASNMR spectra indicated that this phase has a Si/AI ratio varying between 550 and 30 as a function of the sample prepared and also that no extra framework AI is present. SEM micrographs (Figure 4) reveal the presence of small crystallites (Figure 4, C) deposited on the large ot-Al20 3 particles of the thicker layer of the support (layer n~ Figure 1) and of larger ones on the external surface of the ot-Al20 3 layer n~ (Figure 4, B). However, when using as starting material a ceramic tube with a ,/-AI203 toplayer (Figure 1), the formation of crystallites at the external surface is not observed.

Fi~mare4. SEM micrographs of the silicalite-alumina composite material. A: cross-section of the tube. B and C: magnifications of the inner surface of the tube and of the first otAI203 layer.

131 Crosswise SEM-EDX analyses show that the global Si/AI ratio (zeolite + support) varies in a very different way (Figure 5) according to the presence or absence of the ~/-AI203 toplayer in the support.

1 SilAI 0.8

0.4 I

I

10

I

30 L (pm)

I

I

50

Figure 5. Crosswise SEM-EDX analysis in the composite membranes (L, radial distance from the inner surface of the tube). Curve 1 (e) corresponds to support without the ~,AI203 toplayer, curve 2 (r-l) with the 7-A1203 toplayer. N 2 adsorption leads to a type I isotherm (Figure 3, curve 2), typical of microporous systems. The corresponding pore size distribution, as calculated using the Horvath-Kawazoe equation [ 11] is given in Figure 6. A sharp maximum, near 0.6 nm is observed. Other methods of isotherm analysis, such the DFT method [ 12], lead to very similar results in the microporous range, but also reveals the mesoporous domain and part of the macroporous one (pore diameters < 200 nm). Very few pores are present in the mesoporous range. Hg porosimetry experiments completed this characterization and have shown [ 13] that pores corresponding to the intermediate ot-Al203 layers (n ~ 2 and 3, see Figure 1) almost completely disappeared after the zeolite synthesis. 0.003 Porous

volume 0.002 (cm3/g) 0.001

0.5

1.0 Pore

1.5

2.0

2.5

diameter (rim)

Figure 6. Pore size distribution (m) and cumulative pore volumes (..-). Microporous domain, Horvath-Kawazoe equation.

132

3.2 Transport studies The zeolite-alumina tube is no more gas-tight after the thermal treatment. The presence or absence of the y-A1203 toplayer in the starting support does not influence the gas transport properties of the final zeolite-alumina tube. For similar temperature conditions the permeance for nitrogen varies with the applied pressure in a different way according to nature of the sample (i.e. the two starting supports or the zeolite-alumina composite, Figure 7).

1000-

Permeance

100 9

(lO'Tmole/s.Pa.m2 )

_

..2

.

_

10-

i

2

s

Internal pressure (1# Pa)

Figure 7. N2 permeances at 30~ in the tx-Al203 support alone (1, O), in the ot-Al203 support plus the ,/-AI203 toplayer (2, II) and in the zeolite-support composite (3, A). In the zeolite-alumina composite the behaviour of different gases (permeance of pure gases as a function of the temperature, Figure 8) behave very differently from those predicted by ideal Knudsen diffusion processes.

10- ~ c L " ' ~ ' c ~ . . . . . 5 3 . . " Permeance ~ (lO'7mole/s.Pa. 2 )" w _

o' I

i

100

i

i

I

I

200 300 Temperature ('C)

I

I

400

Figure 8. Permeances in the zeolite-support composite for H2 (1, r'l), N 2 (2, A), He (3, II), n-butane (4, O), isobutane (5, o). Intemal pressure 1.2 105 Pa. Gas separations also show non-Knudsen behaviors. In the case of the H2/n-butane mixture, the temperature has a drastic effect on the main permeating gas, at low temperature almost only butane, the heavier component, permeates (Figure 9).

133

When introducing a mixture of n-hexane and 2-2 dimethylbutane (45/55 molar ratio), almost only n-hexane permeates: the permeate contains up to 99.5% of the linear isomer (Figure 10).

Flux 100(10 mole/h)

100

200

300

400

Temperature (~

Figure 9. H 2 (D) / n-butane ( t ) separation with the composite zeolite-alumina membrane (fluxes in the permeate as a function of the temperature). A mixture of hydrogen, n-butane and nitrogen (12 : 14 : 74) was fed in the tube (Fig. 2) with a flow rate of 4.8 l/h. Sweep gas (N2), countercurrent mode, flow rate 4.3 l/h.

12-

Flux

(104mol~)

8-

$_

w

60

i ~

..w

100 Temperature (~

I--

160

l

200

Figure 10. N-hexane (rl) / 2-2 dimethylbutane (m) separation with the composite zeolite-alumina membrane (fluxes in the permeate as a function of the temperature). A mixture of n-hexane, 2-2 dimethylbutane and nitrogen (5 : 6 : 89) was fed in the tube (Fig. 2) with a flow rate of 2 l/h. Sweep gas (N2), countercurrent mode, flow rate 0.5 l/h.

3.3 Catalytic studies Most of the results have been already partly presented in [9] (isobutane dehydrogenation) and [10] (propane oxidative dehydrogenation). Let us recall that the membrane presented in this paper has been associated with a fixed bed catalyst placed within the tube. In the isobutane dehydrogenation the catalytic membrane reactor allows a conversion which is twice the one observed in a conventional reactor operating under similar feed, catalyst and temperature conditions (and for which the performance corresponds to the one calculated from thermodynamics) [9]. In the propane oxidative dehydrogenation, where the membrane separates the two reactants, a 20% increase in the yield was observed with respect to a conventional reactor working at isoconversion [ 10].

134 4. DISCUSSION

4.1 Morphological properties The starting alumina support (Figure 1) is either macroporous or, in the presence of the y-Al203 toplayer, mesoporous. These two materials are highly permeable (Figure 7). The fact that the tubes become completely impervious to any gas aiter the hydrothermal synthesis suggests that the material originated from the synthesis has plugged up the support, resulting in a defect-free gas-tight tube (no cracks or pin-holes). The presence of macropores (Figure 3) suggests however that the porosity (non-passing through pores) has not completely vanished. After the calcination step, experimental data (XRD, 29Si MASNMR) show that a zeolite with the silicalite structure has been formed. 29Si MASNMR indicates for the zeolite material a Si/A1 ratio depending on the sample prepared: it has been observed that both the natures of the silicon source and of the alumina supports may originate these fluctuations. SEM micrographs (Figure 4) show the deposition on the ct-Al203 grains of small crystaUites with the typical hexagonal shape of silicalite. The pore size distribution, as deduced from N 2 adsorption, presents a very narrow peak centred on 0.5 nm, also in good agreement with the pore diameter of silicalite-type zeolites. All these data confirm that a well-defined zeolite silicalite-type crystalline phase has been formed in the presence of the alumina porous tube (which seems indispensable for the zeolite synthesis, as no material is formed in its absence). SEM-EDX analyses (Figure 5) underline the difference of the final material whether the support has a y-Al203 toplayer or not. In its absence, the large zeolite crystals growing on the inner surface of the tube (Figure 4) result in high surface Si/Al ratios. In the porous domain corresponding to layers 2 and 3, the Si/Al ratio is more or less constant and could correspond to a complete filling up of the porous volume (layers 2 and 3 have similar porosities). When analysing the much thicker layer n~ where SEM micrographs show that the zeolite does not fill up the pores, the Si/Al ratio decreases. When using a support with the ~/-AI203 toplayer, the absence of zeolite crystals on the surface, as observed by SEM, is confirmed by the EDX analysis: the Si/Al ratio strongly decreases at the surface. This means that the separative zeolite layer has been here formed only in the bulk of the porous support and not mainly deposited on the outer surface of the support (some infiltration in the support being possible), as is generally the case in zeolite membranes described in the literature [ 14-19]. The presence of the zeolite only within the support could lead to some advantages as far as mechanical and thermal stabilities are concerned. Indeed the separative layer is here protected by the resistant t~-AI203 (thus avoiding damage when introducing for instance catalyst pellets in the tube). The thermal stability of the present zeolite membrane has been also checked: after calcination at 700~ the microporous character remains [ 13 ] and transport properties are not significantly modified after oxidative dehydrogenation of propane at 600~ [ 10] and also after several hours at 850~ under air/steam mixtures [20]. Let us also underline that the zeolite membrane described here is prepared in only one hydrothermal synthesis step. This is not always the case for other preparations reported in the literature for which several successive zeolite synthesis could be needed with the same support to suppress defects.

4.2 Gas transport properties Figure 7 shows that N 2 permeability strongly depends on the pore size. For the macroporous support (curve 1) Poiseuille flow occurs, leading to an increase of the permeance

135 with the pressure. For the mesoporous support (curve 2), Knudsen diffusion rules the transport and permeance becomes pressure independent. When using the microporous zeolite membrane (curve 3) the N 2 permeance decreases when the pressure increases: such a behaviour can be accounted for by activated diffusion mechanisms [21], which are typical of zeolite microporous systems. In such systems the diffusivity depends on the nature and on the concentration of the diffusing molecule which interacts with the surface of the pore. For gases with low activation energies of diffusion, a decrease of the permeability can be observed [22]. Figure 8 indicates that the permeance of n-butane goes through a maximum when increasing the temperature. This observation agrees with the above-mentioned mechanism. At low temperature the nC 4 concentration is high in the zeolite and permeance is low due to the low probability of finding a free site ensuring the mobility. An increase of temperature will then favour the transport. However, when the temperature is too high the coverage of the interacting species decreases and becomes the limiting factor, leading to a global decline of the transmembrane flux. Such phenomena have been also recently described in other zeolite membranes [23], [ 18]. The changes in H2/n-butane separation when increasing the temperature (Figure 9) can be explained on the same basis. At low temperature n-butane interacts strongly with the zeolite and owing to its size completely blocks the pores. Hydrogen has no more room to penetrate the pore and only n-butane diffuses. When increasing the temperature, the occupancy of the zeolite porous volume by n-butane progressively decreases and the hydrogen flux increases. However even at high temperature, the observed separation factor is less than the one expected from Knudsen diffusion. These results underline that the experimental separation factor may strongly differ from the one expected from pure gas permeabilities measurements. However this is not always the case, especially when the two components weakly interact with the surface. When using the membrane to separate a H2/isobutane mixture, the permeation of isobutane, due to its size, is restricted over the entire temperature range and the transmembrane fluxes of the two components of the mixture better follow the permeabilities of the pure gases. Separation factors are here much higher (factors up to 80 have been measured). Molecular sieving effect of the membrane has been evidenced using a mixture of two isomers (i.e. no Knudsen separation can be anticipated), n-hexane and 2-2 dimethylbutane (respective kinetic diameters 0.43 and 0.62 nm). Figure 10 shows the permeate contains almost only the linear species, due to the sieving effect of the zeolite membrane (pore size ca 0.55 nm). This last result also underlines that the present zeolite membrane is almost defect-flee.

4.3 Catalytic properties In isobutane dehydrogenation, when using the zeolite membrane associated with a fixed-bed Pt-based catalyst placed in the inner volume of the tube, the yield is twice the one observed in a conventional reactor [9]. This is especially due to the high separative performance of the membrane which selectively removes the produced hydrogen from the reactor. When using a less selective membrane, for instance the starting support with the 7AI203 toplayer (for which only Knudsen separation occurs), the performance of the CMR is limited by the loss of isobutane through the membrane [9]. The zeolite membrane has also been used in the oxidative dehydrogenation of propane. The goal is here to limit the undesirable complete oxidation owing to the controlled addition of oxygen diffusing through the membrane. Following this mode, it is possible to keep a low oxygen/hydrocarbon ratio all along the fixed bed and then limit complete oxidation. Using the

136 zeolite membrane a 20% increase in the propene yield has been obtained [ 10]. Here also, when using only the starting support as diffusion barrier, it is not possible to control the 0 2 addition and the performance is quite similar to the one observed in a conventional reactor [ 10].

CONCLUSIONS The synthesis of a well-crystallised silicate-type zeolite within the porous volume of an ot-Al20 3 tubular support leads to a defect-flee zeolite membrane showing high thermal and mechanical stabilities. In gaseous mixtures, the separative performance of the membrane greatly depends on the temperature and the nature of the components, especially when strong interactions take place. When occurring, molecular sieving leads to high separative performances. These properties, combined with its high thermal stability, suggest that the zeolite membrane is a very promising material for CMR applications.

Acknowledgements We are indebted to SCT - US Filter for providing alumina support tubes.

REFERENCES [1 ] N. Itoh, AIChE J., 33 (1987) 1576. [2] J.N. Armor, Appl. Catal., 49 (1989) 1. [3] A. Champagnie, T.T. Totsis, R.G. Minet, I.A. Webster, Chem. Eng. Sc., 45 (1990) 2423. [4] V. M. Gryaznov, Plat. Met. Rev., 30 (1986) 68. [5] J.L. Falconer, R.D. Noble and D.P. Sperry Eds, Membrane Separations Technology: Principles and Applications, S.A. Stern and R.D. Noble (Eds), Elsevier, 1994. [6] J-A. Dalmon, Handbook of Heterogeneous Catalysis, G. Ertl, H. Kn6zinger and J. Weitkamp (Eds), VCH, Chap. 9.3 (in press). [7] R. Sofia, Catal. Today, 25 3-4 (1995), 285. [8] J. Ramsay, A. Giroir-Fendler, A. Julbe and J-A. Dalmon, F. Pat. 94 05562 (1994). [9] D. Cazanave, J. Peureux, A. Giroir-Fendler, J. Sanchez, R. Loutaty and J-A. Dalmon, Catal. Today, 25 3-4 (1995), 309. [ 10] A. Pantazidis, J-A. Dalmon and C. Mirodatos, Catal. Today, 25 3-4 (1995), 403. [11] G. Horvarth and K. Kawwazoe, J. Chem. Eng. Japan, 16-6 (1983) 470. [ 12] J.P. Olivier, W.B. Conklin and M.v. Szombathely, Characterization of Porous Solids III, Studies Surf. Sc. Cat., Elsevier, 87 (1994) 81. [ 13] D. Uzio, J. Peureux, A. Giroir-Fendler, J-A. Dalmon and J.D.F. Ramsay, Characterization of Porous Solids III, Studies Surf. Sc. Cat., Elsevier, 87 (1994) 411. [14] H. Suzuki, US Pat 4 699 892 (1987). [15] A. Ishikawa, T.H. Chiang and F. Foda, J. Chem. Soc. Chem. Com., 1989, 12, 764-765. [ 16] M.D. Jia, B. Chen, R.D. Noble and J.L. Falconer, J. ofMembr. Sci., 90 (1994) 1. [17] Y.M. Ma and S. Xiang, US Pat. 5 258 339, 1993. [ 18] F. Kapteijn, W. Bakker, J. Moulijn, H. van Bekkum, Catal. Today, 25 (1995), 213. [ 19] P. M6riaudeau, A. Thangaraj and C. Naccache, Microporous Mat., 4 (1995) 213. [20] A. Giroir-Fendler and J-A. Dalmon (to be published). [21 ] J. Caro, H. Jobic, M. Bialow, J. Karger and B. Zibrowius, Adv. Cat., 39 (1993) 351. [22] J. Peureux, A. Giroir-Fendler, H. Jobic and J-A. Dalmon (to be published). [23] Z.A.E.P. Vroort, K. Keizer, H. Verweij and A.J. Burggraaf, Proc. ICIM4, 1994, 503.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

137

Catalytic Reduction of SO3 Stored in SOx Transfer CatalystsA Temperature-Programmed Reaction Study Gwan Kim a and Michael V. Juskelis b aGrace Davison Division and bResearch Division, W.R. Grace & Co.-Conn. 7500 Grace Drive, Columbia, MD 21044-4098 USA

Abstract A laboratory test method has been developed to assess the efficiency of SO x transfer catalysts for the reduction of SO 3 stored in the catalysts. The test method is based on the temperature-programmed reaction (TPR) with propane over a presulfated catalyst. The reaction products released during the test were determined by means of mass spectrometry (MS). Using this technique, we have appraised the oxides of four transition metals, V, Cr, Fe, and Ce, each impregnated on microspheres of a ternary oxide containing Mg, La, and A1. Based solely on the onset temperature for H2S release, the ranking in catalyst efficiency found was V > Ce > Fe > Cr. In good contrast to this, the ranking based on thermogravimetric analysis (TGA) test for the SO 2 oxidation / SO 3 trapping was Ce > Cr > V > Fe. Using the TPR/MS test method, we have also examined a sample of DESOX TM, the leading commercial SO x transfer catalyst consisting of (MgO)2A1203 and oxides of Ce and V. The onset temperature for H2S release was approximately 450~ when propane was the reactant. It was approximately 580~ when H2 or CH4 was used. These data imply that the rate of H2S release is limited by the supply of a c t i v e hydrogen. Thus, the data further suggest that the primary role of catalytic metal oxide for the reduction of SO3 stored in the catalyst is to facilitate C-H bond cleavage on the catalyst by activating hydrocarbons. 1. I N T R O D U C T I O N The three key steps that determine the performance of a SO x transfer catalyst which controls SO x (SO 2 and SO3) emission from a fluid catalytic cracking unit (FCCU) are (1) the oxidation of SO 2 to SO 3 under the FCCU regenerator condition, typically at 7 0 0 - 730~ (2) the trapping of SO 3 on the catalyst in the form of sulfates, and (3) the reduction of sulfates to release sulfur as H2S in the FCCU riser, typically at 520 - 530~ While the reactions involved in each of these steps are thermodynamically favorable [ 1, 2], steps 1 and 3 must be catalyzed in order to achieve a sufficiently high rate of SO x emission control. The performance of catalysts for steps 1 and 2 combined can be easily assessed by running a TGA test at 700~ [3]. However, no satisfactory test method has been developed for appraising the catalyst efficiency in the laboratory for step 3 under a condition similar to the FCCU riser environment. Some [3, 4] applied TGA to follow the weight loss resulting from the reduction of sulfates with pure or diluted H 2 at 650~ Their data are sufficient to show the reducibility of the sulfates with molecular H 2 under that particular condition, which is far different from a typical FCCU riser condition [5]. Others [6] attempted to simulate their test by including multiple gas oil cracking-stripping-regeneration cycles. Their results, however, lack reproducibility and correlation with catalyst performance in FCCUs. We have developed a laboratory test method whereby the efficiency of a catalyst for step 3 can be readily assessed.

138 To illustrate this test method, we present some of the results obtained from a study of the oxides of four transition metals, V, Cr, Fe, and Ce, which are known to be active for catalyzing step 1 [7].

2. EXPERIMENTAL 2 . 1 . Catalyst Preparation The four catalysts (Samples B-E, Table 1) were prepared by the usual incipient-wetness impregnation with either oxalate or acetate solutions on an oxide containing Mg, La, and A1, followed by 115~ drying and 2h heating in 538~ air. The chemical compositions of all samples were determined by using inductively coupled plasma (ICP) analysis. A master batch of the ternary oxide (Sample A consisting of 37.4% MgO, 18.9% La203, and 42.8% A1203 by weight, Table 1) was prepared by spray drying of a hydroxide slurry, followed by washing, drying and 2h air calcination at 718~ The slurry was prepared by a single-stage coprecipitation at pH 8.5 in a mix-pump reactor, feeding magnesium nitrate, sodium aluminate, and sodium hydroxide solutions simultaneously, followed by aging at pH 9.5. The details of such a preparation have already been disclosed [8]. 2 . 2 . Sulfation of Samples for TPR Test A 0.40g sample was sulfated by running 3h SO 2 oxidation at 704~ in a down-flow Vycor glass reactor with flowing N 2 (126 ml/min) containing 9.50 vol.% O 2 and 0.6000 vol.% SO 2. The sample was cooled in flowing N 2 prior to discharge.

2 . 3 . Propane-TPR/MS Test In the first set of comparison test runs, a 0.10g of each sulfated sample was examined by TPR in a soak (300~ (up to 850~ at 20~ mode, using propane at 14.2 ml/min as the reactant. In the second set of propane-TPR comparison test runs, another 0.10g of each sulfated sample was examined by a soak (300~ (30~ (up to 530~ mode. The reaction products such as SO 2 and H2S released during the course of a TPR run were determined by means of mass spectrometry (Hiden Analytical, HAL-2) in MID mode, monitoring mass numbers 48 for SO fragment from SO 2, and 34 for H2S. 2 . 4 . H 2 and CH4-TPR/MS Tests using H 2 o r C H 4 in Place of Propane In these soak (300~ (up to 900~ mode tests, a 0.10g of sulfated DESOX sample--the leading catalyst based on magnesium aluminate spinel with an excess MgO catalyzed with Ce and V oxides [9]Bwas allowed to react with either n 2 (5%)/A/" (50 ml/min), undiluted n 2 (70 ml/min) or CH 4 (14 ml/min).

2 . 5 . TGA Test for SO 2 Oxidation and SO 3 Trapping Using a TGA unit identical to one already described elsewhere [4], an 8 - 16 mg sample was subjected to flowing N 2 (225 ml/min) at 700~ until the weight stabilized. To this was added an O2/N2 gas mixture followed by an SO~fN2 gas mixture to provide a feed containing 0.73% O 2 and 0.27% SO 2. The weight gain over a period of 15 min exposure to SO 2 oxidation was recorded. 3 . R E S U L T S AND DISCUSSION

3 . 1 . Catalyst Properties The chemical composition from ICP analyses and nitrogen porosimetry data obtained by Quantachrome AUTOSORB-6 are summarized in Table 1 for all the five samples. Note that the catalytic metal oxide loadings in Samples B-E were adjusted so that the efficiency of each catalyst for step 3 can be directly compared on the same basis, per gram atom of metal. The

139 four catalyst samples (B-E) exhibit higher surface areas than Sample A, reflecting a creation of new surfaces resulting from post-impregnations. It is reasonable to expect, however, that the total SO 3 storage capacity as well as the number of SO 3 trapping sites on the surface remain nearly identical for all the five samples. It is also reasonable to assume that none of the four catalytic metal oxides contributes significantly to the SO 3 storage capacity. The powder X-ray diffraction (XRD) data reveal the presence of microcrystaUine MgO and La203 only for all samples. The size of the catalytic metal oxide crystallites is presumed to be quite small at such a low loading, well below the monolayer coverage. Table 1 Property of Catalysts Sample Metal Oxide Loaded (Wt.%) Source S.A. (N2), m2/g Median Pore Diameter, nm

A None 0

B V205 2.49 Oxalate 167 8.3

151 14.5

C D C r 2 0 3 Fe203 2.14 2.13 Acetate Oxalate 193 186 6.6 7.2

E CeO 2 4.22 Acetate 186 6.9

3.2. Propane-TPR/MS Study The propane-TPR/MS results obtained from the soak-ramp-soak mode tests are presented in Figure 1 for the five samples. Some of the results obtained from the soak-ramp-soak mode tests are shown in Figure 2, which emphasizes H:S release at relatively low temperatures

1.0 .

.

.

.

.

.

:

/

---C

I

_ ~()4h

= k-

~/550

"

"

114 I

/ i i(-'"-\V".\~ I .i.."1i

'} I

~'~""+~

600 700 TEMPERATURE ( ~ )

800

Figure 1. Results from propane-TPR/MS tests in soak-ramp mode of sulfated samples,

0

200

400 TIME ( sec. )

600

Figure 2. Results from propane-TPR/MS tests in soak-ramp-soak (530~ mode of sulfated samples.

below 530~ All the data taken from the propane-TPR/MS tests are summarized in Table 2. These data clearly establish V205 to be the best catalyst for step 3, exhibiting the lowest onset temperature for H2S release, approximately 460~ Well-formulated catalysts containing both V and Ce such as DESOX exhibit an even lower onset temperature for H2S release,

140 Table 2 Performance of Catalysts in Laboratory Tests by Propane-TPR/MS and TGA Sample Metal Oxide H2S Release in Propane-TPR Test Onset Temp. (~ Take-off @ or below 530~ Total H2S Released, I.t mol Weight % Gain in TGA Test

A None 600 No 118 3.0

V205 Cr203 Fe203

B

C

E CeO 2

460 Yes 235 5.3

500 No 229 11.3

580 No 212 9.8

D 520 Yes 171 4.5

approximately 450~ which is well below the typical FCCU riser temperature of 520-530~ The next best in the rank solely based on the H2S release onset temperature is CeO v Notice in Figure 2, however, that there is practically no take-off in H2S release over Sample E as long as the temperature remains at 530~ This represents a situation where SO 3 trapped on the surface or stored in the bulk [ 10] is released very slowly, thus creating a condition where the rate of SO x emission control is limited by the number of SO 3 trapping sites. Considerably trailing behind V205 but only slightly behind CeO z is Fe20 3, which shows an onset temperature for H2S release of approximately 520~ This means Fe203 is not expected to be able to adequately catalyze step 3 in the FCCU riser environment because of short contact time [5], even though the temperature at the very bottom of the riser exceeds 530~ In fact, the result of our pilot plant test of such a catalyst is in agreement with this assessment. Thus, it is quite clear that the onset temperatme for H2S release is more critical than the rate of take-off in determining the catalyst efficiency for step 3. Two samples, one (Sample C) with Cr203, the other (Sample A) without any catalytic metal oxide, showed no release at all below 530~ Judging from the pilot plant experience with Fe203-containing catalysts, these two are not expected to be able to function as SO x transfer catalysts. The total amount of H2S released per 0.10 g of sulfated sample during the soak-ramp (up to 800~ mode test, i.e., the integrated area up to 800~ under each plot in Figure 1, is included in Table 2. Because each sample was 3h sulfated and all three steps (1 - 3) are reflected in these data, it is not surprising to see that the results do not parallel with the ranking based on TGA data (also included in Table 2) which cover a 15-rain period of the steps 1 and 2 combined. Nevertheless, we found these data quite useful especially when TGA data for steps 1 and 2 were not immediately available.

3.3. H2-TPR/MS and CH4-TPR/MS Studies The result obtained from a H 2 (5%)/At (95%) - TPR/MS in a soak-ramp mode test is shown in Figure 3 for a sample of DESOX. The onset temperature found for H2S release in this case, approximately 580~ is substantially higher than 450~ the typical onset temperature found in the propane-TPR/MS test. The result was essentially identical in terms of the onset temperature for H2S release even when undiluted H 2 was used as the reactant. Unlike the propane-TPR/MS tests, where the reaction products are essentially H2S only with virtually negligible amounts of SO 2, H2-TPR/MS tests always showed both SO 2 and H2S. These data, notably the pattern of change in the rates of SO 2 and H2S released with temperature in Figure 3, clearly demonstrate, as expected, that the reduction of S § to S 2 in step 3 is a consecutive reaction. The result obtained from a CH4-TPR/MS test of the same DESOX sample is shown in Figure 4. As in H2-TPR, the H2S release is preceded by the evolution of a large amount of SO 2. The onset temperature (580~ of H2S evolution is also comparable to that observed in H2-TPR. It has been presumed that syngas formation [ 11 ] is associated with the generation of H2S in this case.

141

o5-

.,..,

so2 ..

>r

r

~.

"-

SO2 :-"-..

4-

,. .. :9i : :.

H2S i

o 3

9 o

.

~.1

r~

..

2-

r~

<

.

HzS

c~ < ,.d

m 1

..

" '"

.~.

eq .,.,

600

700 800 TEMPERATURE ( ~ )

Figure 3. Results from Hz(5%)/Ar-TPR/MS tests in soak-ramp mode of a sulfated DESOX TM sample.

~

0

, t

400

,,1.~-

.....

60O TEMPERATURE ( ~

800 )

Figure 4. Results from CHn-TPRfMS tests in soakramp mode of a sulfated DESOX TM sample.

3 . 4 . Role of Catalyst in the R e d u c t i o n of S O 3 As we have already pointed out, the onset temperature for the reduction of S O 3 stored in any given catalyst varies with the source of hydrogen. This suggests that the rate of step 3 is limited by the supply of active hydrogen. Since hydrocarbons are the primary source of active hydrogen in the FCCU riser and V +5 can be readily reduced to V +4 [I 2], step 3 can be depicted as consisting of the following scheme of events: Hydrocarbon = (Hydrocarbon - H) § + [H']

(3- 1)

[H'] = H + e

(3-2)

V§ + e = V~

SO 3 + 2 V ' 4 = [SO3-] + 2 V §

(3-3) (3 -4)

[SO3--] = SO2 + O 2

(3- 5)

SO 2 + 0 -2 + 8H = H2S + 3 H20 + 2e

(3 - 6)

(Hydrocarbons - 8H) § + 8e = (Hydrocarbons - 8H)

(3- 7)

The role of catalytic metal oxide is not only to enhance the step from S *6 to S *4, but also, more importantly, to enhance the step from S 44 to S 2, which evidently is relatively slow. To achieve this requires a massive supply of active hydrogen from hydrocarbons by virtue of interactions with the catalyst. This can be represented by combining the first three equations above as follows: Hydrocarbon + V § = (Hydrocarbon - H) § + V ~ + H

(3 - 8)

Our data suggest a possibility that the rate of the entire step 3 is limited by the rate of the event represented by (3 - 8). That is, how fast active hydrogen can be supplied will depend on how

142 readily V § can be reduced to V +4 by hydrocarbons in the FCCU riser. Finally, it should be noted here that propane was chosen in this TPR study not as a representative but as a convenient hydrocarbon that provides hydrogen attached to primary as well as secondary carbon atoms. Had molecular H E been selected as the reactant in this TPR study, the ranking in catalyst efficiency for step 3 would have been quite different. The catalyst efficiency for step 3, then, is likely to be determined by the rate at which H E molecule dissociates as a result of an interaction with the catalyst. 4. CONCLUSIONS (1) In appraising the SO x transfer catalyst efficiency for the reduction of S O 3 (step 3) trapped in the catalyst, molecular hydrogen should not be used as the reductant. (2) TPR/MS technique, using a hydrocarbon other than C~ and C 2 as a reductant, provides a convenient method of ranking the SO x transfer catalyst efficiency for step 3. (3) Based solely on the H2S release onset temperature found in propane-TPR/MS tests, the ranking in catalyst efficiency for step 3 is as follows for the oxides of four metals: V > Ce > Fe > C r. (4) Because the catalyst performance for the overall SO x emission control depends on the number of sites available on the surface for SO 3 trapping at any given time, which, in turn, depends on the rate of step 3 for site regeneration, the rate of step 3 is just as important as that of step 1, the oxidation of SO 2 to SO 3. (5) Step 3 is a consecutive reaction.

REFERENCES 1. J.K. Dixon and J. E. Longfield, in Catalysis, Vol. 7, Ch. 4, P.H. Emmett (ed.), Reinhold Pub. Corp., Baltimore, MD (1960). 2. E.H. Hirschberg and R. J. Bertolacini, in Fluid Catalytic Cracking - Role in Modern Refining, M. L. Occelli (ed.), ACS Symp. Ser. 375, American Chemical Society, Washington, DC (1988) 114. 3. A . A . Bhattacharyya, G. M. Woltermann, J. S. Yoo, J. A. Karch, and W. E. Cormier, Ind. Eng. Chem. Res., 27 (1988) 1356. 4. J.S. Yoo, A. A. Bhattacharyya, C. A. Radlowski, and J. A. Karch, Appl. Catal. B 1 (1992) 169. 5. A.A. Avidan, Studies in Surface Science and Catalysis, Vol. 76, Origin, Development and Scope of FCC Catalysis, J. S. Magee and M. M. Mitchell, Jr. (ed.), Elsevier, New York (1993), p. 36. 6. L. Rheaume and R. E. Ritter, in Fluid Catalytic Cracking - Role in Modem Refining, M. L. Occelli (ed.), ACS Symp. Ser. 375, American Chemical Society, Washington, DC (1988) 146. 7. O.V. Krylov, Catalysis By Nonmetals, Academic Press, New York, (1970), p. 174. 8. G. Kim, SO x Control Composition, US Patent No. 5 288 675 (1994). 9. J.S. Yoo and J. A. Jaecker, Catalyst and Process for Conversion of Hydrocarbons, US Patent No. 4 469 589 (1984). 10. M. Wagif, O. Saur, and J. C. Lavalley, Appl. Catal., 71 (1991) 319. 11. M.V. Juskelis, an unpublished study done using the same TPR/MS technique. 12. G . L . Woolery, A. A. Chin, G. W. Kirker, and A. Huss, Jr., in Fluid Catalytic Cracking - Role in Modem Refining, M. L. Occelli (ed.), ACS Symp. Ser. 375, American Chemical Society, Washington, DC (1988) 215.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 10l 9 1996 Elsevier Science B.V. All rights reserved.

143

Effect of Tunnel Structures of BaTi409 and Na2Ti6013 on Photocatalytic Activity and Photoexcited Charge Separation M. Kohno, S. Ogura, K. Sato and Y. Inoue* Department of Chemistry, Nagaoka University of Technology, Nagaoka, Niigata 940-21, Japan The effect of the surface structures on photocatalytic activity and photoexcited charge separation was studied using titanates with tunnel and non-tunnel structures. BaTi409 with a pentagonal prism tunnel structure, Na2Ti6013 with a rectangular tunnel structure, and Ba4Ti13030 with non-tunnel structure were employed as a titanate, on which RuO2 was supported by an impregnation method to prepare photocatalysts. The stoichiometric production of oxygen and hydrogen from water with irradiation by an Xe lamp occurred only for the photocatalysts with tunnel structure titanates. ESR spectra of BaTi409 and Na2Ti6013 at 77 K in the presence of uv irradiation showed the characteristic peaks assigned to surface O- radicals in the presence of 02 and N 20, whereas no peak appeared for Ba4Ti13030 under similar conditions, thus indicating that the tunnel structure titanates have high efficiency for the separation of photoexcited charges. Raman spectra showed that the TiO6 octhahedra in the tunnel structures had short Ti-O bonds which reflected large distortion of the octahedra. It is demonstrated that the distortion of TiO6 octahedra produces dipole moments nearly directed to the center of the tunnel, which is proposed to form useful electric fields for the separation of photoexcited electrons and holes without recombination, 1. I N T R O D U C T I O N

Recently, we have shown that the combination of barium tetratitanate, BaTi409 and sodium hexatitanate, Na2TisO13, with ruthenium oxides leads to active photocatalysts for water decomposition[I,2]. The unique feature of these photocatalysts is that no reduction of the titanates is required to be activated: this is intrinsically different from conventional photocatalysts using TiO2 which are often heat-treated in a reducing atmosphere. Such different photocatalytic characteristics suggest that efficiency for the separation of photoexcited charges (a pair of electrons and holes) which is the most important step in photocatalysis is

:144 different between the titanates and mio2. Since BaTi409 and Na2TisO13 have a pentagonal prism tunnel and a rectangular tunnel structure, respectively, we have paid attention to the role of the tunnel structures in photocatalysis. A barium titanate, Ba4Ti13030, in a series of Ba-Ti titanates was chosen as a representative with non-tunnel structure, and we have compared the photocatalytic activity and the ability for the production of surface radical species with uv irradiation between the tunnel and the non-tunnel titanates. 2. E X P E R I M E N T A L

SECTION

BaTi4Og, Ba4Ti1303o and Na2Ti6013 were prepared by calcining an equimolar ratio of TiO2(high purity grade, Soekawa Chemical Co.) and the corresponding carbonate(BaCO3, high purity grade, Soekawa Chemical Co. or Na2CO3, extra pure grade, Nakarai Inc.) in air at 1173, 1473 and 1273 K, respectively. The formation of the titanates was confirmed by X-ray diffraction patterns which were obtained with a Rigaku Denki RAD III diffractometer. The photocatalysts were obtained by impregnating the titanates with an aqueous solution of RuCI3, drying at 353 K, then oxidizing at 773 K. The ruthenium metal loading was 1 wt%. The photocatalytic decomposition of water was carried out in a closed gas circulation system connected to a high-vacuum line. About 200 mg of a powdered catalyst was dispersed in pure distilled water and irradiated through a water filter with an Xe lamp operated at 400 W. Hydrogen and oxygen evolved in the gas phase were analyzed by a gas chromatograph which was directly connected to the reaction system. ESR spectra were recorded on a JEOL JES-RE2X spectrometer. About 250 mg of ruthenium-free titanate samples was subjected to heat treatment either in vacuum at 573 K or in a hydrogen atmosphere at 973 K. For measurements of ESR spectra, 30 Torr of 02 or N20 was introduced at room temperature and then cooled to 77 K without evacuation. The spectra were obtained in the dark and under UV irradiation with a 500 W low pressure mercury lamp. UV diffuse reflectance spectra of the titanates were obtained with a JASCO UVIDEC-660 spectrophotometer using a sintered alumina disc as a reference. Raman spectra were recorded at room temperature on a JASCO NR-1100 spectrometer. 3. R E S U L T S A N D D I S C U S S I O N

Figure 1 shows the decomposition of water with uv irradiation on RuO2/BaTi4Og, RuO2/Ba4Ti13030 and RuO2/Na2TisO13 photocatalysts. For RuO2/BaTi409, both oxygen and hydrogen were produced with nearly the stoichiometric ratio. The production continued without deterioration as long as the catalyst was irradiated:

145

120

O

I~ 100

O~

.--, r

80

O q--,

60

O .,i,..,a

o

40

< 20

0

60

120

180 240 Time / min.

300

360

Figure 1. Photodecomposition of water on RuO2/BaTi4Og(O,H2; 0,02), RuO2/Na2TisO13(~,H2; I--I,O2) and RuO2/Ba4Ti13030(A,H2;A,O2). the turnover number of oxygen production amounted to about 50 at a reaction time of 10 h, which indicates that the production of oxygen and hydrogen occurs photocatalytically. RuO2/Na2Ti6013 also showed the stoichiometric production of oxygen and hydrogen, the photocatalytic activity of which was approximately 3 times lower than that of RuO2/BaTi4Og. For RuO2/Ba4Ti13030, a small amount of hydrogen production occurred, but no oxygen evolution was observed. These findings demonstrate that the titanates with tunnel structures have an ability to induce photocatalytic activity for water decomposition. Figure 2 compares ESR spectra of BaTi4Og, Na2TisO13 and Ba4Ti13030. No ESR signals were observed without uv irradiation. In accordance with previous observation[3], uv irradiation of BaTi409 at 77 K in the presence of 30 Torr oxygen

146 O

(A) g=2.018

.

i g=2.004

(B)

O

g=2.018 [ Mn2' marker

] g=2.004

Mn2' marker

x0. g=2.020 -' ~

_

(a) @

~

(b) ,

tip

20G

20G

Figure 2. ESR spectra of BaTi4Og(a), Na2TisO13(b) and Ba4Ti13030(c). In the presence of 30 Torr O2(A) and in the presence of 30 Torr N20(B). All spectra were recorded at 77 K with uv irradiation. produced a strong characteristic peak with g=2.018 and g=2.004, and Na2TisO13 exhibited the same peak in addition to an additional peak with g=2.020.On the other hand, no significant peak was observed for Ba4Ti13030. Figure 2 also shows ESR spectra of BaTi4Og, Na2TisO13 and Ba4Ti13030 exposed to 30 Torr of N20. Exactly the same peaks(g=2.018 and 2.004 for BaTi409 and g=2.018, 2.004, and 2.020 for Na2Ti6013) as those observed in the presence of oxygen were obtained for BaTi409 and Na2TisO13, whereas no peak appeared for Ba4Ti1303o. These g

147

i ~q"

Mn z§ marker

I

Mn 2+ marker O

g= 1.977

I (a)

(b)

20G

Figure 3. ESR spectra of BaTi409 reduced at 973 K in H2 flow and then evacuated at 673 K. In the dark(a), with uv irradiation(b), and in the dark after exposure to 30 Torr of oxygen(c). All spectra were recorded at 77 K.

values of the observed peaks were similar to g=2.021 and g= 2.0026 reported for the adsorption of N20 on ZnO[4] and g=2.020 and g=2.006 for N20 adsorbed on Mo/SiO2[5]. These g values are assigned to an O- species produced by photodecomposition of N20:N20 + e- ---* N2 + O'. Thus, the ESR peaks observed for the adsorption of N20 on BaTi409 and Na2Ti6013 are assigned to a surface Oradical. The most interesting feature is that oxygen also produces the same ESR peak, thus indicative of O- radical formation. The formation of the stable O" radical means that BaTi409 and Na2Ti6013 surfaces have significantly higher efficiency for the separation of photoexcited charges than Ba4Ti1303o, and it is rationally considered that this capability is related to the tunnel structures. Figure 3 shows ESR spectra of BaTi409 which was subjected to reduction in a hydrogen atmosphere at 973 K. Without uv irradiation, there are a large broad peak at g=1.977 and a small peak with g=2.018 and g=2.004. Since the g value of the large peak is consistent with that(g=1.957-1.992) reported for colloidal TiO2 by Howe and Gratzel[6], this peak is assigned to Ti 3+. The small peak is due to O-. With uv irradiation, a considerable decrease in the intensity of Ti 3+ peak and

148

/ /

l(/

~

/// It./

l// I!; ...... _.

9'

I

300

|

400

I

500

Wavelength / nm

Figure 4. UV diffuse reflectance spectra of BaTi40~( Ba4Ti13030(. . . . ).

), Na2Ti6013( . . . . ) and

complete disappearance of the O- peak occurred. The BaTi409 surface was exposed to 30 Torr of oxygen at room temperature, and then ESR spectra were measured at 77 K: the Ti 3+ peak decreased remarkably, but the O- peak recovered. Such changes in peak intensity indicate that the electron transfer occurs between Ti 3+ and O- by uv irradiation. Figure 4 shows the uv diffuse reflectance spectrum of Ba4Ti13030 together with that of BaTi409 and Na2Ti6013. Ba4Ti13030 has a threshold wavelength of absorption at around 400 nm and a maximum at around 320 nm. The general feature of absorption spectrum is similar to that of BaTi409 and Na2Ti6013, and thus it is evident that negligibly small photocatalytic activity of Ba4Ti13030 is not related to differences in the characteristics of absorption. However, there was a difference in that both BaTi409 and Na2Ti6013 had a shoulder peak between 390 and 330 nm, whereas no shoulder appeared for Ba4"1713030.This possibly reflects the difference in the structure between tunnel and non-tunnel. Figure 5 compares Raman spectrum of Ba4Ti13030 with that of BaTi409. In the range of 300 to 1000 cm -1, the Raman spectrum of BaTi4s consists of three main peaks[3]. The first peak appears in the region 430 - 450 cm -1, and the second in

149

,.-:,.

r~

1000

I

I

1,

800

600

400

W a v e n u m b e r / crn ~

Figure 5. Raman spectra of BaTi4Og(a) and Ba4Ti1303o(b). 590-650 cm -1. The most characteristic peak is the third one, because of a strong single peak at a higher frequency of 860 cm -1. Interestingly, the Raman spectrum of Na2TisO13 was similar to that of BaTi4Og, and the characteristic peak appeared at 870 cm -1(not in figure). Ba4Ti13030 shows complex spectrum: a peak observed at 820 cm -1 is weak and broad. Dehnicke reported that the Raman spectrum of TiOCI2 had a strong peak at 836 cm -1, which was associated with the stretching vibration of a short Ti-O bond[7]. From analogy with this assignment, it is demonstrated that short Ti-O bonds are involved in the TiOs octahedra in BaTi409 and Na2TisO13, suggesting that the TiOs octahedra are distorted. Figure 6(a) shows schematic projection of BaTi409 and a pentagonal prism tunnel structure, which were obtained by X-ray diffraction of a BaTi409 single crystal[8]. There are two types of l-lOs octahedra in BaTi 409: each of the octahedra is distorted in different ways to a such extent to which a Ti ion is displaced from a center of gravity of surrounding six oxygens. This leads to the generation of two dipole moments. The shortest Ti-O bond length is 0.177 rim, which gives the characteristic Raman peak at around 860 cm -1. Figure 6(b) shows schematic projection of Na2TisO13 and the rectangular tunnel structure[9]. Na2Ti6013 has three kinds of TiO6 units, each of which is distorted to produce different dipole moments. The directions of the dipole moments are

150

C

C

a

a

Figure 6.

(a) Schematic

(b) projection

and tunnel

structure of

BaTi4Og(a)

and

Na2TisO13(b). 9 atom; O,titanium atom; ~,barium or sodium atom; • ,center of gravity of six oxygen ions. The arrows indicate the dipole moments.

drawn in the figure, and it is interesting to see that the dipole moments of both BaTi409 and Na2Ti6013 are directed nearly toward the center of the tunnel. Table 1 demonstrates the distance of displacement from the center of gravity and the magnitude of the dipole moments. The dipole moments of BaTi409 were D=5.7 and D=4.1 and those of Na2Ti6013 were D=5.3, D=5.8 and D=6.7, whereas all the moments of Ba4Ti13030 was below D=5.0. It is to be noted that there is a clear difference in the dipole moment between tunnel structure titanates(D>5.0) and non-tunnel structure Ba4Ti1303o (D<5.0). The ESR spectra of BaTi409 showed

151 Table

Dipole moments of octahedra

Titanate

BaTi409

Displacement nm

Til

Dipole moment D

0.030

..............................

.T_.i.z ............... .0..:..0...2...1.. .........................................

NazTi6013

Til

Ba4Tii303o

5.7 ..4.....1. .................

0.027 5.3 Ti2 0.030 5.8 Ti~.................0.._0..3_.5. .....................................6....7.................. Til 0.021 4.0 Ti2 0.000 0 Ti3 0.016 3.1 Ti4 0.025 4.8 Ti5 0.023 4.4

a single peak(g=2.018), whereas those of Na2TisO13 two peaks(g=2.018 and g=2.020). Such a different anisotropic character of the O radicals is considered to be associated with the difference in tunnel structure between BaTi,O9 and Na2Ti6013. The presence of the dipole moments corresponds to the formation of internal local electric fields in the TiO6 octahedra, and the interesting feature is that the field works to promote the separation of photoexcited charges. Similar situation holds for Na2Ti6013 with the rectangular tunnel structure. These effects lead to high efficiency for charge separation. The previous study using high resolution transmission electron microscopy showed that the spherical RuO2 particles of 1.43.0 nm in diameter were uniformly distributed over the BaTi 409 surface: a model is proposed in which the pentagonal prism tunnel space provides the form of a nest, i.e., a concave site with a ridge, and offers the accommodation sites of the RuO2 particles. According to this model, a strong interaction between RuO2 and the BaTi409 surface is possible. Since the dipole moment in the TiO6 units forming the tunnel is directed nearly toward the center of the tunnel, the internal electric fields also work to facilitates the transfer of photoexcited electrons to RuO2 particles located at the tunnel sites. This mechanism explains a unique feature that high photocatalytic activities of RuO2/BaTi409 and RuO2/Na2Ti6013 are generated without reduction.

152

Acknowledgement This work was supported by a Grant-in-Aid for Scientific Research on Priority Areas from The Ministry of Education, Science, Sports and Culture.

REFERENCES 1. Y.Inoue, Y. Asai and K. Sato, J.Chem.Soc.Faraday Trans.,90(1994)797. 2. Y.Inoue, T.Kubokawa and K.Sato, J.Phys.Chem.,95(1991)4059. 3. M.Kohno and Y.Inoue, to be submitted to J.Chem.Soc.Faraday Trans. 4. N.B.Wong, Y.B.Taarit and J.H.Lunsford, J.Chem.Phys.,60(1974)2148. 5. V.A.Shvets and V.B.Kazansky, J.Catal.,25(1972) 123. 6. R.F.Howe and M.Gratzel, J.Phys.Chem.,89(1985)4495. 7. K.Dehnicke, Z.Anorg.AIIg.Chem.,309(1961)266. 8. D.H.Templeton and C.H.Dauben, J.Chem.Phys.,32-5(1960) 1515. 9. S.Andersson and A.D.Wadsley, Acta Cryst., 15(1962) 194.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

153

O r g a n o c h r o m i u m C o m p l e x e s in H o m o g e n e o u s Olef'm P o l y m e r i z a t i o n G. Bhandari, J. L. Kersten, R. R. Kucharczyk, P. A. White, Y. Liang, and K. H. Theopold* Department of Chemistry and Biochemistry, Center for Catalytic Science and Technology, University of Delaware, Newark, DE 19716, USA (e-mail: [email protected]) The advent of single-site, so called "metallocene" catalysts is revolutionizing the field of coordination polymerization of ethylene and a-olefins.[1] Molecular design of the organometallic catalyst precursors has enabled unprecedented control of the various structural elements of the polymer (i.e. molecular weight, dispersity, incorporation of side chains, tacticity) that determine its physical properties. However, despite the commercial importance of chromium based polymerization catalysts, and in contrast to group 4 (Ti, Zr, HI) chemistry,[2] little is known about the relevant organometallic chemistry of chromium, and the development of novel organochromium catalysts is only beginning. Part of the problem is the paramagnetism of relevant chromium alkyls in those oxidation states commonly thought to be catalytically active (i.e. Cr II, CrlII). The limited utility of NMR spectroscopy for the characterization of paramagnetic organometallics has generally discouraged the exploration of their chemistry and thus hindered the elucidation of the structural requirements for catalytic activity. Fundamental issues like formal oxidation state, coordination number, and chemical nature of the ligands of active chromium catalysts are still relatively ill defined. Seeking to shed light upon these matters, some time ago we initiated a broadly based investigation of the organometallic chemistry of paramagnetic chromium alkyls, with a particular emphasis upon polymerization catalysis. Herein we summarize recent results from our laboratory concerning homogeneous olefin polymerization catalysts containing chromium. We have previously described a class of cationic chromium(HI) alkyls of the type [Cp*Cr(L)2R]+X -, which catalyze the polymerization of ethylene at ambient temperature and pressure.[3] The most efficient of these catalysts was the complex shown below. i. . . . .

i

1§ BPh4I

~

'~0 ' ' 9Cr

'CH 3

This complex and structurally related molecules served as a functional homogeneous model system for commercially used heterogeneous catalysts based on chromium (e.g. Cp2Cr on silica - Union Carbide catalyst). The kinetics of the polymerization have been studied to elucidate mechanistic features of the catalysis and in order to characterize the potential energy surface of the catalytic reaction. *Correspondence to this author. This research was supported by grants from NSF and Chevron Chemical Co.

154 1. ACTIVE OXIDATION STATE The literature is sharply divided regarding the formal oxidation state of chromium containing active sites. Based, apparently, on some early determinations of the average valence of heterogeneous catalyst preparations,[4] and abetted by the lack of well-characterized homogeneous model systems, the notion of divalent chromium (Cr la) as an active oxidation state remains popular. Neglecting the possibility that the designation "active site" may sometimes be applied to an immediate precursor to, or resting state of, the actual catalyst, we wish to address the oxidation state of the actual chain-propagating species, i.e. presumably a chromium alkyl capable of binding and inserting ethylene into its chromium-carbon bond. While "oxidation state" is acknowledged to be a formalism that does not correspond to an actual charge on the metal atom, nor even to a specific electron density or level of electron deficiency (witness the wide range of Crm/Cr II redox potentials depending on coordination environment), there nevertheless seems to be a correlation between formal oxidation state and reactivity. Among the stable oxidation states of the transition metal chromium (i.e. II, IN, and VI) the +I11 state is by far the most common and the most stable under normal conditions. If compounds in this oxidation state can be shown to be catalytically active and kinetically competent to account for the productivity of commercial catalysts, the need for postulating more esoteric oxidation states would seem to be diminished. Wherever possible, we have sought a direct comparison of the reactivities of structurally related Cr III and Cr II alkyls with ethylene. For example, after having established the catalytic activity of complexes of the type [Cp*CrlII(L)2R] + (see above), we showed that the isostructural neutral compounds Cp*Crn(L)2R did not polymerize ethylene; instead facile 13hydrogen elimination was observed.[3] This difference in reactivity was not due to the charge of the complexes. Thus, we have subsequently shown that neutral Cr m alkyls are also active polymerization catalysts. For example, Cp*CrlIIoT-tF)Bz2 and even anionic Li[Cp*Crrn(Bz)3] (Bz = benzyl) polymerized ethylene at ambient temperature and pressure, while the structurally related CpCrla(bipy)Bz proved inert.[5] Our attempts to prepare chromium hydrides and to evaluate their role in polymerization catalysis eventually led to the isolation of a series of alkyls and hydrides lacking any ancillary ligands besides the cyclopentadienyl moiety (see below).[6] Reduced to the essence of Cr I! alkyls, these complexes provided another piece of evidence in the growing case against polymerization activity of divalent chromium; none of the alkyls even reacted with ethylene. The hydride underwent one insertion and stopped at the stage of an ethyl group.

/Cp" ,,

\-Z/

~)"

9 Cp"

It might well be argued that the pronounced metal-metal bonding exhibited by these extremely electron deficient compounds attenuates the reactivity toward external substrates, such as olefins. A hypothetical, monomeric "Cp*CrR" (with a 12-electron configuration)

155 would certainly be an interesting molecule to try in this regard. Unfortunately, such compounds are incapable of existence, due to their extreme electron deficiency and steric unsaturation, as reflected in the actually observed dimeric structures. However, a sterically more demandin2g ancillary ligand might yet facilitate isolation of a monomeric and coordinatively unsaturated Cr u alkyl (see below). A further dramatic comparison of the comparative reactivities of chromium alkyls in diverse oxidation states was furnished by another set of benzyl complexes.[7] Shown below, these three compound are isomers, yet they range in oxidation state from Cr I to Cr In. Of the three, only the mixed-valent complex Cp*Cr~-rl6:rl3-Bz)Cr(Bz)Cp *, containing a trivalent chromium bound to an rl3-benzyl and a rll-benzyl ligand, catalyzed the polymerization of ethylene.

~

Ph

Ph

PE While we have now found many well-characterized Cr III alkyls which catalyze the rapid polymerization of ethylene under very mild conditions, none of the various Cr II alkyls we have prepared have shown any indication of polymerization activity. Those who would continue to invoke Cr II as an "active oxidation state" for ethylene polymerization should be called upon to provide experimental support for that proposal.

2. (z-OLEFIN POLYMERIZATION An surprising feature of the Cp*-system described initially was its high selectivity for polymerization of ethylene over wolefins; i.e. [Cp*Cr(THF)2CH3]+BPh4" neither polymerized propene nor showed any tendency in copolymerization experiments with ethylene/o~-olefin mixtures to incorporate t~-olef'ms. We have adopted several strategies to overcome this lack of reactivity. Reasoning that 0~-olefms are more weakly bound and thus compete less efficiently with free THF for the coordinatively unsaturated chromium, we have prepared even more labile complexes containing dialkyl ether ligands - e.g. [Cp*Cr(R20)nCH2SiMe3]+BAr'4" (R = Me, Et, ~Pr; A t ' = 3,5-(CF3)2C6H3). +

R --

RP"O ~ R ....~Si.

R'~ R

o Cr

R

v

"~ ....~Si

n

156 In solution, these compounds spontaneously lose ether to generate coordinatively unsaturated cationic 13 electron Cr0II) alkyls, which indeed polymerize a-olef'ms (e.g. propene, 1-hexene) at temperatures slightly below ambient. The molecular weights of the polymers were rather low, however, and so were the productivities and turnover frequencies. Consistent with the notion of efficient discrimination between ethylene and alkyl substituted olefins, copolymerization experiments with these catalysts also yielded ethylene homopolymer without any indication of a-olefin incorporation into the polymer. These complexes are extraordinarily reactive ethylene polymerization catalysts; e.g. [Cp*Cr(Et20)2CH2SiMe3]+BAr4 - will polymerize liquid ethylene (T < -100~ In the presence of any ethylene, a-olefins cannot compete for binding and/or insertion. Finally, the extreme electrophilicity of the coordinatively unsaturated 13-electron chromium ions made the catalysts susceptible to decomposition; degradation of the coordinated ether molecules resulted in the formation of chromium bound alkoxides, leading to catalyst deactivation. Another attempt at encouraging a-olefin binding and polymerization is the preparation of socalled "constrained geometry" catalysts.[8] In such molecules (see below), tight linkage of the cyclopentadienyl group to another ligand via a short tether opens up the remaining coordination sites and is thought to make the metal more accessible for olefin binding. t!

-

t!

--

( X ) ~ " )N~M\

R

,~

--

I

\N....--Cr .,,,, '~R L

R

Constrained geometry catalysts containing group 4 metals are noted for their high rates of r olefin incorporation into copolymers. We were curious, if this effect would carry over to chromium catalysts. Accordingly, we have prepared a series of chromium alkyls containing the chelating (TiS-Me4Cs-SiMe2-NtBu)-ligand. Figure 1 depicts the molecular structure of the representative (rlS-Me4Cs-SiMe2-NtBu)Crrn(THF)Ph, as determined by X-ray diffraction.

Figure 1. The structure of (rlS-Me4Cs-SiMe2-NtBu)CrlII(THF)Ph; selected distances (/I,) and angles (deg): Cr-C(16), 2.126; Cr-(N1), 2.011, Cr-O(1), 2.106; N-Cr-Cp(centroid), 110.6.

157 The THF ligand of (rl5-Me4C5-SiMe2-NtBu)CrIII(THF)Ph is only weakly bound, and dissociates in solution. Indeed, with sterically more demanding alkyl groups base-free fivecoordinate alkyl complexes, e.g. (rl5-Me4C5-SiMe2-NtBu)CrCH2SiMe3, could be prepared. Constrained geometry chromium alkyls catalyzed the polymerization of ethylene; however, the reaction was relatively slow, and elevated pressures (PC2H4 = 500 psi) were required to generate significant amounts of polymer. Not surprisingly then, no homopolymerization or copolymerization of (x-olefms was observed. Instead, catalytic isomerization and dimerization of the alkyl-substituted olefms was found. Based on our observation in these two systems, it would appear that Cp*CrlII-alkyls, if rendered electrophilic and/or sufficiently coordinatively unsaturated, will both bind and insert (x-olefins. However, the more heavily substituted alkyl ligands thus formed (i.e. CrlII-CH2CHfR)-P vs. CrlILCH2-CH2-P resulting from ethylene insertion) seem to be very susceptible to facile [~-hydrogen elimination. Rapid chain transfer and very low molecular weights are the results of this tendency. Whether the latter is an innate property of all chromium alkyls or reflects the particular chemical nature of the Cp*Cr-moiety remains to be established. To this end, investigation of chromium alkyls with a variety of other ancillary ligands are needed.

3. CHROMIUM ALKYLS WITH NITROGEN- AND OXYGEN-LIGANDS 3.1. Tris(pyrazolyl)borate complexes In an attempt to change the electronics of the chromium atom, we are replacing the carbon based cyclopentadienyl ring with ligands containing harder donor atoms. For example, we have employed the tris(pyrazolyl)borate moiety, an isoelectronic replacement for Cp* featuring tridentate N-coordination.J9] Figure 2 shows the molecular structure of Tpt-Bu,MeCr-Ph, a representative Cr rl alkyl. It will be noted, that this complex is mononuclear, due to the steric protection of the extremely bulky tris(pyrazolyl)borate.

Figure 2. The molecular structure of Tpt-Bu,MeCr-Ph; selected distances (/~) and angles (deg): Cr-C(46), 2.123; Cr-Navg, 2.12; N(6)-Cr-C(46), 171.8;

158 Keeping in mind the isoelectronic nature of Tp and Cp, Tpt-Bu,MeCr-Ph represents an analog of the elusive mononuclear "Cp*CrR" fragment (see above). An unusual structural feature of this four-coordinate 12-electron complex is the extreme distortion from the expected C3v geometry. Rather than adopting a pseudo-tetrahedral geometry, the complex assumes a "cisdivacant octahedral" structure, i.e. it can be thought of as being derived from an octahedral complex by removal of two of its ligands. Note the nearly perfect trans-disposition of the phenyl ligand to one of the pyrazole nitrogens (i.e. N(6)). This peculiar structure creates an open coordination site for binding of olefin substrate. Indeed, if there was any doubt about the easy accessibility of the chromium in this coordination environment, it was allayed by the isolation and structural characterization of both Tpt'Bu,MeCr-Cl(a structural analog of Tp tBu,MeCr-Ph) and its five-coordinate pyrazole adduct Tpt'Bu,MeCr(t'Bu,MepzH)C1,and by the observation that the four-coordinate chromium alkyls reversibly add pyridine (see below).

a,

tBu

Cr\ "~'Bu R

t B / ~Cr'*l~'tBu

tBu

The alkyls Tpt-Bu,MeCr-R arethe best test case yet of the catalytic activity of Cr rl alkyls (see Section 1). However, they did not react with ethylene, even at elevated temperature. On the contrary, Tpt-Bu,MeCr-Et eventually decomposed by an apparent [3-hydrogen elimination yielding Tpt-Bu,MeCr-H and ethylene. Thus our notion that divalent chromium alkyls are not the chain propagating species in polymerization catalysis receives further support. The obvious next step was oxidation of the tris(pyrazolyl)borate chromium alkyls to the catalytically active +111oxidation state. However, cyclic voltammetry experiments did not show a reversible oxidation in any case, and all attempts to p r e l ~ e complexes of the type [Tp tBu Me 9 Cr-R] +X- by chemical oxidation failed, yielding [Tp-t flu ,MeCr('H-tF)rd +X- instead. The reasons for the apparent instability of TpCr III alkyls are not clear, and we are continuing our efforts to isolate related compounds. 3.2. [ C p * 3 C r 3 0 4 ] - an oxide support mimic Commercial chromium-based polymerization catalysts consist of chromium complexes supported on oxidic supports, such as silica. Ancillary Ligands featuring oxygen coordination of chromium might thus claim more immediate relevance for modeling catalysts of the Phillips type. Such complexes might also allow us to address some of the unresolved questions surrounding the initiation events, which transform a reduced chromium ion lacking any chromium-carbon bonds into a catalytically active site, simply upon contact with ethylene. Serendipit~ has afforded us a novel oxygen tripod ligand, namely the "open cube" [Cp*3Cr304] z-. Like the tris(pyrazolyl)borates (see Section 3.1.) this ligand is isoelectronic with Cp*-, being a tridentate 6-electron donor. However, the harder O-donors, combined with the lack of n-accepting ability and the dianionic charge of the ligand, should make for a dramaticaUy different reactivity of any chromium alkyls supported by this cluster.

159 The triply protonated form of the new ligand, i.e. [Cp*3Cr3(~t3-O)(kt2-OH)3]+C1 ", was prepared by addition of aqueous NaOH to a solution of [Cp*Cr(I.t-C1)]2. Its crystal structure was determined by X-ray diffraction and the result is depicted in Figure 3.

Figure 3. The molecular structure of the oxygen tripod ligand [Cp*3Cr3(I.t3-O)(I.t2-OH)3]+C1The [Cp*3Cr304] 2- fragment may be envisioned as resulting from excision of a Cp*Cr 2+ group (i.e. a Cr IN fragment) from the intact cube Cp*4Cr4(I.t3-O)4, a well known compound first prepared by Bottomley et al.[10] Indeed, in a proof of principle, [Cp*3Cr3(~t3-O)(~t2OH)3]+C1 - was reacted with Li+[Cp*Cr(NEt2)3] -. The cubane cluster Cp*4Cr4(I.t3-O)4 was formed in high yield, thus proving that metal fragments can be introduced into the empty site through a threefold acid/base elimination reaction. Next, we hope to incorporate a coordinatively unsaturated Cr m alkyl in the same position. It is hoped that such a compound will be both a structural and functional model of oxide supported polymerization catalysts. 4. C O N C L U S I O N S We have prepared and structurally characterized a variety of paramagnetic chromium alkyls in different oxidation states. To be active homogeneous catalysts for the polymerization of ethylene, the formal oxidation state of chromium must be +1II, and an open coordination site for binding of the substrate is required. Cr rl alkyls are inactive, even if highly electron deficient and coordinatively unsaturated. Pentamethylyclopentadienyl-ligated Cr III alkyls react with aolefins, albeit much more slowly than with ethylene. The alkyls resulting from a-olefin insertion tend to suffer facile I$-hydrogen elimination, thus yielding only dimers or oligomers of olefins. Chromium complexes of nitrogen- and oxygen-ligands have been prepared to evaluate the reactivity of chromium alkyls in harder coordination environments. While those studies are still in progress, our results are beginning to define parameters for the design of new chromium-based polymerization catalysts.

160 5. A C K N O W L E D G M E N T S All crystal structures were determined in the laboratory of Prof. A. L. Rheingold. We thank Drs. D. Beach and M. Carney of Chevron Chemical Co. for helpful discussions. 6. REFERENCES

1. A.M. Thayer, Chemical & Engineering News, Sept. 11, 1995, 15 - 20. 2. H.H. Brintzinger, D. Fischer, R. Mtilhaupt, B. Rieger, R. M. Waymouth, Angew.Chem. Int. Ed. Engl., 34 (1995) 1143. 3. B.J. Thomas, S. -K. Noh, G. K. Schulte, S. C. Sendlinger, K. H. Theopold, J. Am. Chem. Soc., 113 (1991) 893. 4. a) F. J. Karol, G. L. Karapinka, C. Wu, A. W. Dow, R. N. Johnson, et al, J. Polym. Sci., Part A-l, 10 (1972) 2621. b) H. L. Krauss, H. Stach, Inorg. Nucl. Chem. Lea., 4 (1968) 393. 5. G. Bhandari, Y. Kim, J. M. McFarland, A. L. Rheingold, K. H. Theopold, Organometallics, 14 (1995) 738. 6. R.A. Heintz, R. L. Ostrander, A. L. Rheingold, K. H. Theopold, J. Am. Chem. Soc., 116 (1994) 11387. 7. G. Bhandari, A. L. Rheingold, K. H. Theopold, Chem. Eur. J., 1 (1995) 199. 8. a) P. J. Shapiro, E. Bunel, W. P. Schaefer, J. E. Bercaw, OrganometaUics, 9 (1990) 867. b) P. J. Shapiro, W. D. Cotter, W. P. Schaefer, J. A. Labinger, J. E. Bercaw, J. Am. Chem. Soc. 116 (1994) 4623. c) J. A. M. Canich, Process for Producing Crystalline Polycx-olef'ms with a Monocyclopentadienyl Transition Metal Catalyst System, US Patent No. 5,026,798 (1991). d) J. C. Stevens, D. R. Neithamer, Metal Complex Compounds, US Patent No. 5,132,380 (1992). 9. S. Trofimenko, Chem. Rev., 93 (1993) 943. 10. F. Bottomley, J. Chen, S. M. Macintosh, R. C. Thompson, Organometallics, 10 (1991) 906

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

161

Studies in Surface Science and Catalysis, Vol. 101 9 1996Elsevier Science B.V. All rights reserved.

S e l e c t i v e c a t a l y t i c o x i d a t i o n w i t h a i r of g l y c e r o l a n d d e r i v a t i v e s on p l a t i n u m m e t a l s

oxygenated

Peter Fordham, R~gis Garcia, Mich~le Besson and Pierre Gallezot Institut de Recherches sur la Catalyse-CNRS, 2 Avenue Albert Einstein, 69626 Villeurbanne, France Abstract Platinum metal catalysts, promoted or not with bismuth, have been designed to favour oxidation of either the primary or the secondary alcohol functions of glycerol in aqueous solution. Thus, on palladium catalyst oxidation of the primary function of glycerol predominates and glyceric acid is the chief product (70% yield at 90% conversion, p H = l l ) w h e r e a s on bismuth-promoted platinum catalyst oxidation of the secondary function prevails to give dihydroxyacetone (37% yield at 75% conversion, pH=2). Subsequent oxidation of glyceric acid on platinum gives the tartronate (61% yield at 94% conversion, pH=10-11); improved yields were obtained with the bismuth-promoted catalyst (83% yield at 90% conversion, pH=10-11). On bismuth-promoted platinum oxidation of glyceric acid gives hydroxypyruvic acid as the main product (64% yield at 75% conversion, pH=3-4) and under the same conditions tartronic acid is oxidised to mesoxalic acid (29% yield at 53% conversion, pH=l.5). Bismuth adatoms have two roles: they prevent over-oxidation by strongly adsorbed acids and orientate the selectivity towards oxidation of the secondary alcohol function. 1. INTRODUCTION The catalytic oxidation of alcohols and polyols in aqueous solution on platinum m e t a l s using molecular oxygen is well known and has been extensively studied [1-4] but its prospective use for the production of fine chemicals and accompanying potential for cleaner technology processes has stimulated renewed interest. Attractive features include the mild conditions required, the lack of noxious effluents and the feasibility for catalyst recycling. Platinum or palladium are generally the most effective metals for this type of reaction but the addition of certain heavy metal promoters (e.g: Bi, Pb) has been shown to play a useful dual role in changing the selectivity of the reaction and reducing c a t a l y s t deactivation. The method has shown great utility in carbohydrate chemistry e.g: for the oxidation of glucose to gluconic acid [5-7]. Glycerol represents a highly functionalised molecule which is also a freely available, biosustainable resource being manufactured from triglyceridecontaining crops. Thus, it is a clear candidate for use as a feedstock in the manufacture of derived oxygenates. Present commercial processes employ stoichiometric reactions using mineral acids, or enzymatic pathways, to oxidise glycerol. However, the use of air and platinum group metal catalysts

162 supported on carbon to oxidise aqueous solutions eliminates the production of potentially polluting or toxic effluents. Furthermore, control of the process by employing promoters, adjusting the pH of the reaction medium, or by carrying out subsequent oxidations of derived products, may yield a wide range of useful products. Preliminary studies on the oxidation of glycerol to glyceric acid and dihydroxyacetone have been published [8] and, elsewhere, the use of this method for the synthesis of dihydroxyacetone, using batch and continuous systems, has been described [9,10]. A number of patents relating to oxidation of the primary functions of glycerol have also recently appeared [11,12]. The main objective of the work described here was to control the chemoselectivity of metal catalysts for the oxidation of the p r i m a r y and secondary alcohol functions of glycerol (GLY) to give the full range of possible oxygenated derivatives: glyceric acid (GLYAC), dihydroxyacetone (DHA), hydroxypyruvic acid (HYPAC), tartronic acid (TARAC) and mesoxalic acid (MESAC) (see Scheme 1) and thus to illustrate the considerable potential of this mode of oxidation for the production of fine chemicals. This was achieved by using tailored platinum or palladium catalysts where the metal surface was modified with bismuth adatoms. OH

OH

O

,

O H HO~OH

/

OH

OH

~d

GLYAC

~,,~

0 GLY

HO

OH

TARAC

OH

OH

MESAC

0 OH

DHA

~

HO

0

OH HYPAC

Scheme 1. Target oxygenated derivatives of glycerol. 2. EXPERIMENTAL 2.1 P r e p a r a t i o n o f catalysts

Platinum catalysts were prepared by ion-exchange of activated charcoal. A powdered support was used for batch experiments (CECA 50S) and a g r a n u l a r form (Norit Rox 0.8) was employed in the continuous reactor. Oxidised sites on the surface of the support were created by t r e a t m e n t with aqueous sodium hypochlorite (3%) and ion-exchange of the associated protons with Pt(NH3)42§ ions was performed as described previously [13,14]. The palladium catalyst mentioned in section 3.1 was prepared by impregnation, as described in [8]. Bimetallic PtBi/C catalysts were prepared by two methods: (1) b i s m u t h was deposited onto a platinum catalyst, previously prepared by the exchange method outlined above, using the surface redox reaction: 3(Pt - H) + BiO + --~ (Pt 3 - Bi) + H3 O+

163 This was achieved by s u s p e n d i n g the catalyst in a glucose solution, in a nitrogen atmosphere, and adding the required volume of BiONO3 dissolved in hydrochloric acid, or (2) by coimpregnation of the s u p p o r t with solutions of H2PtC16 and BiC13 in HC1 (1.2M) at 5~ over a period of five hours, and s u b s e q u e n t reduction by adding aqueous solutions of formaldehyde a n d K O H [8]. The m e t a l loading of c a t a l y s t s was d e t e r m i n e d by atomic a b s o r p t i o n spectroscopy, following dissolution, and the size of the m e t a l particles was established by high resolution transmission electron microscopy (TEM) w i t h a Jeol 100CX microscope. The composition of individual m e t a l particles was d e t e r m i n e d with 1.5 nm spatial resolution, on t h i n sections cut w i t h an u l t r a m i c r o t o m e , u s i n g a field-emission s c a n n i n g t r a n s m i s s i o n e l e c t r o n microscope (STEM) VG HB 501 and an energy dispersive X-ray (EDX) analyzer. 22, Oxidation p r o c e d u r e s Oxidations of aqueous solutions (0.1 - 1 tool 1-1) of glycerol, glyceric acid, d i h y d r o x y a c e t o n e and t a r t r o n i c acid were carried out in a t e m p e r a t u r e controlled glass batch reactor fitted with stirrer, gas supply system, oxygen electrode (Ingold) and pH electrode (Radiometer) (see [8] for a scheme of the reaction vessel and a t t a c h m e n t s ) . The s u b s t r a t e / m e t a l m o l a r ratios w e r e maintained at -500-600, unless stated otherwise. The catalyst was suspended in 300 ml of an aqueous solution of the reactant, with a constant flow of nitrogen bubbling through the slurry, and heated to 50*C whilst stirring continuously at 1200 rpm. Once the required t e m p e r a t u r e was attained, the gas supply was switched to air (0.75 ml rain -1) and, where necessary, the pH was m a i n t a i n e d at a constant value by the addition of a solution of sodium hydroxide via a p u m p controlled by the pH meter. The t e m p e r a t u r e , pH and oxygen p r e s s u r e were recorded throughout, and samples of the reaction mix were t a k e n at i n t e r v a l s for analysis. Aqueous solutions of glycerol were also oxidised in a small-scale continuous r e a c t o r ( M i c r o c a t a t e s t MCB 890, Vinci Technologies). A s c h e m e for the circulation of fluids is shown below.

-"!GAS FLOW ] CONTROLLER

~

REACTOR V=10ml

~ ~

~-

A

/

OVEN GAS

SEPARATOR RESERVOIR

PUMP

LIQUID

Figure 1. Continuous/trickle-bed reactor: circulation of fluids. The reactor consists of a small cylindrical metal tube (length: 100 mm, internal diameter: 7 mm) heated by an electric oven and m o u n t e d w i t h i n an assembly which accommodates supply lines, gas/liquid s e p a r a t o r , p u m p and

164 gas flow controller. An aqueous glycerol solution is mixed with air and passed through the catalyst bed within the reactor. The solution is subsequently collected in a separator and sampled at intervals. Periodic evacuation of the separator was possible by means of a pressure-release valve, thus allowing the system to remain on-stream for periods of up to several weeks. The degree of conversion of the reactant and the concentration of oxidised products were carefully monitored by HPLC analysis. Sample components were separated on an ion-exchange column (Sarasep Car-H, 300mm x 7.8mm i.d.) using a 0.0004M H2SO4 mobile phase and detected by UV and RI detectors mounted in series. 3. RESULTS AND DISCUSSION 3.1 Catalyst characterization TEM of microtome sections of the catalysts prepared by ion-exchange showed t h a t the metal particles were in the size range 1 - 2.5 nm and homogeneously distributed in the microporous support. STEM-EDX analysis established t h a t PtBi/C c a t a l y s t s t h u s p r e p a r e d consisted of particles of h o m o g e n e o u s composition. Figure 2(a) shows a TEM view through a section of a PtBi/C catalyst prepared by ion-exchange/redox and Figure 2(b) shows a very high magnification micrograph of the coimpregnated catalyst which reveals the microstructure of the charcoal support. Those bimetallic catalysts which were prepared by coimpregnation gave a less homogeneous particle size distribution with some big crystalline particles of i r r e g u l a r shape, but a n a l y t i c a l microscopy showed that these particles were also bimetallic.

Figure 2. TEM images of PtBi/C catalysts prepared by (a) ion exchange/surface redox reaction, and (b) coimpregnation.

165 3.2 Reaction conditions Monitoring of the oxygen pressure during reaction indicated t h a t the rate of conversion of glycerol to glyceric acid u n d e r basic conditions (see section 3.3) was limited by oxygen mass transfer. All other reactions were free from gasliquid diffusion control; but this does not exclude the possible l i m i t a t i o n by intra-porous diffusion.

3~ Oxidation of glycerol Oxidation of the p r i m a r y alcohol function of glycerol to give glyceric acid (reaction a, see Scheme 1) was described in an earlier piece of work [8]. The selectivity for glyceric acid was found to be dependent on the pH of the reaction m e d i u m , with the h i g h e s t yields being obtained u n d e r basic conditions on palladium catalyst (see Figure 3 overleaf and Table 1). P l a t i n u m catalysts were found to be more active t h a n palladium, but less selective; the difference in activity being attributed to the higher redox potential of platinum leading to a lower oxygen coverage of the surface and t h u s i n c r e a s i n g the a d s o r p t i o n coefficient of the organic substrate. Table 1 Activity and selectivity data Catalyst Prepn. method

Reaction a pH

Concn. Initial (tool 1-1) rate b

Main Sel. c Conv. c product (%) (%)

Pd/C

impreg,

a

11

1

109

GLYAC

PtBi/C

coimpreg,

b

1.5-2

1

52

DHA

PtBi/C

coimpreg,

c

10-11 0.1

302

TARAC

9C3

93

PtBi/C

coimpreg,

d

3-4

0.1

752

HYPAC

95

37

Pt/C

exch.

e

6

0.3

42 (70d)

HYPAC

82

40

PtBi/C

exch./redox

f

1.5

0.1

218 (363 d) MESAC >95

<10

a d

80

89

>95

<10

see Scheme 1 b mol h -1 mo1-1 (PdfPt) c at m a x i m u m selectivity t u r n o v e r frequency, calculated from m e a n particle size m e a s u r e d by TEM

O x i d a t i o n of t h e s e c o n d a r y alcohol f u n c t i o n of glycerol to give dihydroxyacetone (reaction b, Scheme 1) was performed on PtBi/C catalyst. Figure 4 shows the product composition plotted against time for an aqueous solution of glycerol (1 mol 1-]) on 7%Pt3%Bi/C p r e p a r e d by coimpregnation. When the reaction is performed under acidic conditions (initial p H - 4 , dropping rapidly to pH=l.5-2), selectivities approach 80% at the beginning of the reaction with an initial rate of 52 mol h -1 mo1-1 (Pt). However, selectivity decreases with time as the catalyst undergoes a deactivation process, which m a y be due to the accumulation of strongly adsorbed acid products (hydroxypyruvic acid, etc.) on the metal surface. The highest yield was 37% at 75% conversion.

166

+ GLY

100

o=

60 ~

o,~

o

40

E

~ ~9

9 DHA

80-

8O

20

= O

60

Cn

o

40-

E o C.)

20-

* HYPAC

!-

0 t

0

1

2

3

9G L Y A C

~-

0

4

6

12

18

24

time (h)

time (h) Figure 3. Product composition vs. time for glycerol oxidation under basic conditions on 5%Pd/C.

Figure 4. Product composition vs. time for glycerol oxidation under acidic conditions on 7%Pt3%Bi/C.

In an effort to increase conversion and t h u s improve yields of dihydroxyacetone, this reaction was transferred to a continuous reactor. An aqueous glycerol solution (0.2 mol 1-1) was passed over a 4%Ptl%Bi/C catalyst (2g) supported on charcoal pellets and held within a reactor maintained at 50~ Under 1 MPa air pressure and with an air flow rate which was varied in order to maintain a constant oxygen/glycerol mole ratio of 1, the liquid feed rate was varied. Figure 5 shows the plot of percentage composition against contact time. High conversions of glycerol were achieved but over-oxidation of products ensured t h a t yields in dihydroxyacetone were low. The plot of selectivity for dihydroxyacetone vs. glycerol conversion for batch and continuous systems is depicted in Figure 6. 100 +

~"

8O

= o

83

.,... o,.., ~n

E o

100~

u

40

DHA

~ >

~

4o

a

i

. 10

.

. 20

. 30

i

40

contact time (gPt.h/ml) Figure 5. Composition vs. contact time for glycerol oxidation to dihydroxyacetone on 4%Ptl%Bi/C in a continuous reactor.

t

c

h

continuous~

.

20

0~ 0

80

+GLY

0

!

I

i

!

29

40

60

80

100

conversion (%) Figure 6. Selectivity for dihydroxyacetone vs. conversion of glycerol, on PtBi/C, for batch and continuous systems.

167 3.4 Oxidation of glyceric acid U n d e r basic conditions, oxidation of glyceric acid (calcium salt) on platinum gave high yields of the tartronate (61% at 94% conversion, pH=10-11) (reaction c, Scheme 1). Even higher yields were obtained on bismuth-promoted p l a t i n u m p r e p a r e d by coimpregnation (83% yield at 90% conversion, pH=10-11) (see Figure 8). This improved performance was attributed to reduced over-oxidation of the t a r g e t e d product on PtBi/C; significantly higher levels of the overoxidatiion product, oxalic acid, were evolved on Pt/C. C a l c i u m t a r t r o n a t e was p r e c i p i t a t e d and hence s a m p l e s r e q u i r e d acidification prior to the filtration step necessary to remove the catalyst. The chief product of over-oxidation was oxalic acid. However, conversion to oxalic acid proceeds at a relatively low rate and yields of the former are consequently high. This is probably partly due to the tartronate being precipitated, effectively hindering further oxidation.

100 so 6o

9GLYAC 9T A R A C

4o

= OXAC

0 0

6

12

18

24

time (h) Figure 8. Product composition vs. time for glyceric acid oxidation at pH=10-11 on 5%Pt2%Bi/C. Under acidic conditions, bismuth-promoted platinum may be used to oxidise the secondary function of glyceric acid (calcium salt) (reaction d, Scheme 1). Thus oxidation proceeds rapidly to give high yields of hydroxypyruvic acid (64% at 75% conversion, pH=3-4) (Figure 9). A proposed mechanism for this reaction is depicted in Scheme 2, where the formation of a complex between b i s m u t h atoms and the glycerate anion leads to selective oxidation of the secondary alcohol function. The high activity (see Table 1) is assumed to be partly due to the high affinity of the glycerate for the surface of the b i s m u t h - p r o m o t e d platinum catalyst (see Scheme 2). However, at high conversion, over-oxidation gives glycolic acid and residual levels of formic acid. Likewise, the r e a d y a d s o r p t i o n of hydroxypyruvic acid onto the catalyst's surface probably contributes to its subsequent rapid oxidation to glycolic acid. The reduced catalyst deactivation compared to the analogous oxidations of glycerol and tartronic acid was attributed to the use of the calcium salt r a t h e r t h a n the free acid. A recent publication describes a similar observation for the oxidation of sodium gluconate [15]. Sodium ions were a s s u m e d to counter catalyst deactivation by neutralizing the acid species responsible.

168

1009

"~

GLYAC

80

9H Y P ~ C

60

-aLYco

40

O

,.,0,

O

\\ /CH2OH C--C / / ", O O H

Bi +

~

l

I

i

I

0

6

12

18

24

time (h) Figure 9. Product composition vs. time for glyceric acid oxidation at pH=3-4 on 5%Pt2%Bi/C. 3.5 O x i d a t i o n o f ~ n i c

Scheme 2. Proposed m e c h a n i s m for oxidation of t h e s e c o n d a r y alcohol function of glyceric acid.

acid

Tartronic acid was oxidised to mesoxalic acid on 6%Pt2%Bi/C, p r e p a r e d by exchange/redox, under acidic conditions (reaction f, Scheme 1) (29% yield at 53% conversion, pH=l.5). Figure 10 shows t h a t the conversion rate of tartronic acid is high at first but decreases as the reaction proceeds, probably because the formed mesoxalic acid is more strongly adsorbed on the surface t h a n tartronic acid. The initial high selectivity tapers off due to over-oxidation. i(.}0 -,

~"

8O

o= .p-4

60

9 TARAC

op.., O3

o

- MESAC

40

E ~~9

20

0

6

12

18

24

time (h) Figure 10. Product composition for the oxidation of tartronic acid, obtained at pH=l.5 on 6%Pt2%Bi/C, as a function of time.

169

3.6 Oxidation of dihydroxyacetone Dihydroxyacetone is not stable under the basic conditions preferred for oxidation of the primary function to give hydroxypyruvic acid (reaction e). U n d e r acidic conditions the rate of oxidation of a 1 tool l-I aqueous solution is very slow (5 tool h -I mol'1(Pt)). O n platinum the initial rate of conversion for reduced concentrations of the starting material (0.3 tool I-1), whilst retaining the s a m e amount of catalyst, was 42 tool h -z mol-1(Pt), as might be expected under non-favourable acidic conditions. Hydroxypyruvic acid is evolved with a selectivity of 8 2 % at 40% conversion (see Figure 11).

I00 80-

m

DHA

60-

*

HYPAC

A GLYCO

~

- OXAC

o

f

20 0

~--~~,

0

6

'

I

I

12

18

"'"'......

-

I

24

time (h) Figure 11. Product composition vs. time for dihydroxyacetone oxidation at pH=6 on 5%Pt/C. 4. CONCLUSIONS This work shows t h a t a wide r a n g e of glycerol-derived o x y g e n a t e d substances, which have useful properties p e r s e (e.g: as chelating agents) or can be used as building blocks in industrial synthesis, may be produced by catalytic oxidation with air of aqueous solutions of glycerol on platinum group metals. At present none of these products have a large m arket because they are p r o d u c e d by costly biotechnology processes. The r e s u l t s o b t a i n e d are encouraging since high yields of a given product can be obtained by catalytic oxidation on metals and selectivity may be controlled by modifying the metal surface with bismuth adatoms. However, conversions and selectivities should still be improved to avoid separation processes. Improvements can be obtained by decreasing the deactivation process and preventing the occurrence of side reactions. Deactivation is mainly due to the st rong adsorption of acids, especially diacids, on the surface which block f u r t h e r reaction. B i s m u t h a d a t o m s have two useful effects (1) they decrease over-oxidation, probably because they bind the carboxylic acid function thus preventing them to adsorb on p l a t i n u m and to react further (2) they orientate the selectivity towards oxidation of the secondary alcohol function (e.g: reactions b and d) because of the formation of surface complexes such as t h a t depicted in Scheme 2. A promising future way to block deactivation could be to neutralize the acid

170 functions with cations. Thus, by increasing the rate of the main reaction, the amounts of by-products, essentially smaller molecules and ultimately carbon dioxide and water formed by over-oxidation, will be reduced. Acknowledgements: The award of an EC studentship, under the H u m a n Capital and Mobility scheme, to Peter Fordham is gratefully acknowledged. REFEICJENCES 1. K. Heyns and H. Paulsen, Adv. Carbohydr. Chem., 17, 169-221 (1962). 2. H. v a n Bekkum, in Carbohydrates as Organic Raw Materials (F. W. Lichtenthaler, ed.), 1991, pp.289-310, VCH Verlag, Weinheim. 3. T. Mallat and A. Baiker, Catalysis Today, 19,247-84 (1994). 4. P. Gallezot and M. Besson, in Carbohydrates in Europe (Carbohydr. Res. Foundation, ed.), 13, 10-15, 1995, The Hague, Netherlands. 5. P. C. C. Smits, B. F. M. Kuster, K. van der Wiele and H. S. van der Baan, Appl. Catal., 33, 83-96 (1987). 6. P. Fuertes and G. Fl~che, European Patent 233 816 (1987) 7. M. Besson, F. Lahmer, P. Gallezot, P. Fuertes and G. Fl~che, J. Catal., 152, 116-121 (1995). 8. R. Garcia, M. Besson and P. Gallezot, Appl. Catal., 127, 165-76 (1995). 9. H. Kimura and K. Tsuto, Appl. Catal., 96, 217-28 (1993). 10. H. Kimura and K. Tsuto, Appl. Catal., 105, 147-58 (1993). 11. T. Imanaka, H. Terasaki, A. Fujio and Y. Yokota, Japanese Patent 05 331 100 (1993). 12. H. Kimura, T. Imanaka and Y. Yokota, Japanese Patent 06 279 352 (1993). 13. P. Gallezot, R. de M~sanstourne, Y. Christidis, G. Mattioda and A. Schouteeten, J. Catal., 133, 479-85 (1992). 14. P. Gallezot, F. Fache, R. de M~sanstourne, Y. Christidis, G. Mattioda and A. Schouteeten, Stud. Surf. Sci. Catal., 75, 195-204 (1993). 15. A. Abbadi and H. van Bekkum, Appl. Catal., 124, 409-17 (1995).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

171

Selective methylation of catechol: catalyst development and characterisation L.Kiwi-Minsker, S.Porchet, P.Moeckli *), R.Doepper and A.Renken Institute of Chemical Engineering, Swiss Federal Institute of Technology (EPFL) *JCeramic Laboratory, Dept. of Material Science, EPFL CH-I O15 Lausanne (Switzerland)

Abstract

Gas-phase methylation of eatechol by methanol was studied on y -A1203 modified by the basic elements: K, Li, Mg and Ca. Addition of 7.5 at.% Mg to ),-A1203 was optimal and increased the 3methyl catechol selectivity from 0.26 to 0.65. X-ray diffraction experiments showed the diffusion of Li+ and M g 2+ cations into the 7 -A1203 bulk. This induces a change in the surface species (XPS data) and the surface acid-base properties (rPD cxpcrirnents).Ca 2+ and K § addition to y-alumina was ineffective due to formation of basic oxide layers on the surface.

I. I N T R O D U C T I O N Alkylcatechols are important as chemicals and chemical intermediates in the fine chemistry industry for the synthesis of flavouring agents, agricultural chemicals and pharmaceuticals [ 1]. 3-methyl cateehol has a special value from the industrial point of view. Previously y-alumina was found to be an effective catalyst for the gas-phase methylation of catechol with methanol [2]. The process can be schematically presented as:

H

H + CH3OH

(A1)

(A4) CH3

polymethylated products

H3 (A2)

H (A3)

3

Figure 1. Reaction scheme of catechol methylation.

172 Direct methylation to 3-methyl catechol (A3) and 4-methyl catechol (A4) as well as guaiacol (2-methoxyphenol) (A2) formation are parallel reaction pathways. The reverse reaction of guaiacol leads to the initial products catechol (A~) and methanol. The rearrangement of guaiacol in the sequential reaction yields 3-methylcatechol (A3). To avoid the formation of polymethylated products the reaction conditions were optimised by using a conversion of less than 0.3 and a temperature range of 260-310~ At low conversion (<0.05) the rate of rearrangement was negligible and guaiacol was the main product in this pathway. The carbon (C)-alkylation was mainly ortho-selective ( preferential A3 formation ). The purpose of this work was to increase the A3 selectivity at low conversion through a catalyst modification. Previous studies of phenol alkylation with methanol (the analogue reaction) over oxides and zeolites showed that the reaction is sensitive to acidic and basic properties of the catalysts [3-5]. It is the aim of this study to understand the dependence of catalyst structure and acidity on activity and selectivity in gas phase methylation of catechol. Different cations such as Li § K +, Mg 2§ Ca ~-~, B 3§ incorporated into T-AI203 can markedly modify the polarisation of the lattice and consequently influence the acidic and basic properties of the surface [5-8] which control the mechanism of this reaction.

2. EXPERIMENTAL

2.1. Catalyst preparation Pure y-alumina from Engelhard (A1-3982) was used as starting material. The modified catalysts were prepared by wet capillary impregnation with aqueous solutions of nitrate salts. H3BO3 was used in the case of boron impregnation. The concentration of the solutions were adjusted to get the required ratio cation/Al 3§ in atomic percentage (at.%) by varying the amount of nitrates added to T-alumina. Before the impregnation alumina was dried at 200~ for 2 hours. After impregnation (2 h, 25~ and drying (50~ 12 h) the samples were calcinated in air for 8 hours at 620~

2.2. Catalytic test The catalytic reaction was carried out at 270~ and 101.3 kPa in a stainless steel tubular fixed-bed reactor. The premixed reaction solution, with a molar ratio catechol : methanol : water of 1:1:6, was fed into the reactor using a micro-feed pump. To change the residence time in the reactor, the catechol molar inlet flow ff~0) and the catalyst mass (m~t) were varied in the range 10-2 < F~0 <10 mol.h t and 2.10 -2 < n~t < 3.10 t kg. The products were condensed at the reactor outlet and collected for analysis. The products distribution was determined quantitatively by HPLC (column Nucleosil 5C~s, flow rate, 1 ml.mint, operating pressure, 18 MPa, mobile phase, CH3CN : H20 = 1:9 molar ratio). Experiments at different flow rates and with different catalyst grain sizes confirmed that the reaction kinetics is not influenced by external or internal mass transfer. Catechol conversions (X) were always less than 0.05 allowing the reaction to be carried out in the differential kinetic region. The initial yields (Y~0) for the monomethylated isomers were measured under steadystate conditions (after 8-10 hours of the catalyst activity stabilisation) and were used to compare the catalysts selectivities:

173

s,,o-

Rio -Rio

Yi,o/r' = - - X/r'

where r'- mcat -

F~o R,o ,

-RIo

Y,.o

ri.o

i = 2,3,4

X

is the contact time of catechol (h.kg.moll); are reaction rates of i and catechol under inlet conditions (mol.kg-~.hl); is the initial yield of i.

Imrinsic activities (am) of different catalysts were used to compare catalytic activities: a,,, = - R w / SBer

(mol.m'2.h "1)

where SBer

is a catalyst specific surface area (m2.g'l).

2.3. Catalyst characterisation Specific surface areas of the catalysts used were determined by nitrogen adsorption (77.4 K) employing BET method via Sorptomatic 1900 (Carlo-Erba). X-ray diffraction (XRD) patterns of powdered catalysts were carried out on a Siemens D500 (0 / 20) diffractometer with Cu I ~ monochromatic radiation. For the temperature-programmed desorption (TPD) experiments the catalyst (0.3 g) was pre-treated at different temperatures (100-700 ~ under helium flow (5-20 Nml-min ~) in a micro-catalytic tubular reactor for 3 hours. The treated sample was exposed to methanol vapor (0.01-0.10 kPa) for 2 hours at 260 ~ The system was cooled at room temperature under helium for 30 minutes and then heated at the rate of 4 ~ "~. Effluents were continuously analyzed using a quadruple mass spectrometer (type QMG420, Balzers AG).

3. RESULTS AND DISCUSSION 3.1. Catalytic activity and selectivity Catalytic activity and selectivity of pure and modified aluminas are presented at Figure 2. It can be seen that the activity and selectivity of the aluminas used depend on the cation introduced. 5at.% of basic elements decreases the alumina activity as follows: 2.7 times for Mg 2§ 7 times for Li § 40 times for Ca 2§ 340 times for K § B-modified alumina, being more acid than y-AI203, demonstrates higher starting activity with respect to pure alumina, but it deactivates to a greater extent and has about 10% less activity under steady-state conditions. Pure MgO is almost inactive under the conditions used ( see Table 1 ). The selectivity changes in a more complex way. For K § Ca 2+- and B3+-modified alumina no change in selectivity is observed. In the case when Li + cation is added to y-alumina, A3 product selectivity increases to 0.4 from 0.26 for pure y-AI203. The highest selectivity of A3 product (0.62) is observed for Mg2+-modified alumina. The catalytic activity of Mg2+-modified alumina decreases with increasing Mg 2§ content in the samples as seen from Figure 3.

174

Figure 2. Dependence of intrinsic catalytic activity, a~ and 3-methyl catechol selectivity, $3,0 on the cation type used. All cations were added at 5 at.% content.

Table 1. Characteristics of the different catalysts Catalyst

Ion content

Activity, a~

Initial selectivities

(atomic-%)

(mol-m-2-h "1)

sz0

s3,0

$4,0

A1203

-

2.09-10-6

0.71

0.26

0.03

MgO

-

2.87-10 -9

_

_

_

B 3. / A1203

5

8.71-10 7

0.75

0.24

0.01

Ca 2§ / A1203

5

5.15-10 "s

0.74

0.24

0.02

K + / A1203

5

6.13-10 -9

0.70

0.25

0.05

Li § / A1203

5

2.93-10-7

0.60

0.39

0.01

Mg 2+ / A1203

5

7.56-10 "7

0.34

0.62

0.04

This can be explained by the loss of acidity of pure alumina through the addition of the basic element. Literature reports [2-5] show that acidic oxides such as V205, SiO2, Al203 are active for methylation reaction below 400~ Basic catalysts such as MgO are active only beyond 500~

175

"7

E

2.0-

~

1.0-

"

0.8 ._.

2~ o.e ~t 0.4._~ ~9

0.2 0.0 6

"

~

'

1'0

"

;~

"

s

Mg 2+ content [atomic-%] Figure 3. Catalytic activity of Mg2+-modified alumina; temperature

=

270~

The selectivity s3,0 is strongly dependent on the Mg 2§ content (Fig. 4). It increases with Mg concentrations up to 7.5%, passing through a maximum of s3.0 = 0.65. 3-methyl catechol becomes the main product at low conversion. It is worthwhile to note that only the relative formation of guaiacol and 3-methyl catechol are affected by the percentage of Mg 2+ added to alumina, O-allq, lation is reduced in favour of C-alkylation, but the ring methylation stays preferentially ortho-selective.

0.8 0.7-

l

""" Z

9

A'~"'~

S3,0

....

9

0.6-

0.5"

:~

0.4

o

"~

0.3

A/

' m--- t/

Sz0

.--_ 0.2 0.1

$4,0 V" . . . . . .

0.0

I

0

NiL . . . . . . "

9 "V'--

9 ......

I

'

9

................................ "

5 ;0 ;5 Mg 2* content [atomic-%]

9 "

I

20

Figure 4. Initial selectivity of the monomethylated isomers as function of Mg 2§ content.

176

3.2. Catalyst characterisation

BET Pure alumina exhibits a BET surface area of 165 mZ-g-1. The surface area remains almost unchanged after incorporation of B 3§ Li § and Mg 2+ (5 at.% of ion content), but decreases in the case of Ca 2§ and K + to 145 mE-g-1. Pore size distribution experiments show, in the latter case, a small decrease (from 60 to 55 A) of the average mesoporous diameter.

X-ray T-AI203 has a spinel type structure [9] with a close-packed cubic arrangement of 32 oxygen anions per unit cell and 21-1/3 aluminium cations, inducing 2-2/3 vacant cation positions in the lattice. According to [ 10,11 ] all vacancies are in the octahedral interstices. This leads to a slight deformation of the cubic structure to a tetragonal symmetry. As a consequence ~/-A1203 (Fig.5) shows very broad, overlapping and unresolved diffraction lines. To calculate the unit cell parameters the positions of the diffraction lines were determined by fitting the observed peaks with a Split-Pearson function using the programme FIT [ 12].To reduce the number of fit parameters a cubic lattice symmetry is assumed and the small tetragonal distortion of the lattice is neglected. Fig.5a shows that no additional peaks appear besides the y-AI203 peaks in the spectra of modified alumina catalysts at 5 at.% content of basic elements. Nevertheless, the average cell parameter _a of the lattice increases for Li and Mg modified catalysts as shown below in Table 2. This is not observed in the case of Ca and K modified samples. Table 2. The averaged cell parameters _a of modified catalysts. Catalyst

y-AI203

5 at.% Li §

5 at.% Mg 2§

5 at.% K §

a [/~]

7.907

7.917

7.920

7.910

5 at.% Ca 2+ 7.901

Table 2 indicates that Mg 2§ and Li § cations interact with 'y-AI2O3 lattice taking the tetrahedral free cation positions, or alternatively replacing A13§in octahedral lattice positions respectively. This leads to normal MgAI204, or partially inverse LiAlsOs spinel formation. Fig.5 B) presents the spectra of Mg2+-modified aluminas with different Mg 2+ content in the samples. The lattice parameter _a, calculated for 6 cubic peaks ( hkl indexes: 220, 311, 222, 400, 511, 440 ) of different MgZ+-modified aluminas is presented in Fig.6. As it is seen from this figure the average cell parameter changes with increasing Mg 2§ concentration, approaching the value for regular MgAIzO4 spinel 9a = 8.0831 A [13]. When Ca z§ and K § are added to ~/-A1203 the a parameter does not change. No new additional peaks appear in the spectra for cation concentrations up to 20 at.%. This shows that K § and Ca 2+ cations do not diffuse into the alumina bulk.

177

A)

,=o

)=,,,=(

I0

"

~

~o " ~ o "

"

~o"

~"

~o"

~o " ~o " I~o'

11o

2O M g 2+ ( n % ) / A 1 2 0 3

B)

n=,o :i

IN .,.,,

,.o .z-, r162

9

40

I

I

"

"

I

48

44

"

I

52

"

56

I

"

60

I

"

I

64

"

72

68

2O

~9

"~ ~o

"

_Ca2§ (30%)/A1203

:~

"

~

"

~

"

~

'

do

"

~

"

~

"

~

20

Figure 5. X-ray diffraction patterns of the different catalysts

'

~m

178

8.10 ~..--.o--------____.____

8.05 ,-,o-"

o< ~a,)

o

,.,o---"

8.00 7.957.90-

.o

--o---__~..._

7.85

7.80

S

84

--V-----I"1--

7.75 7.70

Mg 2§ content [atomic-%] Figure 6. The lattice parameter a for Mg2+-modified aluminas from different hkl index peaks. According to the first Pauling rule the cation/anion radii ratio, allowed for octahedral coordination (f~t) has to be in the range of 0.41 - 0.73. For tetrahedral co-ordination this ratio (fro) has to be within the range of 0.22-0.41. To estimate the above ratios for the cations used a value of 1.4 A is taken for the O 2- anion radius value [ 14]. Results are presented in the Tab.3. Table 3. Cation radii and c a t i o n / 0 2. ratio Cation type

Radius [/~] [14] tetrahedral

K§

cation / 02. ratio

octahedral

f~t

1.38

fo~ 0.99

Li §

0.59

0.74

0.42

0.53

Mg z§

0.49

0.72

0.35

0.51

Ca 2§ A13+

1.00 0.39

0.59

0.72 0.28

0.42

Table 3 suggests that AI3§ cation in 7-A1203 structure can be replaced by Mg 2§ both in tetrahedral and octahedral interstices, but Li + can only enter into an octahedral interstices of the lattice. Diffusion for Ca 2§ cation into the 7-A1203 bulk is rather difficult, since fo~t= 0.72 is at the upper limit of the allowed range. For K + cation such a diffusion is impossible because f~t = 0.99 is out of the allowed range. From 30 at.% Ca 2+ content in "/-A1203 new peaks appear in the spectra which can be attributed to Ca~2AI14033 (Mayenite) [ 13] (see Fig.5c ). This mixed calcium-aluminium oxide has a cubic structure with the cell parameter a = 11.982 A.

179 When K is added to ]/-A1203 at a concentration of more than 30 at.% new peaks were also observed in the spectra (Fig.5c) which is attributed to the mixture of different K-oxides [ 13]. Based on this data the process of ?-A1203 modification by basic elements can be describe by the following. When the impregnation technique is used to prepare the catalysts the nitrate salts are decomposed during the calcination step and the basic oxides (CaO, MgO, Li20, K20) are formed and spread onto the y-AI203 surface. The formation of a monolayer or submonolayer of such solids during oxide-oxide interaction has also been reported [ 15]. This process is thermodynamically favourable due to the formation of strong surface bonds between the oxide ions and the alumina surface. MgO and Li20 at 620~ diffuse into the alumina bulk [ 16]. This leads to MgAI204 and LiAlsOs spinels formation. CaO interacts with the 'y-Al203 surface forming the mixed oxide, Ca12A114033 layer. K modification of ?-A1203 results in potassium oxides formation on the alumina surface. Different structures induce different acidbase properties in the materials prepared. TPD

TPD experiments were carded out with methanol in order to study the influence of Mg modification on ?-A1203 surface acidity. It is known [18,19] that CH3OH dissociatively adsorbs on aluminium ions losing its proton and forming a methoxy group. With the increase of surface Lewis acidity the methoxy group adsorption becomes stronger. This stronger adsorption leads to an increase of CH3 - group reactivity, since the CH3 - O bond becomes weaker and is more easy to break. CH3OH desorption proceeds through the dimethylether (DME) formation from the alumina surface. The temperature of DME desorption is a characteristic of CH3 - g r o u p ch~nic, al reactivity and reflects the relative surface acidity. The results for pure ?-A1203 and Mg modified (5 at.%) alumina are presented on Fig.8. The temperature for DME desorption increases, indicating the decreases of Lewis acidity and methyl group reactivity.

~,-alumina

Im .,=.,

_

.'. , . . : " . ^ d -.~'"

".~ x

.,.r

.',,,'...:'....

~. .~. . ~ . . _ . , , . . . ,

"". . 4 ." . ".',.',

'

0

I

100

'

I

200

'

I

300

'

I

400

'

I

500

.{,,~ , . . . ' . . . . 9 .,...

'

I

600

Temperature [ ~ Figure 8 Desorption of DME in the TPD process of methanol: comparison between ?-alumina and Mg 2§ (5%) / y-alumina

180 4. CONCLUSIONS For the studied catechol methylation reaction the catalyst structure and surface properties can explain the catalytic behaviour. As mentioned above, the reaction at 260-350~ has to be performed over the acid catalysts. Porchet et al. [2] have shown, by FTIR experiments, the strong adsorption of catechol on Lewis acid/basic sites of the 7-A1203 surface. These sites control the reaction mechanism. Mg 2+ and Li + cations in the ~/-A1203 structure reduce the acidity of the surface. This causes an increase in the energy needed to activate adsorbed methanol and results in the decrease of the observed catalytic activity. Since the catalyst structure remains the same type, it is possible to control the surface acidity by varying the amount of cation added. The strength and amount of active acid/basic sites on the surface influence also the catechol adsorbed structure as well as the ring polarisation and therefore determine the O/C-alkylation ratio. The best catalyst for 3-methyl catechol selectivity was found to be Mg2+-modified alumina. Addition of 7.5 at.% Mg 2§ to ~/-A1203 was optimal and increased the A3 initial selectivity up to 0.65. Due to the formation of Ca/A1 mixed oxide on the surface, the Ca2+-modified alumina has a completely different structure compared to the spinel one. This leads to a different type of surface Lewis acid/basic sites, rendering the catalyst 30 times less active. Modification of ~'-A1203 by K + cation was shown to form a potassium oxides layer. This layer blocks the alumina surface and decreases drastically (> 99%) the catalytic activity with respect to pure y-AI203. REFERENCES

1. Fiege, H., et al., Phenol Derivatives, Ullmann' s Encyclopaedia of Industrial Chemistry, Vol. A 19, VCH Vedagsgesellschaff, Weinheirn, 1991, p. 313. 2. Porchet, S., Doepper, R., Renken, A., Chem. Eng. Technol., 17 (1994) 108. 3. Durgakumafi V., Narayanan S., J. of Molecular Catal., 65 (1991) 385 (and references cited therein) 4. Tanabe K., Misono M., Ono Y., Hattori H. (eds.), New Solid Acids and Bases (Studies in Surface Science and Catalysis), vol. 51, Elsevier, Amsterdam, 1989, p.231. 5. Velu S., Swamy C.S., Appl. Catal. A: General, 119 (1994) 223. 6. Berteau, P., Ceckiewicz, S., Delmon, B., AppL Catal., 31 (1987) 361. 7. Jurczyk, K., Kania, W., Appl. Catal., 56 (1989) 203. 8. Lercher, J. A., Colombier, C., Noller, H., J. Chem. Soc., Faraday Trans. 1, 80 (1984) 949. 9. Knozinger H., Ratnasamy P., CataLRev.-Sci.Eng., 17(1) (1978) 31. 10. Deyu Li, O'Cormors B.H., Acta Crystallographia, A 46 (1990) C-61. 11. Gray T.J., High Temperature Oxides, Academic Press, New York - London, 1971, p. 77. 12. Siemens Analysentechnische Mitteilung 311 ( N ~ A19100-E689-B7-V-1). 13. ICDD PDF-2 Database, International centre for diffraction data. 14. Shannon R.D., Prewitt C.T., Acta Crystallographia, B 25 (1969) 925 15. Xie Youchang, Gui Linlin, et al., Adsorption and Catalysis on Oxide Surface, Elsevier Sci., Amsterdam, 1985, p. 139. 16. Schmalzried H., Solid State Reactions, Verlag Chemic, Basel, 1981, p. 105. 17. Moulder G.F., et al., Handbook of X-ray Photoelectron Spectroscopy, Perkin-Elmer Corporation, USA, 1992, p.55. 18. Busca G., Rossi P.F., Lorenzelli V., J. Phys. Chem., 89 (1985) 5433 19. Greenler, R.G., J. Phys. Chem., 37 (1962) 2094

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) I l th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

181

SELECTIVE OXIDATION WITH COPPER COMPLEXES INCORPORATED IN MOLECULAR SIEVES Robert Raja and Paul Ratnasamy National Chemical Laboratory Pune - 411 008, India

The selective oxidation, with molecular dioxygen as well as aqueous H202,of n-hexane, naphthalene and phenols using complexes of copper (copper acetate, copper phthalocyanine, copper chlorophthalocyanine and copper nitrophthalocyanine) encapsulated in molecular sieves (X, Y, MCM-22 and VPI-5) is reported. The integrity of the copper complexes in the cavities of the molecular sieves was confirmed by IR, UV-VIS, ESCA and ESR spectroscopic techniques. Dimeric copper acetate complexes encapsulated in molecular sieves oxidise phenols to ortho diphenols and diphenols to quinones using 02 as the oxidant. Encapsulated chloro- and nitro- phthalocyanines oxidise n-hexane selectively to 1-ol and 1-al using molecular oxygen and to 2-ol and -on, 3-ol and -on using aqueous H202 as the oxidants, respectively. Similar oxidation of naphthalene yields beta naphthol with high selectivity using H202 as the oxidant. The copper phthalocyanines, encapsulated in the zeolites Na-X or Na-Y, oxidise phenols to hydroquinone and catechol with high efficiency. Copper complexes encapsulated in molecular sieves are a promising class of selective oxidation catalysts. 1. INTRODUCTION Even though transition metal complexes (such as copper acetate) were used as oxidising agents for organic compounds during the era of the alchemists [1], their use as selective oxidation catalysts is of more recent vintage. The advantages of the encapsulation of such complexes inside the cavities of molecular sieves is also well recognised [2]. Hybrid catalysts of this type (termed variously as 'ship-in-bottle' complexes, "zeozymes", etc.) possess, in principle, the advantages of both heterogeneous (easy catalyst separation, shape selectivity, ruggedness, etc.) and homogeneous (high selectivity, chirality) catalysts. Compared to the wealth of information available on iron complexes encapsulated in zeolites [3], the chemistry of similar copper compounds is known in less detail. This is surprising in view of the major role of copper in catalysing a large number of redox reactions in the biosphere. The synthesis and characterisation of copper (II) perfluorophthalocyanines encapsulated in zeolites Na-X and Na-Y was studied by Balkus Jr. et al [4]. They reported that the Cu(II)/Cu(1) redox process was clearly well-defined when the complexes were entrapped within the zeolite cavities, likely as a consequence oftheir site isolation, while the same process for the dissolved molecules was extremely ill-defined and the electrochemical response (in voltametry) was not very reproducible because of aggregation. Earlier, Ellis and Lyons had reported that replacement of the easily oxidisable hydrogen atoms in the phthalocyanine or porphyrin ligand in transition metal complexes by halogen atoms enhances the oxidative stability and improves the catalytic activity in redox reactions[5]. More recently, Lyons et al. [6], have reported that halogenated porphyrins of Fe, Mn, Cr and Co are catalysts for the selective oxidation of acyclic alkanes with O2. In the present study, we report the synthesis, characterisation and catalytic properties (in selective oxidation reactions) of copper acetate, copper tetradecachlorophthalocyanine and copper tetranitrophthalocyanine encapsulated in molecular sieves Na-X, Na-Y, MCM-22 and VPI-5. Both molecular oxygen and aqueous H202 have been used as the oxidants. The

182 results obtained using n-hexane, naphthalene and phenol (representative of n-paraffins, aromatic hydrocarbons and substituted aromatics, respectively) indicate that copper complexes with appropriately chosen ligands and incorporated in molecular sieves are a novel class of catalysts which can selectively oxidise a wide variety of organic substrates using both 0 2 and H202 as oxidants. 2. E X P E R I M E N T A L 2.1 Materials (1) Solid copper acetate monohydrate (Cu-Ac), and copper acetate in (2) zeolite Na-Y (Si/AI = 2.5; 0.08% wt. Cu; designated as CuAc-Na-Y) (3) zeolite H-Y (Si/AI = 2.5; 0.12 % wt. Cu; designated as CuAc-H-Y) (4) zeolite H-MCM-22 (Si/AI = 60; 0.05% wt. Cu; designated as CuAc-MCM-22) and (5) the aluminophosphate molecular sieve VPI-5 (0.06% wt. Cu, designated as CuAc-VPI-5) were used as catalysts for the hydroxylation of phenol using molecular O2 as the oxidant. CuAc-Na-Y was prepared by stirring 3.5 g of copper acetate and 7 g ofNa-Y (PQ Corporation, USA) in distilled, deionized water for 8 hrs., filtering and washing with distilled water till the washings are free of copper. The above procedure was again repeated with 3.5 g of copper acetate monohydrate. The catalyst was later evacuated (10 a Torr.) and dried at 383 K for 24 hrs. CuAc-H-Y was prepared by first exchanging the Na-Y (twice) with 1 M. ammonium acetate for 8 hours. The ammonium form was calcined at 753 K for 24 hrs. to obtain H-Y. CuAc-H-Y was prepared by the same procedure described earlier for CuAc-Na-Y. CuAc-MCM-22 was prepared by stirring 5.0 g of copper acetate and 9.0 g of H-MCM-22 in distilled, deionized water for 8 hrs., filtering, washing with distilled water till the filtrate is free from copper, evacuation (10 .3 Torr.) and drying at 383 K for 24 hrs. CuAc-VPI-5 was prepared by stirring 4 g of copper acetate and 6.5 g VPI-5 molecular sieve in distilled, deionized water for 8 hrs, filtering and washing with distilled water to remove excess copper acetate. The above procedure was repeated with 4 g of copper acetate. The material was then evacuated and dried at 373 K for 36 hrs. After incorporation of the copper acetate, the solids were first dried in vacuum, then in nitrogen at 298 K and stored in a desiccator. CtlClt4 Pc and Cu(NOz)4 Pc (where Pc stands for phthalocyanine) complexes were obtained from M/s. Lona Industries, Bombay. Details of the procedure for the synthesis of CuCI~4 Pc or Cu(NO2)4Pc encapsulated in zeolite NaX is described below : Aluminium isopropoxide and NaOH (Aldrich) were used without further purification. The silicate gel was prepared from 4.0g of fumed silica (Sigma), 3.2 g NaOH, 0.30 g of CuCI~, Pc or Cu(NO2)4 Pc and 8.0 ml of H20. Addition of the aluminate solution (9.0 g of AI (iOPr)3 , 3.2 g NaOH, 6.0 ml H20) to the silicate gel resulted in a slurry with an intense green/blue colour. An additional 36 ml of deionized water was added. The gel was then transferred to a polypropylene bottle. The mixture with a molar composition of SiO2 : A1203 : Na20 : H20 : CuCI14 Pc/Cu (NO2)4 Pc = 3 : 1 : 3.6 : 141 : 015 was aged at room temperature with stirring for 24 hrs and then heated at 363K for 15 hrs. The mixture was then allowed to cool to room temperature and diluted with copious amounts of deionized water. The solid crystals were isolated by centrifugation at 8000 rpm for 2 hrs. The light, green/blue solid was dried at 363K for 24 hrs in air and extracted (soxhlet) first with acetone, then with pyridine and again with acetone for 72 hrs. It was finally dried at 363 K under vacuum (10 .3 Torr) for 15 h. The X-ray diffraction pattern of the material confirmed that it contains the zeolite, Na-X. Zeolite NaX with varying loadings of CuCI~4Pc or Cu(NO2)4Pc were prepared in a similar way as described above. The synthesis of zeolite Na-Y- encapsulated CuCI~4 Pc and Cu(NO2)4 Pc was done in a similar manner. The silica source was sodium silicate (Lona) and aluminium sulfate (Aldrich) was used as the alumina source. The SiO2/AI203 molar ratio was 2.6. The phthalocyanine based catalysts are designated by the following notation [(Complex) (Zeolite) (copper content in the zeolite, % wt)]. Thus, CuCI~4Pc-Na-Y(0.26) designates a Na-Y zeolite containing 0.26% wt copper in the form of a tetradecachloro copper phthalocyanine complex encapsulated in the supercages of the faujasite structure. Similarly, Cu(NO2)4Pc designates tetra nitro copper phthalocyanine.

183

2.2 Procedures 2.2.1 Catalytic Reaction Reactions with copper acetate-based catalysts have been described earlier [7]. In a typical oxidation reaction, the solid catalyst (0.2 to 0.75g) was added to the substrate in a suitable solvent [.phosphate buffer (0.05 M, pH = 6.5), acetonitrile, etc]. Aqueous H202 (25% wt) was added after the desired temperature was attained. In the case of oxidations with phthalocyanine-based catalysts using O2 as the oxidant, tertiary butyl hydroperoxide (70% aqueous solution, {Aldrich}) equivalent to 1% by weight of the substrate was added to the reaction mixture before air was admitted into the Parr autoclave (300 ml capacity). The catalytic runs were carried out in a three necked flash (100 ml capacity) fitted with a condenser (circulating chilled water) and magnetic stirring. The temperature of the reaction vessel was maintained using an oil bath. Periodically, samples were removed and centriguged to remove the solid catalyst. Copper was not detected (by atomic absorption spectroscopy, Hitachi Model Z-8000) in the colourless reaction product when using any of the solid catalysts used in the present study. The substrates chosen were monohydroxy aromatic compounds like phenol (AR grade, S.D. Fine Chemicals, India), meta and ortho cresols (B.D.H., 15 % wt), n-hexane and naphthalene.

2.2.2 Product analysis The products of the oxidation reaction were analysed by gas chromatography (Hewlett Packard, 5880 A), employing a FID detector and equipped with a capillary column (50 m x 0.25 mm crosslinked methyl silicone gum). The reactants and products of n-hexane oxidation were analysed by gas chromatography (Hewlett Packard, 5890) equipped with a FFAP column (30 m x 0.25 ram). The identity of the products was further confined by GC-MS (Shimadzu QCMC-QP 2000A).

2.3 Catalyst Characterisation The ESR spectra of the solid catalysts were measured at room and liquid N2 temperatures using a Bruker ESR spectrometer (200 D). The second derivative spectra were calculated from the digitised absorption spectra. X-ray diffractograms of the solid catalysts were recorded using a Rigaku D-max III, X-ray diffractometer with a CuKa target. IR and UV spectroscopy of the solid catalysts were recorded using a Perkin Elmer 1600 FTIR and Shimadzu UV-2101 UV-VIS spectrophotometers, respectively. The IR spectra of the solid catalysts were recorded in fluorolube or nujol media (Perkin Elmer). BaSO4 was used as the reference material for recording the diffuse reflectance spectra in the 200-900 nm region. XPS of the solid catalysts were recorded with a VG Scientific ESCA III, Mark (II) spectrometer with MgKcz (1253.6 A,) as the excitation source. 3. RESULTS AND DISCUSSION 3.1 Catalyst characterisation The copper content and chemical composition of the catalysts are given in the Experimental section. The x-ray diffractograms of the catalysts containing the copper complexes did not reveal any significant differences from those of the pure zeolites indicating that the molecular sieves had not undergone any major structural changes due to the incorporation of the copper complexes. The XPS binding energies of copper in the molecular sieves were also similar to the values in the 'neat' complexes. There is, however, a major difference in the XPS spectra of copper ions in CuCI~4Pc and Cu(NOz)4 Pc. While all the copper ions in Cu(NO2)4Pc are in the divalent state, our XPS data indicate that about 25% of the copper ions in CuCI~4 Pc are in the monovalent state (Fig. 1B). In the non-chlorinated

184

A

E

-4(

-20

I i--

-2C

F II

_

935

940

J

935

940

2 0 0 0 1500 2 0 0 0

4000

BINDING ENERGY, eV

I000

cM-I

F~g. 1. ESCA spectra of CuPC (A)and Fig. 2. IR spectra of CuAc (A), Cu-H-Y (B), CuCll4Pc (B) at 298 K Cu-MCM-22 (C), Cu-VPI-5 (D), fluorolube (E), CuCll4Pc (F) and CuCll4Pc-Na-Y {0.26} (G). * is used to denote the nujol peaks.

A

'

B

C

5 I-

3

o

2

200

T

5oo

9oo

WAVELENGTH(nm)

Fig. 3. Diffuse reflectance UV-Vis spectra of catalysts : Curves 1-5 represent Na-Y, CuCi~4Pc, CuCIt4Pc-Na-Y (0.26), Cu(NOz)4Pc and Cu(NO2)4Pc-Na-Y (0.16), respectively.

l,

',

' ...... s~oo ' L ~.soo ~,0o

~.Soo ~5oo 25oo

0E

Fig. 4. ESR spectra at 298 K. A : Curves 1-4 refer to CuAc, Cu-H-Y, Cu-MCM-22 and Cu-VPI-5. B : Curves 1-4 refer to the second derivative spectra of curves 1-4 of A. C : Curves 1-4 represent the absorption curves of CuCit4Pc, CuCil4Pc-Na-Y (0.26), Cu(NO2)4Pc and Cu(NOz)4Pc-Na-Y (0.16), respectively.

185 precursor CuPc, however, there is no evidence for the presence of monovalent copper (Fig. 1A). On treatment of CuCI~4Pc with 02 at 373k for 3 hr 'in situ' in the Esca spectrometer, the peak at 935.8 e.v. (due to Cu § disappeared. The binding energy of Cu 2§ in CuCI14Pc is 937.2 e.v. Its value in the chorine-free CuPc (Fig. 1A) is 936.9 e.v. Fig. 2 shows the infrared spectra of the encapsulated copper acetate and copper chlorophthalocyanine complexes. It is seen that structures of both the complexes are intact even when they are incorporated in the molecular sieves. For example, the bands at 2960 and 2920 cm 1 (due to the asymmetric and symmetric C-H vibrations of the CH3 group ) as well as those at 1630 cm ~ (due to the carboxylate group) are clearly seen in the spectra of encapsulated copper acetate complexes (Fig. 2A). The diffuse reflectance UV-Vis spectra of both the chloro and nitro complexes in the 'neat' as well as the encapsulated states also reveal that the integrity and structure of the complexes are more or less, preserved when they are encapsulated in the zeolites (Fig.3). The absorption maxima for CuCII4Pc-Na-Y (0.26) are at 681.0 and 384 nm, respectively. The corresponding values for CuCI~aPc are 665.5 and 374 nm respectively. Similarily the values for Cu(NOz)4Pc and Cu(NOz)4Pc-Na-Y (0.16) are 628.5, 354 and 652, 373 nm, respectively. There is, hence, a shift to lower energy for these ligand-based electronic transitions (g--->~t') on encapsulation of the complexes inside the zeolite cavities. Balkus et al [3] had also observed a similar phenomena in the case of cobalt (II) and copper (II) perfluorophthalocyanines encapsulated in Na-Y zeolites. They attributed this red shift to the distortion of the phthalocyanine ligand in the supercage of the zeolite. It may be noted that the size of the planar phthalocyanine ligand (- 13 A) exceeds the effective dimensions of the zeolite supercage (about 12.4,). The red shift in the UV-Vis spectra, due to the deformation of the planar phthalocyanine ligands, is indicative of the encapsulation of the copper complex inside the zeolite cavities and not merely adsorbed on the extemal surface of the zeolite. The integrity of the copper complexes is also supported by their ESR parameters (Fig. 4 and Table 1). The second-derivative of their ESRspectra (Fig. 4B) exhibit the seven lines characteristic of a pair of magnetically interacting Cu ~§ ions confirming that the dimeric structure of copper acetate monohydrate is intact even when it is incorporated in the zeolites. A more detailed analysis of the spectroscopic features of these encapsulated complexes will be published later. Table I : ESR parameters (at 298 K) of copper complexes

System CuCll4Pc Cu(N Oz)4 Pc CuCll4 Pc-Na-Y (0.26) Cu(NOz)4 Pc-Na-Y (0.16) CuAc CuAc-H-Y CuAc-MCM-22 CuAc-VPI-5

gt

g~

A~

Ax~

2.08 2.06 2.04 2.03 2.03 2.00 2.02 2.03

2.12 2.16 2.12 2.16 2.21 2.22 2.21 2.22

45 38 51 44 36 43 50 43

61 58 66 62 63 58 55 66

3.2 Catalytic activity 3.2.1 Oxidation with O z

Table 2 illustrates the oxidation of n-hexane with 0 2 (from air) over CuCll4Pc , Cu(NO2)4Pc and the two complexes encapsulated in Na-Y. One interesting finding is that both CuCI~4Pc and its encapsulated analog are able to oxidise even the primary C-H bonds of the end-methyl groups of n-hexane to C-OH groups and the corresponding aldehyde (Table 2). The nitro complex, on the other hand, is unable to oxidise the CH 3 groups. In this context, the halogenated phthalocyanines seem to be more active than their porphyrin analogs in activating primary C-H bonds. Lyons et al [6] found that, while their perfluoro porphyrin complexes were active catalysts for the oxidation of unactivated acyclic alkanes having either

186 secondary or tertiary C-H bonds, the primary C-H bonds were quite resistant to oxidative attack in the presence of their catalysts. It may be noted from Table 2 that though the "neat" complexes are more active than their encapsulated analogs on a weight basis, the catalytic efficiency of the latter (in terms of activity per copper atom) is higher than the former. This is, probably, due to the isolation of the copper sites in the zeolite cavities. An additional feature of the results in Table 2 is that oxidation at the 3-position increases with temperature. Similar results were observed in the oxidation of n-alkanes in Pd-Fe-zeolites by the DuPont workers [8] who attributed it to the end-on diffusion of the linear paraffins in the narrow channels of the molecular sieve. Table 2 : O x i d a t i o n of n-hexane with Oz over phthalocyanines

Catalyst

Temp (K)

CuCI~4Pc CuCI~4Pc-Na-Y (0.26) Cu(NO2)4Pc Cu(NO2)4Pc-NaY (0.16)

n-C6 Conv (%)

Products (%) 1-ol

aid

2-ol

333 343 353 333

14.1 18.2 19.9 8.9

4.9 4.7 2.7 2.8

9.2 9.9 6.4 6.1

. 1.1 2.3 .

353 333 353 333

10.2 6.2 8.1 4.8

4.6 -

5.6 -

. 1.7 1.2 0.7

353

6.9

-

-

1.4

2-on .

.

.

.

1.2 .

2.5 5.4 . .

3-ol

.

3-on 1.9

2.6 3.1 1.2

. 1.2 1.4 0.9

1.7 2.4 2.0

1.6

1.7

2.2

Reaction conditions : n-hexane = 15 g; catalyst = 1.2 g; air = 7 bar; acetonitrile = 15 g; reaction time = 8 hrs. Table 3 : Oxidation of phenols (TON') with O z over copper acetatebased catalysts at 298 K

CATALYST

Cu-Ac CuAc- Na-Y CuAc-H-Y CuAc-MCM-22 CuAc-VPI-5

TON l phenol

o-cresol

m-cresol

L-tyrosine

3.76 19.71 22.10 35.60 20.20

7.20 18.61 56.42 71.5 45.73

5.20 21.62 50.94 55.80 33.80

2.85 12.90 15.15 27.61 10.46

TON" : moles of substrate converted per mole of copper in the catalyst. The oxidation of phenol, ortho/meta cresols and tyrosine with Oz over copper acetate-based catalysts at 298 K is shown in Table 3 [7]. In all the cases, the main product was the ortho hydroxylated diphenol product (and the corresponding orthoquinones). Again, the catalytic efficiency (turnover numbers) of the copper atoms are higher in the encapsulated state compared to that in the "neat" copper acetate. From a linear correlation observed [7] between the concentration of the copper acetate dimers in the molecular sieves (from ESR spectroscopic data) and the conversion of various phenols (Fig. 5), we had postulated [8] that dimeric copper atoms are the active sites in the activation of dioxygen in zeolite catalysts containing encapsulated copper acetate complexes. The high substrate specificity (for mono-

187 or diphenols) and high regioselectivity (oxidation only of the position ortho to the phenolic group) were the distinctive features of this catalyst system, which is a solid mimic, in a restricted sense, of both the monophenolase and diphenolase catalytic activity of the monooxygenase enzyme, tyrosinase [7]. 3.2.2 Oxidation with HzO z Oxidation of n-hexane and naphthalene, with a singlet oxygen source, H202, as the oxidant, is illustrated in Tables 4 and 5, respectively. The most striking difference between the oxidation (of n-hexane) with 02 and H202 (Tables 2 and 4, respectively) is the total absence of n-hexan-l-ol and the n-hexan-l-aldehyde when H202 is used as the oxidant in the case of both CuClz4Pc and CuCI~4Pc-Na-Y (0.26). By contrast, these two compounds were the most important constituents of the products when 02 was the oxidant (Table 2). Another notable feature of Table 4 is the higher concentration of the 3-ol and 3-on at high conversion levels, especially with CuCl~4Pc-based catalysts. There is no significant variation in the ratio of (2+3)-ol / (2+3)-on between the chloro- and nitro complexes as well as at different temperatures. (The value varies between 0.8 and 1.0). This is not surprising since the rate of oxidation of the seconday alcohol to the ketone is expected to be much faster than that of n-hexane to the alcohol. A surprising observation in the oxidation of naphthalene (Table 5) is the preponderance of the beta rather than the alpha isomer among the naphthol products. The latter would have been the predominant species if the attacking moiety had been a simple electrophilic cation (like HO § for example). Oxidaion by ion-radicals is, perhaps, involved. The results of the oxidation of phenol with H202 over CuPc-based catalysts is shown in Tables 6 and 7 as well as in Fig. 6. We have found that both CuClt4Pc and Cu(NO2)4Pc are active catalysts for the oxidation of phenol to hydroquinone/catechol using H202 as the oxidant. The catalytic nature of the reaction is confirmed by the data in Fig. 6A which illustrates the influence of catalyst weight on the conversion of phenol and product distribution. Phenol conversion is negligible in the absence of any catalyst. The kinetics of the oxidation of phenol over CuCI~4Pc-Na-Y (0.26) (Fig. 6B) reveals that, while oxidation occurs at both the ortho and para positions from the beginning of the reaction, parabenzoquinone (rather than hydroquinone) is the predominant para-substituted product in the early stages of the reaction when H202 is present in relatively larger quantities. Table 6 illustrates the influence of copper content (and, hence the concentration of the copper complex) in the zeolite on the rate of oxidation of phenol as well as the product distribution. The following salient features may be noted : (1) the catalytic efficiency (turnover number) of the copper ions is higher in the encapsulated state than in the "neat" complex; (2) the chloro complex is more active than the nitro complex; (3) the turnover number does not vary significantly with copper content suggesting that the copper complexes are well isolated from each other (the conversion of phenol, of course, increases with the concentration of the copper complex in the zeolite); (4) catechol, hydroquinone and parabenzoquinone are the only major products obtained in the oxidation of phenol. The amount of 'tar' products is less than 1%. The product was almost colourless (the product distribution in Table 6 is given on a 'tar free' basis); and (5) parabenzoquinone is observed amongst the products with zeolites containing relatively larger amounts of encapsulated CuCl~4Pc complexes (with 0.27 and 0.28 % wt of Cu, respectively). For the maximum utilization of H202 in the conversion of phenol to (hydroquinone + catechol) and minimum formation of quinones, the phenol : H202 ratio must be kept as high as possible (Table 7). In the case of CuClt4Pc-Na-Y, the ortho-para ratio [catechol/(hydroquinone + parabenzoquinone) ratio in Table 3] is lower at very high concentrations of H20 z (phenol/H202 = 1). This behaviour is different from that observed in the case of titanosilicates. Thangaraj et al [9], had observed that, in the case of oxidation of phenol with H202 over TS-1 molecular sieves, the relative concentration of catechol was higher at higher concentrations of H202. For example, at phenol : H202 mole ratios of 3 and

188

A

Fig.5. Correlation between the integrated area of the seven line ESR spectrum (of Cu-H-Y, Cu-MCM-22 and Cu-VPI-5) and conversion of L-tyrosine, phenol, o-cresol and m-cresol (Curves A-D, respectively).

z o u

Y

MCM-22

I..

.

I

I0

.... I

VPI-5 I

-

2O

ARE.A E.SR ( ARB UNITS)

Table 4 : Oxidation of n-hexane with Catalyst

Temp (K)

HzOzover phthalocyanines n-C6

Products (%)

Conv (%)

CuCll4Pc CuCII4Pc-Na-Y (0.26) Cu(NOz)4Pc Cu(NO2),Pc-Na-Y (0.16)

2-ol

2-on

3-ol

3-on

333 343 353 333

19.2 22.5 26.2 16.5

6.5 9.2 7.1 6.1

8.9 11.0 8.3 7.2

2.1 1.7 5.9 1.1

1.7 0.6 4.9 2.1

343 353 333 353 333

19.1 22.4 10.5 11.2 6.3

6.8 6.2 1.9 2.2 1.1

9.1 7.6 3.1 4.6 1.7

1.6 4.1 3.2 1.7 2.0

1.6 4.5 2.3 2.7 1.5

353

8.2

1.2

1.9

2.9

2.2

Reaction condition 9n-hexane = 15 g; n-hexane/H202 = 3 mole; acetonitrile = 15 g; reaction time = 8 h.

Table 5 : Oxidation of naphthalene with HzO z over phthalocyanines Products (%)

Catalyst

Temp (K')

naphth Conv (%) A

B

C

D

CuCII4Pc

353 363 373 373

3.1 5.6 7.2 6.8

1.8 2.1 4.2 3.8

0.7 1.8 1.7 2.0

0.3 0.8 0.5 0.7

0.2 0.9 0.6 0.2

373

7.1

4.6

1.1

0.9

0.3

CuCI~4Pc-Na-Y (0.26) CuCI~4Pc-Na-X (0.28)

Reaction conditions : naphthalene = 15 g; naphth/H202 = 3 mole; acetonitrile = 15 g; reaction time = 8 h; A = beta naphthol; B = alpha naphthol; C = naphthaquinones; D = phthalic anhydride.

189 Table 6 : Oxidation of phenol with HzO z over phthalocyanines

Catalyst

Cu Cu Cu Cu Cu Cu Cu Cu

Cu (% wt)

C114Pc (NO2)4 Pc CI14Pc-Na-Y (0.II) CI14Pc-Na-Y (0.17) Cll4Pc-Na-Y (0.26) Cll4 Pc-Na-Y (0.27) Cll4Pc-Na-X (0.14) Cll4PC-Na-X (0.28)

Cu(NO2) 4 Pc-Na-Y Cu(NO2), Pc-Na-Y Cu(NO2)4 Pc-Na-X Cu(NO2)4 Pc-Na-X

0.11 0.17, 0.26 0.27 0.14 0.28 0.09 0.16 0.11 0.14

(0.09) (0.16) (0.11) (0.14)

TON

2.1 1.21 7.95 8.27 6.77 7.35 8.25 7.36 4.70 3.95 4.25 3.95

Products (% wt) PBQ

CAT

HQ

2.0 12.5 0 0 0 6.1 0 11.2 0 0 0 0

59.5 0 65.8 63.1 50.9 49.1 62.8 46.4 0 0 0 0

38.5 87.5 34.2 36.9 49.1 44.8 37.2 43.4 100 100 100 100

Reaction conditions : Catalyst wt = 0.75 g; phenol : H 2 0 2 = 3:1, mole; reaction time - 8 h; Temp = 353 K; solvent = acetonitrile; PBQ = para benzoquinone; CAT - Catechol; HQ hydroquinone; TON -- moles of phenol converted per mole of copper in the catalyst Table 7 : Oxidation of phenol over CuCI~4 Pc-Na-Y (0.26) ; Influence of HzOz concentration

Phenol/H202 moles

Phenol Cony. %

1 2 3 4 5 10

Products (% wt)

20.9 21.4 19.7 17.6 15.1 7.9

PBQ

CAT

HQ

16.4 2.9 0 0 0 0

40.9 48.5 50.9 51.2 53.6 52.8

42.7 48.6 49.1 48.8 46.4 47.2

Reaction conditions : Catalyst wt = 0.75 g; phenol : H 2 0 2 = 3:1, mole; reaction time = 8 h; Temp = 353 K; solvent = acetonitrile; PBQ = para benzoquinone; CAT - Catechol; HQ = hydroquinone. CONV

A

tB

CONoV

5

5 O

I

2

5

CATALYST WT. (g)

i,Ir 0

~

5

~ o.I0

REACTION TIME, hrs

Fig. 6. Influence of catalyst weight and residence time on phenol conversion (CONV), and formation of catechol (CAT), hydroquinone (HQ) and para-benzoquinone (PBQ) over CuCi~4Pc-Na-Y (0.26) at 353 K; see Table 6 for reaction conditions.

190 1, the ratios of catechol to (hydroquinone + parabenzoquinone) were 0.94 and 1.56, respectively. The corresponding values for CuCIt,Pc-Na-Y are 1.03 and 0.69, respectively (Table 6). The differences in the shape selectivity and relative hydrophobicity/hydrophilicity of TS-1 and Na-Y are probably additional factors that play a role. In addition, the fast conversion of hydroquinone to parabenzoquinone at high concentrations of H202 enhances the formation of the para-substituted product. It may be recalled that the "neat" chloro complex, CuCI~4Pc, produces 1.5 times more catechol than hydroquinone (Table 6) in accordance with expectations based on statistical probability. The higher, relative, concentration of hydroquinone observed over our zeolite-encapsulated chloro complexes is probably due to the geometric constraints imposed by the zeolite matrix. 4. CONCLUSIONS We have shown that copper complexes (like acetates and phthalocyanines) encapsulated in molecular sieves are novel and versatile catalysts for a variety of selective oxidation reactions. Both dioxygen and singlet oxygen sources like H202, can be used as oxidants. A variety of substrates including paraffinic and aromatic hydrocarbons as well as substituted aromatics like phenols can be oxidised in a selective manner. An unique feature of halogenated copper phthalocyanines is that they are able to oxidise the C-H bonds of even the primary carbon atoms, including those in the -CH3groups in normal paraffins, using O2 as the oxidant. Copper acetate dimers encapsulated in molecular sieves are regioselective, solid, aromatic hydroxylation catalysts, which mimic, in a restricted sense, the catalytic activity and specificity of the monooxygenase enzyme, tyrosinase. One distinguishing and unique feature of all the copper-based catalyst systems studied in this paper is that the catalyst remains in the solid phase during the entire course of the reaction and can be easily filtered off after the reaction is over, thereby providing significant processing advantages in their large scale appplication. The details of the mechanism of activation of dioxygen by these catalysts and the course of the oxidation reaction are not clear and are yet to be elucidated. 5. A C K N O W L E D G E M E N T We thank the European Commission (Contract No. Cll-CT93-0361) for financial support. The experimental help ofS. Badrinarayanan and R.F. Shinde in collecting the ESCA and ESR spectra, respectively, is gratefully acknowledged. REFERENCES

5(a) (b) 6. 7. 8. 9

E. Vogel, Schweigger's Journal, 13 (1815) 162. N. Herron, G.D. Stucky and C.A. Tolman, J. Chem. Soc., Chem. Commun., (1986) 1521. F.T. Starzyk, R.F. Parton and P.A. Jacobs, Stud. Surf. Sci. Catal., 84 (1994) 1419. K.J. Balkus Jr., A.G. Gabrielov, S.L. Bell, F. Bedioui, L. Rone and J. Devyuck, Inorg. Chem., 33 (1994) 67. P.E. Ellis Jr. and J.E. Lyons, Coord. Chem. Rev., 105 (1990) 181. J.E. Lyons and P.E. Ellis Jr., Appl. Catal., A :Gen., 84 (1992) L1. J.E. Lyons, P.E. Ellis Jr. and H.K. Myers Jr., J. Catal., 155 (1995) 59. R. Robert and P. Ratnasamy, J. Mol. Catal., 100 (1995) D.R. Corbin and N. Herron, J. Moi. Catal., 86 (1994) 343. A. Thangaraj, R. Kumar and P. Ratnasamy, J. Catal., 131 (1991) 294. |0

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis -40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

191

Discovering the Role of Au and KOAc in the Catalysis of Vinyl Acetate Synthesis William D. Provine, Patrick L. Mills, and Jan J. Lerou DuPont Central Research & Development Experimental Station, Wilmington, Delaware USA 19880-0262 1. I N T R O D U C T I O N The commercial process for the production of vinyl acetate monomer (VAM) has evolved over the years. In the 1930s, Wacker developed a process based upon the gas-phase conversion of acetylene and acetic acid over a zinc acetate carbon-supported catalyst. This chemistry and process eventually gave way in the late 1960s to a more economically favorable gas-phase conversion of ethylene and acetic acid over a palladium-based silica-supported catalyst. Today, most of the world's vinyl acetate is derived from the ethylene-based process. The end uses of vinyl acetate are diverse and range from the protective laminate film used in automotive safety glass to polymer-based paints and adhesives. The chemistry of vinyl acetate synthesis from the gas-phase oxidative coupling of acetic acid with ethylene has been shown to be facilitated by many co-catalysts. Since the inception of the ethylenebased homogeneous liquid-phase process by Moiseev et al. (1960), the active catalytic species in both the liquid and gas-phase process has always been seen to be some form of palladium acetate [Nakamura et al, 1971; Augustine and Blitz, 1993]. Many co-catalysts which help to enhance the productivity or selectivity of the catalyst have appeared in the literature over the years. The most notable promoters being gold (Au) [Sennewald et al., 1971; Bissot, 1977], cadmium acetate (Cd(OAc) 2) [Hoechst, 1967], and potassium acetate (KOAc) [Sennewald et al., 1971; Bissot, 1977]. A vinyl acetate catalyst technology that has been extensively patented is a palladium-gold (Pd-Au) bimetallic silica-supported catalyst promoted with potassium acetate (KOAc). Patent claims and journal literature have shown that addition of Au to a Pd catalyst improves the intrinsic selectivity of the catalyst to vinyl acetate [Bartley et al., 1993] along with boosting its production rate [Sennewald et al., 1971 ]. KOAc has been shown to boost the vinyl acetate production rate of the catalyst [Debellefontaine and Besombes-Vailhe, 1978; Samonos et al., 1971] and promote combustion of acetic acid [Nakamura and Yasui, 1980]. Most, if not all, of the previous research describe the overall reactor performance effect upon adding promoters to the catalyst mix, but none addressed the fundamental questions of "how the chemistry changes" with addition of such promoters. The primary purpose of the work described herein was to search for some experimental evidence on how Au and KOAc manipulate the underlying mechanism of vinyl acetate synthesis when these components are added to a palladium-based catalyst. 2. E X P E R I M E N T A L M E T H O D O L O G Y & CATALYST SYNTHESIS In order to probe the influence of Au and KOAc on the vinyl acetate synthesis chemistry, four different catalysts were synthesized. All of these catalysts were prepared in a manner exemplified in prior patent technology [Bissot, 1977], and each contained the same palladium loading in an egg-shell layer on the surface of a spherical silica support. The palladium content in the catalyst was easily controlled by adjusting the solution strength of palladium chloride (PdCI2) added to the porous silica beads prior to its precipitation onto the support by reaction with sodium metasilicate (Na2SiO3). The other two catalyst components (Au and KOAc) were either present or absent in order to complete the independent evaluation of their effect on the process chemistry, e.g., (1) Pd+Au+KOAc, (2) Pd+KOAc, (3) Pd+Au, and (4) Pd only.

192

3. REACTOR EVALUATION RESULTS AND DISCUSSION The chemistry of vinyl acetate synthesis was studied using both a laboratory-scale high pressure fixed bed reactor along with a TAP (Temporal Analysis of Products) reactor [Gleaves et al., 1995] operated at atmospheric pressure (Figure 1). The fixed bed reactor was used to determine productivity results at high pressure, while the TAP reactor provided evidence on the transformation of the vinyl acetate chemistry upon addition of gold (Au) and potassium acetate (KOAc) to a palladium (Pd) catalyst. Each of the four catalysts were tested under a variety of conditions with both 1,2 ~SC-ethylene and unlabelled ethylene. The labeled ethylene was used to allow segregation and discrimination of competitive pathways, e.g. CO 2 formation from ethylene or acetic acid while the unlabeled ethylene was used for the high pressure productivity evaluations.

3.1 Fixed Bed Reactor Experiments Each catalyst was evaluated in the fixed bed reactor their relative performance via the metrics of VAM STY (Space-Time-Yield) and VAM SEL as derived from eqs. 1 and 2. [=] g/liter-hour [=] %

VAM STY = 86*Pb*rv~ VAM SEL = 100*rv~,~/(rvAu+0.5*rco2)

(1) (2)

In eqs. 1 and 2, r v ~ and rco2 denote the molar hourly production of each species per gram of catalyst while Pb denotes the catalyst's bulk density (grams catalyst per reactor liter). Each of the catalysts from the Pd-Au-KOAc catalyst series was tested and evaluated under steady-state conditions in the fixed bed reactor under process conditions typical of vinyl acetate synthesis and the VAM STY and VAM SEL results are included as Table 1.

Table 1 Fixed bed reactor performance of Pd-Au-KOAc catalyst series after 40 hours on-stream. Test Conditions: 165 ~ 115 psig, with a feed stream consisting of C2H4, AcOH, 02, and N 2.

Catalyst

VAM STY (g/iite r-hr)

VAM SEL

Pd-Au w/KOAc Pd w/KOAc Pd-Au Pd

764 100 594 124

93.6 95.4 91.6 94.7

,, ,,,i,ll it

( %)

,,,,,,,,,,, ,,,,,,,,,,,,,,, ,

These results show that important performance benefits are obtained by the addition of Au and KOAc to a palladium catalyst. Gold (Au), when added to palladium (Pd), enhanced the VAM production rate of the catalyst (VAM STY) substantially while it decreased the overall selectivity of the catalyst. This is true in both cases: Pd-Au w/KOAc vs. Pd w/KOAc (764 vs. 100 & 93.6% vs. 95.4%) and Pd-Au vs. Pd (594 vs. 124 & 91.6% vs. 94.7%). Conversely, KOAc increased the selectivity of the catalyst whether or not Au is used. In the most appetizing example (Pd-Au w/KOAc vs. Pd-Au), KOAc improved the selectivity of the catalyst by 2.0% while it additionally improved the production rate of the catalyst 30%. KOAc only increased the production rate of the catalyst in the presence of Au. KOAc decreased the VAM production rate of the catalyst on its addition to the Pd catalyst where Au was absent. In summary, the results from the fixed bed reactor study provided evidence as to the effect of Au and KOAc on the performance of the catalyst, though, these experiments did not give any information on the perturbation of the reaction pathways with the addition of Au and KOAc. For this type of information, additional experiments were performed using the TAP reactor with 1,2 C-labeled ethylene used as an isotopic tracer of the kinetics. 13

193

Figure 1. Key elements of the TAP reactor (A) and high pressure fixed bed reactor 03) experimental systems. The TAP reactor schematic shows the heated valve manifold and reactor with the elevated pressure attachment located in the main high vacuum chamber. The fixed bed reactor shows the feed system, liquid vaporizer, oxygen disperser, reactor, and waste recovery system.

194

3.2. TAP Reactor Experiments A TAP reactor allows transient response experiments to be performed under vacuum or atmospheric pressures. Figure 2 shows typical response curves when pulses of ethylene, oxygen, and helium are introduced at atmospheric pressure into the TAP reactor via a pump-probe experiment. Ethylene and oxygen are pulsed, from separate reservoirs, directly into a 5 cm~ stream of nitrogen saturated with acetic acid at 60~ Pump-probe time interval (time between ethylene and oxygen pulses) was chosen to be 0.01 second and the delay between sets of ethylene and oxygen pulses was 15-30 seconds depending on the washout time for the particular component under study. Helium is used as an inert tracer and is pre-mixed with both the ethylene and oxygen pulses and enters the feed stream with these pulses. The feed stream is then passed through a 0.45 cm ~ (0.25 g) bed of crushed catalyst particles (250-425 ~tm) which were held at 170~ The reactor effluent exits the system either through a vent line or through a small capillary bleed line to the mass spectrometry chamber where the response curves were detected. The vinyl acetate evolved as a product of the oxidative coupling of the ethylene with the acetic acid and it exited the reactor after the ethylene and oxygen. The TAP reactor's pulse sequences allowed each catalyst to be evaluated with expensive labeled reactants since typical pulse intensities were approximately 1 torr of gas per pulse. Results from the TAP reactor results were obtained by signal averaging l0 pump-probe ethylene + oxygen pulses that are then smoothed and adjusted to zero baseline by subtraction of the background. The 0.01 second pump-probe time interval was determined to be sufficiently short as to behave as if the pulses of ethylene and oxygen entered the reactor together. Their segregation also enabled easy manipulation of the ethylene to oxygen ratio entering the reactor without pre-mixing large volumes of gases Figure 2. Reactant & product response curves for a pump-probe which would be explosive, pulse sequence in the TAP reactor. One of the challenges to correctly interpreting the TAP reactor data in this particular application is the discrimination between the mass spectra data for the various components. The assignments used here for the results included in this paper are listed in Table 2.

Table 2. Mass fragments used in identification of components described in this article. Reactants used in these experiments were 13C2H4, CH3COOH , 02, and He. a.m.u. 4 12 44 45 60 74 75 88

Compounds (Intensity) He (1) CO (s) CO 2(l), CH3OOCCH 3 (s) 13CO2(l), CH3COOH (1), 13CH3OOCCH3 (l) CH3COOH (1) CH3OOCCH 3 (m) 13CH3OOCCH, (m) CH3COO13CH13CH2 (s)

195

The first set of results from the TAP reactor, as shown in Figure 3, shows the 1,2 ~3C vinyl acetate (MW=88) response curve for the four catalyst samples. The transient response suggests that KOAc dramatically accelerated the desorption of the vinyl acetate off the surface of the catalyst (peak maximum at 7.5 seconds without KOAc, 5.5 seconds with KOAc on average). In addition, Au enhanced the desorption rate of the vinyl acetate, but to a much lesser extent. It is also seen that Au improved the production rate Figure 3. Transient response of 1,2 ~3C-vinyl acetate (MW=88) of the catalyst (peak areas: obtained from a typical pump-probe experiment with the Pd-Au > Pd and Pd-Au palladium-based catalysts and 1,2 ~3C-ethylene. w/KOAc > Pd w/KOAc). The total amount of vinyl acetate produced per pulse agreed with the VAM STY data obtained from the fixed bed reactor with the exception that the Pd-Au catalyst produced more vinyl acetate in the TAP reactor than the Pd-Au w/KOAc catalyst. This can be attributed to the Pd-Au catalyst's VAM production rate being limited by the desorption of VAM when operated at elevated pressures and with a constant flow of C2H 4 and 02. Several long term multipulse experiments were used to help to bridge the gap between the steadystate/high surface coverage fixed bed reactor work included in Table 1 and the transientflow surface coverage pump-probe TAP reactor experiments of Figure 3. The long term multipulse experiments were conducted with a mixture of ethylene, d4-acetic acid, and oxygen characteristic of VAM synthesis conditions. This mixture was pulsed into a 9 cm3/min stream of nitrogen at an intensity of 2.6 torr/pulse of gas. Each experiment consisted of 200 pulses spaced 1 second apart from one another. As can be seen from Figure 4, the Pd-Au w/KOAc catalyst produced the most vinyl after 200 pulses in direct agreement with the fixed bed reactor results. It is also seen that vinyl acetate production is limited at the early stages of this experiment when KOAc is absent from the catalyst mix. Again, this can be attributed to the slow desorption rate of VAM when KOAc is not present. Figure 4. d3-vinyl acetate (MW=89) response curves derived from a long term multipulse sequence with a premixed d3-acetic acid + ethylene + oxygen feed.

196 Figure 5 shows VAM evolution curves which depict the desorption rate of VAM as a function of acetic acid concentrations. These results were obtained when unlabelled ethylene and oxygen are pulsed into a carrier gas that contains various amounts of acetic acid. As the concentration of acetic acid in the feed stream is increased, the desorption of vinyl acetate is enhanced. Thus, the role of KOAc seems to be linked with its ability to keep acetic acid on the catalyst surface. This Figure 5. Vinyl acetate transient response curve (MW=86) for result complements nicely the obtained from a typical pump-probe experiment with the hypothesis of Tamura and palladium-based catalysts and unlabeled ethylene. Results were Yasui (1979), who theorized obtained with the Pd-Au w/KOAc catalyst. that KOAc forms double salts (e.g., KH(OAc)2) with the acetic acid on the surface of the catalyst. Moreover, these complexes would form a molten salt layer on the surface of the catalyst which would set-up an environment referred to as supported liquidphase catalysis. Carbon dioxide can form as a result of combustion of ethylene, acetic acid, or vinyl acetate. The 1,2 t3Cethylene experirnents provided evidence on the role of Au and KOAc and its influence over the formation of carbon dioxide in the process, t2CO2 would be a characteristic of acetate decomposition, while 13CO2 would be characteristic of ethylene decomposition. Figure 6 compares the 12CO2 response curves obtained from each of the catalysts used. The maximum in the ~CO 2 response curve is attributed to the combustion Figure 6. 12CO2transient response curves (MW--44) obtained of acetic acid, while the from a typical pump-probe experiment with the palladium-based extended tail of the Pd catalysts and 1,2 ~3C-ethylene. w/KOAc and Pd curves is attributed to a secondary combustion pathway. The secondary acetate decomposition pathway is most likely attributed to the conversion of carbon monoxide to carbon dioxide.

197 Figure 7 compares the transient responses for evolution for riCO obtained from a typical TAP pumpprobe experiment with the four catalyst samples. 12CO can be directly attributed to the combustion of acetic acid. The combustion of acetic acid (CH3COOH) is believed to evolve carbon dioxide via the carboxylic (COOH) fragment of the molecule. Then, the remaining carbon (CH 3 group) is converted to carbon monoxide at low oxygen coverages and onward to Figure 7. nCOtransient response curves (MW=12) obtained carbon dioxide at higher from a typical pump-probe experiment with the palladium-based oxygen coverages [Davis and catalysts and 1,2 ~3C-ethylene. Barteau, 1991; Aas and Bowker, 1993]. This is supported by the TAP reactor experiments since the maximum of the riCO response curve occurs between 5.5-6.0 seconds, while the CO 2 evolution occurs earlier between 4.55.0 seconds. It is also seen that both KOAc and Au enhance the formation of CO. This agrees with the earlier observations with the CO 2 response curves, and supports the hypothesis that the secondary CO 2 peak occurs from the conversion of CO to CO 2. The results from Figures 6 and 7 support the observation that acetic acid combustion is accelerated by the presence of Au and KOAc. The evolution of carbon dioxide is enhanced by both Au and KOAc, while the evolution of carbon monoxide is enhanced by the presence of KOAc and suppressed by the presence of Au. This shows that acetic acid combustion is more complete with a Pd-Au alloy versus Pd alone which is important since carbon monoxide can act as a temporary catalyst poison in the process. These results agree with Nakamura and Yasui's (1980) on acetic acid oxidation which showed an increase in acetic acid combustion when KOAc is added to a Pd catalyst. Figure 8 shows the t3CO2 (MW--45) response curve for the standard pump-probe experimental conditions. ~3CO2would be generated from either the combustion of 1,2 13C2H 4 or the product 1,2 ~3C-vinyl acetate. Since the curve has a maximum at 4.5 seconds which is before the evolution of vinyl acetate at 5.5-7.5 seconds, ethylene combustion seems to be the pathway that can be attributed to the formation of this peak. Moreover, on interpretation of Figure 7, it is apparent that both the Pd-Au w/KOAc and the Pd w/KOAc catalysts produced less 13CO 2 than their Pd-Au and Pd counterparts. Therefore, it can be concluded that KOAc impeded the combustion of ethylene.

198

Figure 8. ~3CO2 transient response curves (MW--45) obtained from a typical pump-probe experiment with the palladium-based catalysts and 1,2 ~3C-ethylene. Figure 9 shows the results from experiments with the Pd-Au w/KOAc catalyst and the dependence of acetic acid and ethylene combustion on surface oxygen coverage. Figure 9a shows the ~3CO2 evolution curve, while Figure 9b shows the ~2CO2 evolution curve. The experiments were conducted with the same size ~3C2H4 pulse intensity while the 02 pulse intensity was increased so that the C2H 4 to 0 2 molar ratios were varied between 2 to 1 and 4 to 1. It shows that increased oxygen concentration enhances both the combustion of acetic acid and ethylene. As can be seen from the figure, ethylene combustion is virtually eliminated at low oxygen coverages. This suggests that the main route to carbon dioxide formation in the synthesis of vinyl acetate is the combustion of acetic acid. This agrees with previous researchers [Davidson et al., 1984; Crawthome et al., 1994] who had observed similar behavior on other palladium-based catalysts.

Figure 9. ~3CO2(MW--45) and ~2CO2 (MW--44) response curves for a ethylene + oxygen pulse to a continuous acetic acid saturated nitrogen stream after it is passed over the Pd-Au w/KOAc catalyst with different ethylene to oxygen ratios (e.g., 4 to 1).

199

4.

CONCLUSIONS AND SUMMARY

The effect of the catalyst composition upon the catalyst activity, selectivity, and reaction pathways was examined using a conventional high pressure fixed bed reactor and a TA_P reactor. Particular emphasis was placed upon the effect of Au and KOAc on the acceleration or impedance of the pathways associated with vinyl acetate synthesis. A summary of the key findings is given below: (1) KOAc enhanced the combustion of acetic acid to carbon dioxide and carbon monoxide (2) KOAc suppressed the combustion of ethylene to carbon dioxide (3) KOAc and Au enhanced the desorption rate of vinyl acetate (4) Au and KOAc enhanced the formation of vinyl acetate (5) Au suppressed the formation of carbon monoxide from the combustion of acetic acid (6) Au enhanced the conversion of carbon monoxide to carbon dioxide On conclusion of this experimental work, key insight was learned as to the role of Au and KOAc in the synthesis of vinyl acetate. Additionally, this work helped to identify how Au and KOAc physically perturbed the chemistry. For example, Neurock (unpublished results) has shown by density functional theory calculations that the binding energy of CO and other adsorbates are higher on Pd-Au clusters than on Pd clusters alone. This supports the hypothesis that Au has an electronic stabilization effect upon Pd complexes in the system which increased the catalyst's ability to turnover these moieties to vinyl acetate. On the other hand, Crathorne et al. (1994) have shown that acetic acid exits in high concentrations near the surface of Pd-KOAc catalysts. It is believed that KOAc created a molten salt layer with the acetic acid on the surface of the catalyst which promoted high Pd-OAc surface coverages. This multi-layer coverage provided a protective coating on the surface of the catalyst which impeded ethylene combustion by isolating discrete palladium sites for preferential adsorption of ethylene and oxygen that leads to VAM formation rather than combustion. The formation of vinyl acetate is viewed to occur either on small clusters of palladium acetate dissolved in the "supported liquid phase" or on palladium acetate dense surfaces. The overall picture of the role of a Pd-Au-KOAc-SiO 2 catalyst in the synthesis of vinyl acetate is included as Figure 10.

Figure 10. Role of Pd-Au-KOAc-SiO 2catalyst in the manufacturing of vinyl acetate

200 5.

ACKNOWLEDGMENTS

The authors would like to acknowledge the contributions of several individuals for their insight and hard work in achieving the data included in this report: J. Scott McCracken (TAP reactor), Kevin S. Slusser (fixed bed reactor), and Tom Borecki (catalysis synthesis). The authors would also like to thank DuPont's vinyl acetate business and manufacturing teams for allowing this work to be published. 6. REFERENCES N. Aas and M. Bowker, J. Chem. Soc. Faraday Trans., 89 (1993), 1249. S.M. Augustine, and J.P. Blitz, J. Catal., 142 (1993) 312. W.J. Bartley, S. Jobson, G.G. Harkreader, M. Kitson, and M. Lemanski, U.S. Patent 5,274,181 (1993). T.C. Bissot, U.S. Patent 4,048,096 (1977). E.A. Crathorne, D. MacGowan, S.R. Morris, and A.P. Rawlinson, J. Catal., 149 (1994) 254. J.M. Davidson, P.C. Mitchell, N.S. Raghavan, Front. Chem. React. Eng. (Proc. - Int. Chem. React. Eng. Conf.), 1 (1984), 300. J.L. Davis and M.A. Barteau, Surf. Sci., 256 (1991) 50. H. Debellefontaine, and J. Besombes-Vailhe, J. Chim. Phys., 75 (1978) 801. Farbwerke Hoechst Artiengesellschaft, Great Britain Patent, 1,188,737 (1967). J.T. Gleaves, J.R. Ebner, and P.L. Mills, paper presented at USPC-2, St. Louis, Mo., Sept. 1995. T. Kunugi, K. Fujimoto, H. Arai, T. Kono, and A. Namatame, Kogyo Kagaku Zasshi, 71 (1968) 2007. I.I. Moiseev, M.N. Vargaftik, J.K. Syrkin, and Aka Doklady, Nauk SSR, 133 (1960), 377. M. Nakamura, Y. Fujiwara, and T. Yasui, U.S. Patent 4,087,622 (1978). S. Nakamura, and T. Yasui, J. Catal., 17 (1970) 366. S. Nakamura, and T. Yasui, J. Catal., 23 (1971) 315. S. Nakamura, and T. Yasui, J. Japan Pet. Inst., 23 (1980), 416. M. Neurock (unpublished results). B. Samanos, P.Boutry, and R. Montarnal, J. Catal., 23 (1971) 19. Sennewald, U.S. Patent 3,631,079 (1971). W. Schwerdtel, Chem. Ind., (1968), 1559. M. Tamura, and T. Yasui, Shokubai, 21 (1979), 54. M.N. Vargaftik, V.P. Zagorodnikov, and I.I. Moiseev, Kinet. Katal, 22 (1981), 743. P. Wirtz, K. Woemer, F. Wunder, K. Eichler, G. Roscher, and I.Nicolau, Canadian Patent Application 2,071,699 (1992).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 1 lth International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

201

F o r m a t i o n of c i t r a c o n i c a n h y d r i d e b y v a p o r - p h a s e d e c a r b o x y c o n d e n s a t i o n of p y r u v i c acid 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 Indstrial Ltd., 1978 Kogushi, Ube 755, Japan

Citraconic anhydride (Methyl maleic anhydride) was found to be produced from pyruvic acid by an oxidative decarboxy-condensation. The best catalyst is iron phosphate with a P/Fe atomic ratio of 1.2. The presence of oxygen is required to promote the reaction. The main side-reaction is formation of acetic acid and CO2 by oxidative C-C bond fission. The best results are obtained at a temperature of 200~ The yield of citraconic anhydride reaches 71 mol% at a pyruvic acid conversion of 98%. 1. INTRODUCTION Pyruvic acid is the simplest homologue of the a-keto acid, whose established procedures for synthesis are the dehydrative decarboxylation of tartaric acid and the hydrolysis of acetyl cyanide. On the other hand, vapor-phase contact oxidation of alkyl lactates to corresponding alkyl pyruvates using V205- and MoO3-baseds mixed oxide catalysts has also been known [1-4]. Recently we found that pyruvic acid is obtained directly from a vapor-phase oxidativedehydrogenation of lactic acid over iron phosphate catalysts with a P/Fe atomic ratio of 1.2 at a temperature around 230~ [5]. CH3-CH(OH)-COOH + 0.502 ; CH3-CO-COOH + H20 The one-pass yield of pyruvic acid reached 50 mol%. The main side reactions are the formation of acetaldehyde and CO2 by oxydative C-C bond fission of lactic acid and that of acetic acid and CO2 by oxidative C-C bond fission of the produced pyruvic acid. CH3-CH(OH)-COOH + 0.502 ; CH3CHO + CO2 + H20 CH3-CO-COOH + 0.502 , CH3COOH + CO2 Over the V205- and MoO3-based mixed oxide catalysts, the main part of lactic acid is converted to form acetaldehyde and CO2.

202 In the reaction of lactic acid to form pyruvic acid over the iron phosphate catalysts, formation of a new c o m p o u n d was observed. As the extent of reaction increased, the amount of pyruvic acid increased to a m a x i m u m and then decreased, while that of the new compound increased steadily. It was therefore concluded that the new compound is formed from pyruvic acid in parallel with acetic acid and CO2. According to gas-mass analyses, the molecular weight was determined as 112. However, there are many c o m p o u n d s with molecular weigth of 112. After the NMR analyses and X-ray diffraction analyses for the single crystal, the new compound was determined to be citraconic anhydride, i.e., mono-methyl maleic anhydride. 2 CH3-CO-COOH + 0.502

; CH3-C-C--O \ + CO2 + 2 H 2 0

II .,o

HC-C=O Citraconic anhydride and citraconic acid are used as raw materials of various chemicals, resines, surface-active agents, and dyes. as is maleic anhydride. They have generally been produced from itaconic acid which is produced by fermentation of cane sugar and grape sugar. Therefore, the price is much higher than that of maleic anhydride. Indeed, by analogy with the formation of maleic anhydride from n-butene, there have been attempts to produce citraconic anhydride by vapor-phase contact oxidation of olefinic hydrocarbons with carbon number of more than five, such as isoprene, 2-methyl butene, toluene, and xylene, using V205-based mixed oxide catalysts. However, the yields were clearly lower than that oberved in the oxidation of n-butene. Citraconic anhydride formation from pyruvic acid by oxidative decarboxycondensation has not been k n o w n prior to these studies. Therefore, in this paper, we attempted to get more insight into the new reaction.

2. EXPERIMENTAL An iron phosphate catalyst with a P/Fe atomic ratio of 1.2 used in this study was prepared according to the procedures described in the previous studies [6-8]. On the other hand, a V-P oxide catalyst with a P/V atomic ratio of 1.06 and pumice supported 12-molybdophosphoric acid (H3PMo12040) and its cesium salt (Cs2HPMo12040) catalysts were the same as used in a previous study [9]. Pumice supported WO3-based mixed oxide catalysts were the same as used in a previous study [101. The contact oxidation of pyruvic acid was carried out with a continuous-flow system. 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. Air or a mixture of nitrogen and oxygen was fed in from the top of the reactor and an aqueous solution containing 100 g of pyruvic acid in 1000 ml was introduced into the preheating section of the reactor by means of a syringe pump. The feed rates of pyruvic acid, oxygen, nitrogen, and water vapor are either 10.5, 70, 280, and 480 or 10.5, 13.4,

203 350, and 480 m m o l / h , respectively, unless otherwise indicated The reaction t e m p e r a t u r e was in the rage of 180 to 270~ The extent of reaction was varied by changing the amount of catalyst used from 1 to 36 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 recovered solution (about 50 ml) and the effluent gas were analyzed by four GC's.

3. RESULTS 3. 1 Performance of various oxide catalysts Since no information has been reported on the reaction, it seems necessary to make a character sketch of the catalytic function. Therefore, various kinds of metal oxides were tested as the catalysts. The results are s u m m a r i z e d in Table 1. It is clear that the V-P oxide and Mo-P heteropoly compound catalysts are not effective for the production of citraconic anhydride. The acidic catalysts such as Table 1 Performance of oxide catalysts for production of citraconic anhydride

Catalyst (Atomic ratio) P/V

(1/1.06)

H3PMo12040 Cs2HPMo12040

Pyruvic acid conversion

Citraconic anhyd. yield (mol%)

(g)

(%)

3 5 10 20 20

23 72 16 77 31

5 5 0 4 6

38 54 77 89 52

0 4 6 0 0

P/Si AI/Si

(1/9) (15/85)

K/Si

(5/95)

20 5 10 20 1

P/Ni P/Fe

(1/1) (1.2/1)

5 10

30 86

2 40

W P/W Mo/W Ti/W Ti / W Sn/W K/W

(5/95) (5/95) (10/90) (70 / 30) (10/90) (10/90)

20 20 20 20 20 20 20

92 25 83 93 59 92 93

61 20 57 60 36 61 52

Reaction temperature = 230~ = 10.5, 350, and 480 mmol / h.

Feed rates of pyruvic acid, oxygen, air, and water

204 Si-P and Si-A1 are not effective, either. The basic catalyst such as Si-K is very active for decomposition of pyruvic acid, but citraconic a n h y d r i d e is not produced. The best results are obtained with W and W-based catalyst. The onepass yield of citraconic anhydride reaches about 60 mol% at a pyruvic acid conversion of about 92%. The combination of P to W decreases the activity markedly. The combination of Mo or K decreases the selectivity to citraconic anhydride. The combination of Ti or Sn is scarcely effective, w h e n the a m o u n t is less than 10 atomic %. When the a m o u n t of Sn or Ti is high, the selectivity falls. It should also be noted that the next best results are obtained with the iron phosphate catalyst u n d e r the reaction conditions used. 3.2. Performace of WO3-based oxide catalysts Since the best results were obtained with the W and W-based oxide catalysts, the reaction was studied in more detail using 20 g portions of these catalysts. The reaction was performed at 230~ with feed rates of pyruvic acid, air, and water = 10.5, 350, and 480 m m o l / h . The contact time defined as volume of catalyst (ml)/rate of gaseous feed (ml/s) was about 5.2 s. The main products were citraconic anhydride and CO2. The amount of acetic acid was very small. No other products were detected except for very small amounts of CO, acetone, and acetaldehyde. A relatively large discrepancy was observed between the amount of consumed pyruvic acid and that of the sum of produced citraconic a n h y d r i d e and acetic acid. This discrepancy was defined as "loss". The product distributions are shown in Figure 1 as a function of the time-onstream. At the begining of reaction, that is in the first 1 h on stream, the conversion of pyruvic acid reached 92% and the yields of citraconic anhydride and acetic acid were 61 and 5 mol%, respectively. The amount of loss was about 26 mol%. It should be noted that the catalytic activity falls markedly with an increase in the time-on-stream. Similar falls in catalytic activity were also observed in the cases of the other W-based mixed oxides, such as W-Sn, W-Ti, and W-Mo. It was also found that the deactivated catalysts can easily be regenerated by a heattreatment at 500~ in air for I h.

3. 3. Performance of iron phosphate catalyst The reaction was studied using the iron phosphate catalyst at 230~ with feed rates of pyruvic acid, air, and water = 10.5, 350, and 480 m m o l / h . The main products were citraconic anhydride, acetic acid, and CO2. W h e n the a m o u n t of catalyst used was 10g, that is, when the contact time is about 2.6 s, the conversion of pyruvic acid reached 95% and the yields of citraconic anhydride and acetic acid were 50 and 28 mol%, respectively; the loss was about 17 mol%. The selectivity to citraconic anhydride is clearly lower and that to acetic acid is higher than in the case of the W-based oxide catalysts. However, the catalytic activity was very stable. No clear change in the yield of citraconic anhydride was observed during the reaction for 10 h. The reaction was then performed using different amounts of catalyst from I to 20g. The yields of citraconic anhydride and acetic acid and the loss are plotted as a

205 100 Catalyst = 20g T = 230~ 80

Conversion

Citraconic Anhyd.

30

Citraconic

0

anhyd.

- T = 230oc

40

x,

60

50

/

\

~ 4O

20

1

AcOH /

~

>, Loss

20

T

0

0

AcOH

1

10! NA.~ A |

-

2

3

7.-!

4

Time-on-stream / h Figure 1. Performance of WO3 catalyst

I I 0 -! 40 60 80 100 Conversion of pyruvic acid / %

Figure 2. Performance of iron phosohate catalyst

function of the conversion of pyruvic acid in Figure 2. The slopes of lines from the origin indicate the selectivities to each product. The selectivities to citraconic anhydride, acetic acid, and loss are about 51, 29, and 20 mol%, respectively. It should also be noted that the selectivities remain unchanged with a large variation in the extent of reaction. This indicates that the citraconic a n h y d r i d e and acetic acid produced are stable enough under the reaction conditions used. 3. 4. Effects of oxygen concentration on the reaction o v e r i r o n phosphate catalyst The reaction was performed over the iron phosphate catalyst by changing the feed rate of oxygen from zero to 350 m m o l / h , while fixing the sum of feed rates of oxygen and nitrogen at 350 m m o l / h . The feed rate of pyruvic acid was fixed at 10.5 m m o l / h . The yields of citraconic anhydride obtained at a temperature of 230~ and a short contact time of 0.52 s (amount of catalyst used = 2 g) are plotted as a function of the feed rate of oxygen in Figure 3. The formation of citraconic a n h y d r i d e increases with an increase in the feed rate of oxygen up to about 70 m m o l / h (air). However, with a further increase in oxygen feed rate, the formation of citraconic anhydride levels off. It is clear that the presence of oxygen is required to form citraconic anhydride from pyruvic acid. Another series of tests were performed in a low oxygen concentration: the feed rates of pyruvic anhydride, oxygen, nitrogen, and water were 10.5, 13.4, 350, 480 m m o l / h , respectively. The extent of reaction was varied by changing the a m o u n t

206

~9,

"~ 9

20

50

02 feed (mmol/h) 13. O , 70. o 9

40

15

o~.

~

../ ~

O// . /

t"

//~itraconic Anhyd.

31-

_

6 I"

-

o

-~

t( ~

0 fj

I~

20 40 60 o.,,~ oxygen feed / m m o l / h

Figure 3. Effect of oxygen on the rate

o

Aco8 '

'

-

,

40 60 80 100 Conversion of pyruvic acid / % Figure 4. Effect of oxygen on the selectivity

of catalyst used. The yields of citraconic anhydride and acetic acid are compared with those obtained with a oxygen feed rate of 70 m m o l / h by using air as the oxygen source (Figure 4). It is clear that the selectivity to citraconic anhydride increases and that to acetic acid decreases with a decrease in the oxygen concentration. 3. 5. Effects of reaction temperature on the reaction over iron phosphate catalyst The reaction was performed over the iron phosphate catalyst by changing the reaction temperature from 180 to 270~ and the amount of catalyst used from 1 to 36 g, while fixing the feed rates of pyruvic anhydride, oxygen, nitrogen, and water at 10.5, 13.4, 350, 480 m m o l / h , respectively. The yields of citraconic anhydride obtained at different temperatures are plotted as a function of the conversion of pyruvic acid in Figure 5. And the yields of acetic acid at different temperatures are also plotted in Figure 6. The selectivity to citraconic anhydride decreases and that to acetic acid increases as the temperature is raised. The results indicate that the activation energy for the formation of citraconic a n h y d r i d e is much lower than that for the formation of acetic acid. The selectivity to acetic acid decreases steadily with a lowering of the temperature. However, the highest selectivity to citraconic anhydride is obtained at 200~ Possibly vaporization of pyruvic acid may become difficult at temperatures below 200~ The yield of citraconic anhydride reached 71 mol% and that of acetic acid was 7 mol% at the pyruvic acid conversion of 98%; the loss was about 20 mol%.

207

!

9o*c

70

50F

-

- 70

/& 40

-

"0 20 "~ 10

o

20

~-/ , ,

10

O[~

~ ~ ! I 20 40 60 80 100 Conversion of pyruvic acid / %

Figure 5. Effect of temperature on the yield of citraconic anhydride

20

0 40 60 80 100 Conversion of pyruvic acid / % Figure 6. Effect of temperature on the yield of acetic acid

4. DISCUSSION Since formation of citraconic anhydride from pyruvic acid is one of "acid to acid type" transformations, such as reactions from isobutyric acid to methacrylic acid and from lactic acid to pyruvic acid, the required catalysts must be acidic [11]. If the catalysts are basic, it may be impossible to obtained acidic products, because basic catalysts activate selectively acidic molecules and, as a result, they show a very high activity for the decomposition of acidic products [11]. As mentioned before, two reactions of pyruvic acid take place in parallel as follows: 2 CH3-CO-COOH + 0.502 2 CH3-CO-COOH + 02

-

~ citraconic anhydride + CO2 + 2 H 2 0 ~ 2 acetic acid + 2 CO2

WO3-based oxide catalysts show a high selectivity to form citraconic anhydride even at a relatively high temperature of 230~ This means that their activity for the oxidative C-C bond fission is completely suppressed. However, they lose quickly the catalytic activity. Possibly, they cannot generate enough of the oxygen species required for the reaction, because WO3 possesses redox function in a very poor extent unlike MOO3, V205, and iron phosphate. As seen in Figure 3, the presence of oxygen is required to promote the formation of citraconic anhydride.

208 Similarly, acidic oxides without redox function, such as Si-A1, Si-P, and Ni-P oxides, are not effective as catalysts for this reaction. MOO3- and V205- based oxides possess both acidic and redox functions. Therefore, they show a good performance as catalysts in many partial oxidations for producing especially acidic compounds. However, they possess double bond oxygen species, that is, M=O species, which promote the oxygen insertion reactions as well as the dehydrogenation. Therefore, in the reaction of lactic acid to pyruvic acid, they promote preferentially the C-C bond fission by oxygen insertion to form acetaldehyde and CO2 [5]. Similarly, in the reaction of pyruvic acid, they promote preferentially the formation of acetic acid and CO2 by the C-C bond fission rather than the condensation to form citraconic anhydride. On the other hand, iron phosphate possesses both acidic and redox functions, though the functions are considerd to be much lower than those of MoO3 and V205. However, it possesses no double bond oxygen species unlike MoO3 and V205. Therefore, its function to promote oxygen insertion processes is very weak. As a result, in the reaction of pyruvic acid, the formation of acetic acid and CO2 by the C-C bond fission is suppressed, to a certain extent, and a relatively good performance is obtained for production of citraconic anhydride. The results shown in Figure 4 indicate that the oxygen dependency of the acetic acid formation is higher than that of the citraconic anhydride formation. Therefore, use of a low oxygen concentration is beneficial to the selectivity to citraconic anhydride, though it is disadvantageous to the reaction rate. It should also be noted that the iron phosphate catalyst is inactive for consecutive decomposition of the produced citraconic anhydride and acetic acid. As may be seen in Figures 5 and 6, when iron phosphate is used as the catalyst, the selectivity is dependent largely on the reaction temperature. Therefore, the activation energy for the citraconic anhydride formation is considered to be much lower than that for the acetic acid formation. Indeed, the selectivity to acetic acid decreases steadily with a decrease in the temperature. However, the selectivity to citraconic anhydride shows a maximum at about 200~ Possibly, the vaporization of pyruvic acid may become difficult at temperatures below 200~ As for the reaction path from pyruvic acid to citraconic anhydride, it is considered that a condensation reaction first takes place by a reaction between an oxygen atom of carbonyl group and two hydrogn atoms of methyl group in another molecule, followed by oxidative decarboxylation to form citraconic acid. The produced citraconic acid is dehydrated under the reaction conditions used. The proposed reaction path is shown in Figure 7.

209

H3C,"

H3C

#O

o..C-C.oH

>

I,-I .,(3 H-C-ClJ C"OH FI 0

H3C\c- C~""O H + II C:_.OH

H,C-

"~0

zO

"C- C~.IOH

H"C-C..O il C"OH

-I- H20

0.5 02

-

>

0

COz

H3C,, C - C ~ 0 II /"~ /O +

H~C- C % 0

H20

Figure 7. Estimated reaction path to citraconic anhydride

REFERENCES

1. T. Yokoyama and K. Matsuoka, Jpn. Patent No. 56-19 854 (1981). 2. Y. Yamaguchi, Jpn. Patent No. 57-24 336 (1982). 3. S. Sugiyama, N. Shigemoto, N. Masaoka, S. Suetoh, H. Kawami, H. Kawami, K. Miyaura, K. Miyaura, and H. Hayashi, Bull. Chem. Soc. Jpn., 66 (1993) 1542. 4. H. Hayashi, N. Shigemoto, S. Sugiyama, N. Masaoka, and K. Saitoh, Catal. Lett., 67 (1994) 273. 5. M. Ai and K. Ohdan, Chem. Lett., (1995) 405. 6. M. Ai, E. Muneyama, A. Kunishige, and K. Ohdan, Bull. Chem. Soc. Jpn., 67 (1994) 551. 7. E. Muneyama, A. Kunishige, K. Ohdan, and M. Ai, J. Mol. Catal., 89 (1994) 371. 8. E. Muneyama, A. Kunishige, K. Ohdan, and M. Ai, Appl. Catal., 116 (1994) 165. 9. M. Ai, J. Catal., 89 (1984) 413. 10. M. Ai, in Proc. 10th International Congress on Catalysis, Budapest 1992, eds.. L. Guci, F. Solymosi, and P. T6t6nyi, Akademiai Kiadpo, Budapest, 1993, p. 1199. 11. M. Ai, in Proc. 7th International Congress on Catalysis, Tokyo 1980, eds., T. Seiyama and K. Tanabe, Elsevier, Amsterdam, 1981, p. 1060.

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J.W. Hightower,W.N. Delgass,E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. I01 9 1996ElsevierScienceB.V. All rights reserved.

211

Heterogeneous enantioselective dehydration of butan-2-ol Saskia Feast, Donald Bethell, Philip C. Bulman Page, M.Rafiq H.Siddiqui, David J. Willock, Graham J. Hutchings, Frank King ~, and Colin H. Rochester b Leverhulme Centre for Innovative Catalysis, Robert Robinson Laboratories, Department of Chemistry Liverpool University, P.O. Box 147, Liverpool, L69 3BX, U.K. "ICI Katalco, Research and Technology Group, P.O.Box 1, Billingham, Cleveland, TS23 1LB, U.K. b Department of Chemistry, University of Dundee, Dundee, DD 1 4HN, U.K.

ABSTRACT: Zeolite Y modified with chiral sulfoxides has been found catalytically to dehydrate racemic butan-2-ol enantioselectively depending on the chiral modifier used. Zeolite Y modified with R-1,3-dithiane-1-oxide shows a higher selectivity towards conversion of S-butan-2-ol and the zeolite modified with S-2-phenyl-3-dithiane-l-oxide reacts preferentially with Rbutan-2-ol. Zeolite Y modified with dithiane oxide demonstrates a significantly higher catalytic activity when compared to the unmodified zeolite. Computational simulations are described and a model for the catalytic site is discussed.

Introduction:

The activity of molecules used as pharmaceuticals, agrochemicals and food additives can depend strongly upon stereochemistry. Accordingly the synthesis of pure enantiomers is becoming increasingly important in the development of new products. Although enantioselective catalysis in homogeneous systems is fairly well established ~2. The problems associated with these systems are the drawbacks normally associated with homogeneous systems i.e. product recovery and catalyst separation. Recently some progress has been made using heterogenised homogeneous catalysis by immobilising homogenous systems 3. Some success has also been reported by the modification of zeolites as catalyst for heterogeneous enantioselective

212 catalysis in the liquid phase *~. In this presentation we describe a method by which zeolite structures can be made chiral by the inclusion within them of a chiral species. Specifically we discuss new catalytic systems consisting of zeolite Y chirally modified by 1,3-dithiane-l-oxide with R=H and R=phenyl ( Fig. 1). With these modified zeolites we have observed the enantioselective dehydration of butan-2-ol in the gas phase and in particular, we demonstrate the use of the combination of experimental studies and computational simulation to address a complex catalytic process.

Experimental: MATERIALS AND REACTION SYSTEM: The method for the synthesis of zeolite Y and the incorporation of the chira] modifier R- 1,3-dithiane- 1oxide Fig. 1 has been reported 7. A ~~ similar method was used to modify s . y s . oSvS zeolite Y with S-phenyl- 1,3-dithianeR 1-oxide. Catalytic experiments were I II carried out in a conventional glass microreactor using on-line GC analysis. The activity of the catalyst (zeolite Y modified with 1,3-dithiane-l-oxide) and the comparison with unmodified zeolite Y, zeolite Y modified with 1,3-dithiane, silica/alumina modified with 1,3- dithiane oxide, and boron nitride modified with 1,3- dithiane-l-oxide was carried out using the following conditions. For a typical catalyic run, catalyst (0.3 g) was reacted with prevaporized butan-2-ol (3.3 x 10 ~ mol h "1) using nitrogen (3.7 x 10 .2 tool h 1) as a diluent. The enantioselectivity reactions were tested using the following conditions. The catalyst (zeolite Y modified with enantiomerically enriched dithiane oxides) 0.1 g, was reacted in a glass microreactor with prevaporized racemic butan-2-ol (7.35 x 10 ~ mol h 1) diluted in nitrogen (6.7 x 10 ~ mol hl). The products were analysed using an on-line GC with a 40m capillary T-cyclodextrin column with trifluoroacetyl stationary phase. The GC was temperature programmed from 25 - 70 ~ with a split ratio of 120:1 to achieve the separation of butan-2-ol enantiomers. TECHNIQUES: X- ray diffTaction (XRD) was carried out using a Hilton Brooks modified Philips 1050W diffractometer with a Cu Ka source (40 keV and 20 mA). Thermogravimetric analysis (TGA) was carried out using a Perkin Elmer TGA 7 analyser. All solid state MAS NMR spectral data were aquired on a Bruker MSL 400 MHz spectrometer. The 13C solution NMR spectra were recorded on a B ~ e r AMX 400 spectrometer and referenced to TMS. B IOSYM molecular modelling was carried out on a Silicon Graphics Indigo 2 work station, using Monte Carlo docking and energy minimisation Figure 1

§

/

213 progromme Discover to visualise the interaction of the sulfoxides within the zeolite fromework.

R E S U L T S AND D I S C U S S I O N : Characterization by X-ray diffraction demonstrated that the modifier did not affect the crystallinity of the zeolite, and studies using 1~C MAS NMR spectroscopy showed that the dithiane oxide had been molecularly adsorbed into the zeolite. The dithiane oxide was remarkably stable inside the acidic zeolite: it did not undergo acid catalysed elimination, hydrolysis or rearrangement. Heating the modified zeolite in flowing nitrogen at 180 ~ did not lead to any decomposition, and the dithiane oxide could be recovered intact in high yield by solvent extraction. Thermogravimetric analysis showed that the dithiane oxide modifier either decomposed or desobed between 500-600 ~ Of particular importance is that such t r e a t m e n t of enantiomerically enriched R-dithiane oxide did not result in racemization. Decomposition of compound I (R = H) adsorbed on the zeolite was only observed at temperatures well in excess of 400 ~ The initial experiments were carried out using zeolite Y modified with racemic 1,3-dithiane-l-oxide, and this was compared as a catalyst for the dehydration of racemic butan-2-ol with a control sample of zeolite Y prepared by an analogous method but without the modifier. The results (Table 1) demonstrate that the dithiane oxide modified catalyst was considerably more active than the control sample by several orders of magnitude, and t h a t this activity was maintained over several days of testing without significant loss of the enhanced activity. Under the reaction conditions studied, the control somple only become active at temperatures above 150 ~ and required a reaction temperature of 225 ~ to achieve 90% conversion, whereas the modified zeolite gave 90% conversion at 115 ~ For the modified zeolite, the distribution of butene isomers is close to that expected for equilibrium; 8 subsequent experiments revealed that the modified zeolite was also an active catalyst for butene isomerization. To confirm that the rate enhancement observed was due to an interaction between the modified zeolite and the substrate, a n u m b e r of control experiments were carried out (Table 1). Zeolite Y was modified using 1,3-dithiane II, and, under the same conditions used for the sulfoxide-modified zeolite, no rate enhancement was observed, indicating t h a t the sulfoxide oxygen atom is a vital feature of this catalyst system. Investigation of a non microporous silica/alumina catalyst t h a t had a Si/A] ratio identical to that of the zeolite Y used in this study indicated t h a t modification by the dithiane oxide acted as a poison. The microporous aluminosilicate framework is therefore also important. In addition, 1,3-dithiane-l-oxide supported on the

214 i n e r t m a t e r i a l boron nitride was found to be totally inactive for butan-2-ol dehydration, i n d i c a t i n g t h a t it is the combination of the d i t h i a n e oxide w i t h the zeolite w h i c h is essential.

Table 1 Reaction of racemic butan-2-ol over modified catalysts Catalyst

Y"

TempfC Conversion/% Selectivity/% but- 1-ene E-but-2-ene Z:but:2:ene

115 0 -

Y-SO b

Y-S ~

SiO/

AI,0~'

BN-SO f

AI~.O.~ ~

SiOJ

225 90

115 90

110 3.6

175 35.6

200 31.5

200 24.5

115/225 0

17.7 40.6

8.3 53.8

15.4 53.8

14.7 46.5

16.5 31.7

21.6 33.9

-

................... : ........... 4 1 . 6

...... 3 7 . 8

........... . 3 . 0 . 7 ......... 3 8 . 8

.......... 5 1 = 8 .............. 4 6 . . 5 ......................... =..............

"Zeolite Y ( u l t r a s t a b i l i s e d LZY 82, Union Carbide) b Zeolite Y modified w i t h (+/-)-l,3-dithiane-l-oxide, 1 molecule per s u p e r cage (7.6 wt%) ~Zeolite Y modified w i t h 1,3-dithiane, 1 molecule per supercage (7.2 wt%) d Non microporous silica alumina, SiO]A1203 =5.7 ' Non microporous silica alumina, SiO2/A1203 =5.7, modified w i t h 1,3-dithiane1-oxide (6.9 wt%). f Boron nitride modified with 1,3-dithiane-l-oxide (7.4 wt%). Boron nitride alone also showed no activity. In a f u r t h e r set of e x p e r i m e n t s racemic butan-2-ol w a s reacted w i t h zeolite Y modified w i t h chiral sulfoxides. In the first set of e x p e r i m e n t s racemic butan2-ol was r e a c t e d over zeolite Y modified by R-1,3-dithiane-l-oxide (83% ee) a n d the results s h o w n in Table 2 indicate t h a t S- butan-2-ol reacts preferentially. In a second set of experiments, the s~me reaction was i n v e s t i g a t e d for zeolite Y modified w i t h S-2-phenyl-l,3-dithiane-l-oxide (99% ee) 9 . The results, shown in Table 2, confirm t h a t the s~me effect is observed, however, in this case the R- butan-2-ol reacts in preference to the S- enantiomer. To u n d e r s t a n d the origin of this high catalytic activity and the enantioselectivity a series of computer simulation studies was carried out to d e t e r m i n e the m o s t favourable locations for 1,3-dithiane-l-oxide (R=H and

215 R=phenyl) w i t h i n th e zeolite framework. The Biosym docking p a c k a g e x~w a s used to g e n e r a t e a r a n d o m s et of ten energetically favourable configurations

Table 2 a Reactivity of r a c e m i c butan-2-ol over zeolite Y modified by homochiral d i t h i a n e oxides. Modifierb

Temp.

Convc.

Butan-2-ol conv./x 10~ tool h "1

~ % {R} R=H 110 0.5 {R} R=H 120 1.3 {R} R=H 150 9.9 {S} R=phenyl 110 4.2 {S} R=phenyl 120 7.5 note : {R} and {S} refer to the chiral centre at the

R S 0.002 0.035 0.006 0.094 0.018 0.707 0.276 0.015 0.415 0.081 sulfoxide sulfur.

Relative rate" 17.5 15.7 39.3 18.4 5.1

a Reaction conditions: 0.1 g of the zeolite Y modified catalyst, tested in a conventional glass microreactor with racemic butan-2-ol (7.35 x 10 .3 tool h-l), prevaporized in a nitrogen diluent (6.2 -6.7 x 10 .3 tool h-l). Products were analyzed u si n g on-line GC w i t h a 40m capillary y- cyclodextrin column with trifluoroacetyl s t a t i o n a r y phase, t e m p e r a t u r e p r o g r a m m e d from 25-70 ~ with a split ratio of 120:1. bZeolite Y modified w i t h R- 1,3-dithiane- 1-oxide/S-2-phenyl- 1,3-dithianeoxide ~ conversion of R- an d S-butan-2-ol d Relative ratio of p r e f e r e n t i a l enantiomeric reaction of R- and S-butan-2-ol for each of the d i t h i a n e oxide molecules in a host zeolite Y s t r u c t u r e w i t h a Si:A1 ratio of 1:1. E a c h s t r u c t u r e was energy- minimized a n d the lowest energy conformation t a k e n as the most likely d i t h i a n e position in the zeolite. Within this s t r u c t u r e , the f r a m e w o r k proton n e a r e s t to the sulfoxide oxygen was t r a n s f e r r e d to form the hydroxy sulfonium cation in the case of R=H and the R - h y d r o x y - l , 3 - d i t h i a n e for R=phenyl; the cation position w a s t h e n reoptimized (Figures 2 and 3). D u r i n g this process the atomic charges for the adsorbed molecules were t a k e n from a Mulliken a n a l y s i s of a DZP basis set H a r t r e e Fock calculation (using CADPAC Xlon the MOPAC ( P M 3 ) o p t i m i z e d molecular structure. Table 3 gives the cation c h a r g es for the sulfoxide group and C2 for each cation. In both cases the sulfur and t r a n s f e r r e d proton are positively charged but at sulfur the

216 addition of the phenyl group at the two position has reduced the charge by almost 0. le.

Fig. 2 Calculated low energy conformation of the protonated dithiane oxide cation (R=H) in zeolite Y (Si/Al = 1). The bottom view shows a view through the twelve ring containing the deprotonated framework oxygen, the top view is perpendicular to this. For clarity the zeolite framework is shown using a stick model and the adsorbed molecule is drawn in space filled form represented by the Van der Waals radii for the atoms being in the order S>O>C>H.

When comparing the relaxed structures of figures 2 and 3, it should be noted that the phenyl group has two effects: it greatly increases the size of the dithiane molecule, leading to greater steric hindrance, and it reduces the effective charge at the sulphoxide sulfur, lowering the electrostatic interaction between the sulfoxide group and the deprotonated framework oxygen. These effects are reflected in the relaxed geometries; in both cases the two closest molecule/framework contacts occur between the S and H atoms of the sulfoxide cation and the deprotonated framework oxygen, the S..O distances being 3.01/k in the phenyl case compared to 2.84A for R=H. For the 1,3-dithiane-1-oxide (R=H) case molecular dynamics simulations at the experimental temperature revealed that the R-hydroxysulfomum cation was considerably more stable than the more weakly adsorbed 1,3-dithiane molecule TM. We consider that the hydroxydithiane cation may act as a proton transfer agent and this may account for the enhanced reactivity of this system.

217

Fig. 3 Calculated low energy conformation of the protonated dithiane oxide cation (R=phenyl) in zeolite Y (Si/A1 = 1). The views and presentational details are as for figure 2.

Table 3 Cation charges for the sulfoxide group and C2 for each cation.

C2 (inter sulfur C) $3 (S of SOH) 0 (O of SOH) H (H of SOH)

R=H charge (le l)

R=phenyl charge (le [)

-0.504 +0.825 -0.611 +0.393

-0.513 +0.727 -0.624 +0.404

The structure shown in Figure 2 was used as the host for f u r t h e r docking calculations to introduce each of the enantiomers of butan-2-ol; the docked structures were energy minimized and the lowest energy structures selected as the most likely butan-2-ol binding positions; these are shown in Figure 4. These structures involve the framework and the R-hydroxysulfonium cation acting in concert to bind the butan-2-ol molecules, and calculations indicate that the S-enantiomer is bound more tightly t h a n the R form ( 16.4 k J mo1-1) and we consider that this may play some role in the observed enantioselective reactions with these catalyst systems.

218 F i g u r e 4 Calculated low energy conformations for enantiomers of butan-2-ol in dithiane oxide (R=H) loaded zeolite Y (Si/A1 = 1). Each of the two butan-2-ol enantiomers was docked into the zeolite Y/dithiane oxide cation system (see Figure 2), and the lowest energy structures are shown here. Binding energies for butan-2-ol: S form, -64.7 kJmol-1; R form,-48.3 kJmo1-1. Two views of each structure are shown, with most of the zeolite framework cut away for clarity; the bridging oxygen atom which donated the proton is highlighted as a small sphere. Throughout this work the zeolite potentials and atomic charges were taken from the consistent force field library provided by Biosym Technologies 13'14. The potentials for the guest molecules were also taken from this library but with charges assigned from an SCF calculation on the isolated molecule performed with a DZP basis set within the CADPAC package TM Taken together, these results provide the first example of a gas phase enantioselective reaction heterogeneously catalysed by a zeolite. Although the initial approach has been to design a catalyst that consumes chiral molecules, we have demonstrated that a catalyst can be designed that is capable of discriminating between enantiomers and preferentially transforming one of them. This important effect is achieved by enantioselective rate enhancement, i.e. both enantiomers react faster in the chiral environment than in the absence of the chiral modifier, but one reacts faster than the other. In view of the vast range of microporous materials and potential enantiomerically pure modifiers available, together with the recent advances in theoretical methods, we believe that this approach will provide the basis for a general advance in the design of ultraselective chiral catalysts. We are indebted to the SERC Catalysis and Interfaces Imtiative and ICI Katalco for financial support. Computational results were obtained using software programs from Biosym Technologies of San Diego. References 1 H.B. Kagan in Asymmetric Synthesis, ed. J. D. Morrison, Academic Press, Orlando vol. 5 (1985)

219 2 S. Akutagawa, Appl. Catal. A., 128 (1995) 171 3 K.T. Wan and M. E. Davis, Nature, 370 (1994) 449 4 W. Reschetilowski, U. Bohmer, and J. Wiehl, Stud. Surf. Sci. Catal.,84 (1994) 2021 5 R. Mahrwald, U. Lohse, I. Girnus, and J. Caro, Zeolites, 14 (1994) 486 6 A. Corma, M. Iglesias, C. del Pino, and F. Sanchez, Stud. Surf. Sci. Catal., 75C (1993) 2293 7 S. Feast, D. Bethell, P. C. B. Page, F. King, C. H. Rochester, M. R. H. Siddiqui, D. J. Willock, and G. J. Hutchings, J. Chem. Soc., Chem. Comm., (1995) in press 8 A.C. Butler and C. P. Nicolaides, Catal. Today, 18 (1993) 443 9 P.C.B. Page, M. T. Gareh, and R. A. Porter, Tetrahedron: Asymmetry, (1993) 2139. 10 Catalysis Users Guide, release 236, Biosym Technologies, San Diego (1994). 11 CADPAC5: The Cambridge Analytic Derivatives Package Issue 5, A suite of quantum che~mistry programs with contributions from IL Alberts, JS Andrews, SM Conwell, NC Handy, D Jayatilaka, PJ Knowles, R Kobayashi, N Koga, KE Laidig, PE Maslen, CW Murray, JE Rice, J Sanz, ED Simandiras, AJ Stone and M-D Su, (1992). 12 D. Willock, S. Feast, P. C. B. Page, D. BetheU, and G. J. Hutchings, Topics in Catalysis, (1995) in press. 13 J. R. Hill and J. Sauer, J. Phys. Chem., 98 (1994) 1238 14 J. R. Maple, T. S. Thatcher, U. Dinur, and A. T. Hagler, Chemical Design Automation News, 9 (1990) 10 15 R. D. Amos and J. E. Rice, CADPAC: The Cambridge Analytic Derivatives Package 5, issue 4.0, Cambridge, (1987)

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J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All fights reserved.

221

Enantioselective hydrogenation catalysed by palladium T.J. Hall, P. Johnston, W.A.H. Vermeer, S.R. Watson and P.B. Wells School of Chemistry, University of Hull, Hull, HU6 7RX, United Kingdom The enantioselective hydrogenations of five unsaturated carboxylic acids (including tiglic and angelic acids) catalysed by 1% Pd/SiO2 and of methyl pyruvate catalysed by 4% Pd/Fe203 are described. Enantioselectivity was induced by adsorption of the alkaloids cinchonidine and cinchonine onto the catalysts. Reactions occurred at room temperature and 10 bar pressure. In all reactions, modification with cinchonidine gave S-product in excess, whereas modification with cinchonine gave R-product in excess. Best values of the enantiomeric excess were 27% in acid hydrogenation and 15% in ketoester hydrogenation. The unsaturated acids undergo selective enantioface adsorption adjacent to the modifier; the process has been modelled and the sense of the observed enantioselectivity interpreted. A D-tracer experiment has shown that the ketoester undergoes dissociative adsorption. The reaction kinetics and the deuterium distribution in the product are consistent with the major route to product being conversion to adsorbed-enol and subsequent C=C hydrogenation.

1. INTRODUCTION The enantioselective hydrogenation of the methyl and ethyl esters of pyruvic acid catalysed by cinchona-modified Pt and Ir has been the subject of detailed study and the mechanism of chiral induction is understood [ 1-6]. In these reactions, modification of the catalyst by the adsorption of cinchonidine gives R-lactate in excess in the product, whereas cinchonine gives S-lactate in excess. High values of the enantiomeric excess, above 80~ are achievable. In 1988, Blaser and co-workers reported reversed enantioselectivity over a Pd/C catalyst, i.e. adsorption of cinchonidine induced an enantiomeric excess of 4% in favour of the S-enantiomer [7]. This remarkable result is here confirmed using Pd/Fe203 as catalyst and it is shown that the Pdcatalysed reaction differs in every significant particular from the Pt-catalysed reaction. It became evident during this study that enolisation of the reactant was occurring and that the reaction was one of C=C hydrogenation. Accordingly, we examined for comparison the hydrogenation of several esters of unsaturated carboxylic acids (e.g. methyl tiglate, methyl angelate) but observed no enantioselectivity. However, following Nitta and co-workers' report of the enantioselective hydrogenation of E-a-phenylcinnamic acid catalysed by cinchonidine-modified Pd [8], the free acids were examined and enantioselectivity observed. This paper accordingly reports several examples of the enantioselective hydrogenation of the carbon-carbon double bond.

222 2. EXPERIMENTAL

2.1 Materials 1% Pd/SiO2 was prepared at ICI Katalco and 4% Pd/Fe203 at Johnson Matthey. Reactants used are shown in Table 1. Trifluorotiglic acid and its ester were synthesised by D. Hunter at Glasgow University. The other unsaturated carboxylic acids (Aldrich), methyl pyruvate (Fluka), but-3-en-2-one (Aldrich), cinchonidine (Aldrich), cinchonine (Hopkin and Williams), ethanol (AnalaR grade, BDH) and THF (non stabilised, Fisons) were used as received. Table I Reactants Ri/_

COOH

R2x ..--c ~/ R 3

O/~

H3C

acids (see below)

/OCH3

--c

%0

methyl pyruvate

H2C~-~-CH

H3c / but-3-en-2-one

R1

R2

R3

Name

Trivial description

H H Me Et Et

Me CF3 H H H

Me Me Me Me Et

E-2-methyl-but-2-enoic acid E-2-methyl-4,4,4-trifluoro-but-2-enoic acid Z-2-methyl-but-2-enoic acid Z-2-methyl-pent-2-enoic acid Z-2-ethyl-hex-2-enoic acid

Tiglic acid Trifluorotiglic acid Angelic acid 2MP2 acid 2EH2 acid

2.2 Procedure: hydrogenation of unsaturated carboxylic acids and their methyl esters 0.3 g samples of Pd/SiO2 were evacuated and reduced in 1 bar hydrogen at room temperature for 0.5 h. The catalyst was then wetted by injection into the reduction vessel of a 10 ml aliquot of alkaloid solution (0.05 g in 50 ml THF). The thoroughly wetted catalyst was transferred to the glass liner of a Baskerville steel autoclave and the remaining 40 ml alkaloid solution added. The liner was positioned in the autoclave, 5 mmol reactant added, and the autoclave sealed and purged with pure nitrogen. Immediately thereafter hydrogen was introduced to the desired pressure and stirring at 1200 rpm commenced. Hydrogen pressure was maintained at the set value by a computer-controlled admission system and a data system logged the progress of reaction. Reactions were stopped after 20 h reaction time.

2.3 Procedure: hydrogenation of methyl pyruvate 0.1 g samples of Pd/Fe203 were evacuated and reduced for 1 h in 1 bar hydrogen at the required temperature (normally 293 K). After brief evacuation of the hydrogen the catalyst was wetted by injection into the reduction vessel of a 10 ml aliquot of alkaloid solution (0.2 o in 40 ml ethanol). The thoroughly wetted catalyst was transferred into an open glass beaker, the remaining 30 ml alkaloid solution was added and the slurry was stirred in air for 1 h. The catalyst was then separated by centrifugation and decantation, transferred to a Fischer Porter glass high-pressure reactor, and 20 ml fresh solvent and 113 mmol pyruvate added. The

223 reactor was then sealed, purged, hydrogen introduced, and stirring commenced A computercontrolled hydrogen admission system maintained the set pressure and recorded the course of reaction as a function of time. When unmodified catalysts were examined the procedures were as described above except that the alkaloid was omitted.

2.4 Analysis When the desired hydrogen uptake had been achieved, the vessel was opened, catalyst separated by filtration, and the reaction solution analysed by chiral gas chromatography (column: Cydex B, 50 m, SGE Ltd). Analysis gave conversion and enantiomeric excess. Enantiomeric excess is defined a s [ R - S [/(R+S).

3. RESULTS

3.1 Hydrogenation of unsaturated acids and their methyl esters The hydrogenation of the esters methyl tiglate, methyl trifluorotiglate, and methyl angelate occurred slowly over 1% Pd/SiO2 at room temperature and l0 bar pressure. Modification of the catalyst by cinchonidine or cinchonine induced no enantioselectivity under any conditions Typical results are given in Table 2. Hydrogenation of the free acids over unmodified catalyst occurred slowly, proceeded to completion in 20 h and gave racemic product as expected Enantioselective hydrogenation occurred at a slower rate over alkaloid-modified catalyst, cinchonidine modification providing an excess of S-product and cinchonine an excess of R-product. Racemic and enantioselective hydrogenations of tiglic acid each exhibited an apparent activation energy of 17 kJ tool"1 (268 to 308 K). Enantiomeric excess was constant at 20 to 23% over the range 273 to 308 K but lower, 13%, at 268 K. Enantioselective hydrogenation of trifluorotiglic acid exhibited an activation energy of 23 kJ toolI (253 to 323 K) and a temperature-independent enantiomeric excess of 13 • 2%. No geometrical isomerisation of tiglic acid to angelic acid, or vice versa, accompanied hydrogenation Enantioselective hydrogenation Z-2-methyl-pent-2-enoic and Z-2-ethyl-hex-2-enoic acids occurred over alkaloid-modified PdYSiO2 as described in Table 2. Enantioselectivity was favoured by an increase in hydrogen pressure to 50 bar. The enantiomeric excess of 27% in Z2-methyl-pent-2-enoic acid hydrogenation was the highest value recorded in this study. Enantioselectivity in acid hydrogenation was not sensitive to the reduction temperature of the catalyst Modification by N-benzylcinchonidinium chloride substantially deactivated the catalyst and eliminated enantioselectivity. 3.2 Hydrogenation of Methyl Pyruvate The hydrogenation of methyl pyruvate proceeded over 4% Pd/Fe203 at 293 K and 10 bar when the catalyst was prepared by reduction at room temperature. Racemic product was obtained over unmodified catalyst; modification of the catalyst with a cinchona alkaloid reduced reaction rate and rendered the reaction enantioselective. S-lactate was formed in excess when the modifier was cinchonidine, and R-lactate when the modifier was cinchonine

224 Table 2 Hydrogenation of various unsaturated acids and esters catalysed by 1% Pd/SiO2 and 10 bar pressure and 293 K

Reactant

Modifier a

Initial Conversion rate /% /mmol h'lg "1

Methyl tiglate Ethyl trifluorotiglate Methyl angelate

CD CD CD

10 2 8

35 6 33

Tiglic acid Tiglic acid Tiglic acid Trifluorotiglic acid

none CD CN CD

18 7 5 14

100 65 60 95

0 22 (S) 20 (R) 15 (S)

Angelic acid Angelic acid Angelic acid

none CD CN

29 20 18

100 66 56

0 15 (S) 14 (R)

2MP2 2MP2 2MP2 2MP2

CD CD CN CN

18 (10 33 (50 15 (10 19 (50

61 65 38 60

20 (S) 27 (S) 19 (R) 22 (R)

none CD CN

29 18 15

100 62 51

0 12 (S) 10 (R)

acid acid acid acid

2EH2 acid 2EH2 acid 2EH2 acid

bar) bar) bar) bar)

ee /%

0 0 0

aCD = cinchonidine; CN = cinchonine (Table 3). Enantioselective reaction was of order 0.7 in hydrogen by the initial rate method (over the range 2 to 50 bar, 293 K, cinchonine modifier) and 0.2 in pyruvate (0.1 to 3.0 M, 293 K, 10 bar pressure, cinchonine modifier). Enantiomeric excess was independent of reactant concentrations within these ranges. Reactions exhibited self-poisoning so that complete conversion was not achieved within 20 h reaction time. As the quantity of cinchonine modifier added to the catalyst was increased from zero to 1 gram per gram so the initial reaction rate fell from 180 to 50 mmol h l g 1 and enantiomeric excess rose to 15% (293 K, 10 bar pressure).

225

Table 3 Hydrogenations of methyl pyruvate (A) and ofbut-3-en-2-one 03) over 4% Pd/Fe203 at 293 K Catalyst reduction T/K

Modifer a

293 293 293 293 293 293 673

None CD CN CN CN CN CN

A A A A A A A

10 10 2 10 20 50 10

180 36 57 102 147 193 v.slow

50 25

293

CD

B

10

>9,500

100 b

Reactant

Pressure /bar

Initial rate /mmol h l g -1

Cony. /%

Enantiomeric excess /% 0 5 (S) 12 (R) 13 (R) 13 (R) 14 (R) 0

4

~CD = cinchonidine; CN = cinchonine bProduct = butan-2-one, 100%

Table 4 Deuterium distributions in methyl pyruvate and methyl lactate (X - H or D)

x =

CX3COCOOCH3/% 0 1 2

-dx obs. 6 -dx calc. a 12

39 38

47 38

3

0

8 12

15 3

CX3CX(OX)COOCH3/% 1 2 3 4 13 16

31 31

32 31

8 16

2 3

"binomial distributions for complete exchange at each site. H = 50%, D = 50%

A D-tracer experiment was conducted under standard conditions (293 K, 10 bar pressure) using D2 in place of H2 and C2HsOD as solvent in place of unlabelled ethanol. Dincorporation in reactant and each enantiomer of the product was determined by chiral-gc/ms. Exchange of up to three H-atoms for D occurred in the reactant. The product enantiomers gave identical mass spectra indicating the same D-content and distribution; the product contained 0 to 5 D-atoms. Fragmentation in the mass spectrometer showed that no deuterium was located in the ester group. Thus, the retrieved reactant was CX3COCOOCH3 and the product was CX3CX(OX)COOCH3 (X = H or D). The deuterium distribution in the reactant and products was determined from the mass spectra by application of 13C, 180 and ion fragmentation corrections in the usual way. The isotopic distributions so obtained are shown in Table 4. This fragmentation correction is made on the (usual) assumption that the

226 probability of rupture of H-C, D-C, H-O, and D-O bonds in the mass spectrometer is the same, which is unlikely to be the case. For this reason, and the fact that the parent ion currents of the lightest ions are subject to a greater degree of fragmentation correction, the values quoted for the concentrations of pyruvate-do and lactate-do and -dl are less reliable than those of the more extensively exchanged species. The experimental D-distributions are in modest agreement with those calculated for the cases in which (i) pyruvate has undergone complete exchange at three positions with a pool of adsorbed 'hydrogen' of composition H=50%, D=50%, (ii) lactate has undergone complete exchange at five positions with the same pool of adsorbed 'hydrogen' (Table 4). Pd/Fe203 prepared by reduction at elevated temperature was less active and enantioselective than samples reduced at 293 K (Table 3, entry 7).

4. DISCUSSION

4.1 Hydrogenation of unsaturated carboxylic acids The prerequisite for enantioselective hydrogenation at a metal surface is that selective enantioface adsorption of the prochiral reactant should occur at metal atom sites in the neighbourhood of the adsorbed chiral modifier. The modifiers cinchonidine and cinchonine, in the free state, each exhibit three configurations of comparable minimum energy [1, 9]. When these alkaloids are adsorbed at a Pt surface, one of these minimum energy configurations provides the chiral environment for the selective enantioface adsorption of methyl pyruvate. In that reaction, the rate at the enantioselective sites is enhanced over that at sites elsewhere on the surface that catalyse racemic reaction, and values of enantiomeric excess above 80% are obtainable. The enhanced rate has been attributed to an effect of H-bonding between the OHgroup of the adsorbed half-hydrogenated state and the quinuclidine-N of the adsorbed modifier

[1]. In this Discussion the assumption is made that the adsorption of cinchonidine and cinchonine on Pd is similar to that on Pt [10]. In the hydrogenation of the unsaturated acids and esters over cinchona-modified Pd the key observation is the failure to achieve enantioselectivity in ester hydrogenation and the success encountered in the hydrogenation of the free acids (Table 2). Also, the enantioselective hydrogenation of E-o~-phenylcinnamic acid over cinchonidine-modified Pd has been reported [8, 11 ]. It appears either that the free acid is able to undergo selective enantioface adsorption in the vicinity of the modifier in a manner forbidden to the methyl ester for steric reasons, or that an acid-base interaction between reactant and alkaloid occurs as a precursor state to selective enantioface adsorption which, again, would be unavailable to the ester. Such an acidbase interaction is most easily envisaged as H-bonding between the acid-H atom of the reactant and the quinuclidine-N of the alkaloid. The latter situation has been modelled [12]. Calculations have been carried out in which tiglic acid has been docked with cinchonidine in such a way as to simulate such H-bonding. The docking was carried out for the specific condition in which (i) the alkaloid molecule was in its appropriate minimum energy state, and (ii) the quinoline moiety of the alkaloid and the carbon/oxygen skeleton of the tiglic acid were maintained in the same plane. [This relationship is the closest approximation achievable to that which might obtain if the complex of alkaloid and acid was adsorbed at a plane metal surface.

227 No allowance for the presence of the surface has been made in these calculations.] Two configurations meet these criteria; they are shown schematically in Figure 1. Figure I a shows tiglic acid adsorbed by the enantioface which, on hydrogenation would give S-2-methyl butanoic acid, whereas Figure l b represents the enantioface that would give R-2-methyl butanoic acid. The interconversion of these two states by rotation of the whole tiglic acid molecule about the =--N. . . . H- hydrogen bond was modelled, and showed that the energy of the state represented in Figure l a was lower than that in Figure l b by about 5 kcal mol l moreover, the state represented in Figure l a is located within a broad region of minimum energy. It can thus be inferred that, for the adsorption of tiglic acid adjacent to adsorbed cinchonidine, selective enantioface adsorption occurs such that, on hydrogenation, S-product formation is favoured, in agreement with experiment. H

Hj

0 ~ / C H 3 (S)-2-Methyl Butanoic Acid CH3

(a) H

""'H

H

(R)-2-Methyl Butanoic Acid CH3

(b)

Figure 1. Possible precursor states to the selective enantioface adsorption of tiglic acid

228 The carbon-carbon double bond that undergoes hydrogenation is remote from the modifier and no rate enhancement for the enantioselective process is to be expected. None was observed. Moreover, since the rate at the enantioselective sites is the same as that at other sites on the surface that experience no chiral environment and so give racemic product, the overall enantiomeric excess should be modest, as is the case. To obtain higher enantioselectivity it would be necessary selectively to poison the sites for racemic hydrogenation. The values of enantiomeric excess observed under comparable conditions vary on passing from one acid to another (Table 2) but show no significant trends. The mechanism proposed for the hydrogenation of tiglic acid is applicable to the other acid hydrogenations studied.

4.2 Hydrogenation of methyl pyruvate Although cinchona-modified Pd showed no enantioselectivity in the hydrogenation of the methyl esters of the unsaturated acids, the hydrogenation of methyl pyruvate occurred with a modest enantiomeric excess. The reduction of this ester over Pd d~fered from the corresponding reaction over Pt in every important particular. Enantiomeric excess was low (high over Pt) and in the reverse sense (e.g. cinchonidine modification provided an S-excess in the product over Pd but an Rexcess over Pt). Enantioselective reactions underwent self-poisoning over Pd (proceeded to completion over Pt), were of non-integral order (integral over Pt) and proceeded more slowly than reaction over unmodified catalyst (enhanced rate over Pt). Enantioselective reaction was solvent-specific over Pd (not over Pt) and was favoured by low catalyst reduction temperature (high reduction temperature for Pt). The exchange of H for D in the pyruvate methyl group, which occurred over Pd but not over Pt [13], holds the key to these profound differences in behaviour. This exchange indicates that adsorbed pyruvate (species A, Figure 2) underwent dissociative adsorption at the Pd surface (species B is one formulation of the product of dissociation) and interconversion of species A and B in the presence of a pool of adsorbed-H and -D brought about complete exchange in that methyl group. Occasional formation of the adsorbed enol, species C, is expected as an alternative product of hydrogen-atom addition to species B. Lactate will be formed by hydrogenation of species A or species C, or both, depending upon their relative reactivity. When the C=C and C=O functions were present in the same molecule, but-3-en-2-one, hydrogenation of the former function occurred exclusively and at a rate almost too fast to measure (Table 3). Thus removal of species C is expected to be kinetically fast and that of species A kinetically slow. Strong evidence that the lactate product was formed by enol hydrogenation is provided by the observed order of 0.7 in hydrogen. The rate determining step would necessarily be the H-atom addition to species B to give C; supposing the H-coverage to be described by the appropriate Langmuir equation for dissociative hydrogen adsorption, a fractional order in hydrogen in the region of one-half is expected. It is thus concluded that methyl pyruvate hydrogenation over Pd is a kinetically fast hydrogenation of adsorbed enol formed via the dissociative adsorption of the ot-ketoester. Pd catalysts were active and enantioselective only when reduced at low temperature, suggesting that dissociative adsorption of the reactant was dependent on the presence of surface Pd atoms in a positive oxidation state.

229

O•C

H3C/

HO R

+2H Slow

A

H ~C--

R

H3~

o\ (D)H,--)/C H2c

"/

R B r.

§

HO C

H2

~, ~"

+2H _ Fast -

HO H 9 C

R

H3C/

Figure 2. Mechanism for exchange in methyl pyruvate and for its hydrogenation to methyl lactate via an enol intermediate. R = COOCH3 Enolisation of ketones is favoured in alkanol solution, and the observed solvent specificity in this reaction may indicate that the formation of the enolic species C is favoured when ethanol is used as solvent. No rate enhancement of the enantioselective hydrogenation pathway is expected, in the manner adduced for the Pt-catalysed reaction, because the process is not one of simple H-atom addition across a carbon-oxygen double bond. The self-poisoning character of the reaction was most evident when the modifier was present, from which it may be inferred that adsorbed cinchonidine and cinchonine were convened to an ineffective form and poisoned the surface as reaction progressed. In view of the very high activity of this Pd catalyst for C=C hydrogenation, it may be that the quinoline system of the alkaloids was partially hydrogenated under reaction conditions. Such a proposal would be consistent with our observation that H/D exchange in cinchonidine (over a different Pd catalyst) was accompanied by ring hydrogenation [ 10]. Cinchona alkaloid derivatives having a partially saturated quinoline ring system are poor modifiers of Pt [ 14]. No discussion is offered, concerning the sense of the observed enantioselectivity in this reaction because of a need to be cautious. We have observed that the sense of the enantioselectivity can be reversed simply by a variation of the procedure used for catalyst modification [15], and this has been confirmed by others [16]. Thus it appears that the state of the Pd surface, as well as the nature of the species adsorbed upon it and their spatial relationship, contributes to chiral direction in this reaction.

230 ACKNOWLEDGMENTS

We thank EPSRC, ICI Katalco, Zeneca, and Johnson Matthey for financial support. Acid hydrogenation was conducted as part of the EPSRC/DTI LINK Programme in 'New Catalysts and Catalytic Processes': our colleagues were G. Webb, E. Colvin, E. Allan, D. Hunter and N. Young of the University of Glasgow, S.D. Jackson of ICI Katalco, W. Moss and G. Robinson of Zeneca Pharmaceuticals and S. Korn of Zeneca Fine Chemicals. Pyruvate ester hydrogenation was carried out as part of a programme funded by EPSRC and Johnson Matthey: our colleagues were I.L. Dodgson, A. Fulford, K.G. Griffin and B. Harrison from the Company. We also acknowledge stimulating discussions with A. Ibbotson.

REFERENCES

1.

2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16.

K.E. Simons, P.A. Meheux, S.P. Griffiths, I.M. Sutherland, P. Johnston, P.B. Wells, A.F. Carley, M.K. Rajumon, M.W. Roberts and A. Ibbotson, Recl. Trav. Chim. PaysBas, 113 (1994) 465. O. Schwalm, B. Minder, J. Weber and A. BaJker, Catal. Letts., 23 (1994) 271. K.E. Simons, A. Ibbotson, P. Johnston, H. Plum and P.B. Wells, J. Catal., 150 (1994) 321. H-U. Blaser, H.P. Jalett, D.M. Monti, A. Balker and J.T .Wehrli, Stud. Surf. Sci. Catal., 67 (1991) 147. H-U. Blaser, Tetrahedron: Asymmetry, 2 (1991 ) 843. G. Webb and P.B. Wells, Catal. Today, 12 (1992) 319. H-U. Blaser, H.P. Jalett, D.M. Monti, J.F. Reber and J.T. Wehrli, Stud. Surf. Sci. Catal., 41 (1988) 153. Y. Nitta, Y. Ueda and T. Imanaka, Chem. Lett., (1994) 1095. K.E. Simons, PhD thesis, University of Hull (1994). G. Bond and P.B. Wells, J. Catal., 150 (1994) 329 J.R.G. Perez, J. Malthete and J. Jacques, C.R. Acad. Sci. Paris Serie II, (1985) 169. S.R. Watson, PhD thesis, University of Hull (1995). I.M. Sutherland, A. Ibbotson, R.B. Moyes and P.B. Wells, J. Catal., 125 (1990), 77. H-U. Blaser, H.P. Jalett, D.M. Monti, A. Baiker and J.T. Wehrli, Stud. Surf. Sci. Catal., 67 (1991) 147. S.P. Gfiffiths, P. Johnston and P.B. Wells, unpublished work. P. Collier, J. Iggo and R. Whyman, personal communication.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996Elsevier Science B.V. All rights reserved.

231

S t e r e o c h e m i c a l S t u d i e s of the E n a n t i o - d i f f e r e n t i a t i n g H y d r o g e n a t i o n of Various Prochiral Ketones over Tartaric Acid-Modified Nickel Catalyst T. Sugimura, a T. Osawa, bt S. Nakagawa,ar T. Harada c and A. Tai a* aFaculty of Science, Himeji Institute of Technology; Kamigori, Ako-gun, Hyogo 67812 Japan bFaculty of General Education, Tottori University; Koyamacho-minami, Tottori 680 Japan CFaculty of Science and Technology, Ryukoku University, Seta, Otsu 520-21 Japan Abstract

The stereochemistry and enantiomeric excess (ee) of the products of the enantiodifferentiating hydrogenation of various prochiral ketones over (R,R)-tartaric acid modified Raney nickel catalyst (TA-MNi) were studied in detail. The results were consistent with our newly introduced concept of stereocontrol in that there were two types of interaction modes between adsorbed TA and the substrate on (R,R)-TA-MNi; a two point interaction to give a product of R configuration and an one point interaction to favor a product of S configuration. The relative contribution of these two modes determined the optical purity and stereochemistry of the product depending on the structure of the substrate. In the hydrogenation of a 3-oxoalkanoate where the contribution of two points interaction predominated, elimination of the minor contribution of the one point interaction by tuning the substrate structure enabled us to achieve 96% of ee, the highest value so far obtained using a heterogeneous catalyst.

1. INTRODUCTION Asymmetrically modified nickel catalyst (MNi) is one of the promising and wellinvestigated heterogeneous catalysts for the enantioface-differentiating hydrogenation of prochiral ketones. Many kinds of catalyst preparations and mechanisms of enantio-differentiation with this catalyst have been proposed by various research groups [1]. The best catalyst available at present is TA-NaBr-MRNi (tartaric acid-NaBr-modified Raney nickel catalyst) developed [2] and improved [3] by our group. Under well-optimized conditions, this catalyst gave 86% enantiomeric excess (ee) in the hydrogenation of methyl acetoacetate (MAA) to methyl t Present address: Facultyof Science,ToyamaUniversity,Gofuku,Toyama930 Japan ~:ReserchFellow fromToyoKasei KogyoCo. Ltd., Sone,Takasago676 Japan

232 3-hydroxybutanoate (MHB) and 80% of ee in the hydrogenation of 2-octanone to 2-octanol in the presence of an excess amount of pivalic acid [4]. The reaction site of TA-MNi could not be homogeneous judging from its existence on the solid surface. Through various investigations, it has been deduced that there are two types of sites on the catalyst; one is an enantio-differentiating site (E), where an optically active product is produced with the aid of an adsorbed chiral auxiliary (TA) and the other is a non-enantio-differentiating site (N), where the chiral auxiliary has no effect and the product becomes racemic. When the intrinsic enantiodifferentiating ability of the E site (factor i), which is related to the efficiency of the mutual interaction between TA and the substrate, is taken into account, the observed enantio-differentiating ability (e.d.a.) determined as the ee of the product is formulated according to eq. 1 (reaction site model) [5].

(1)

e.d.a. (%) = [i E / ( E + N)]xl00

The improvements in MNi so far achieved were mostly due to our efforts to eliminate the N site from MNi by changing preparation variables of the catalyst,. while the other important factor i has not been satisfactorily considered. In the present study, hydrogenation of various prochiral ketones with TA-MNi almost freed from the N site were carried out in order to gain insight into the mode of stereo control on MNi, which was expected to determine the stereochemistry of the reaction and to take part in the origin of the factor i.

2. RESULTS Two series of the reactions carried out in this study are shown in eqs 2 and 3.

H2 CH3 C(CH2)n COOR O

= (R,R)-TA-MNi

CH3 C*H(CH2)n COOR

(2)

OH

n = 0,1,2,3

RlCCH2COOR 2 O

H2 (R,R)-TA-MNi

R1 C.HCH2COOR 2

(3)

OH

R, R1,R2 : alkyl group of various chain lengths and branching The results of the series of reactions shown in eq. 2 are listed in Table I together with our early reported data on the h y d r o g e n a t i o n of 2-octanone (7) [4]. The hydrogenation on all substrates proceeded smoothly and gave the corresponding chiral secondary alcohol. In the case of 3_, 4_,5_, and 6_, some amounts of lactone were produced as by-product. From this study, quite interesting stereochemical behavior

233 T a b l e 1. Enantioface-Differentiating H y d r o g e n a t i o n of Various Keto Esters a n d 2Octanone reaction conditions .......p. r ~ configuration entry substrate temp.(~ addition of confi ee Pivalic acid am

ii

l i,,

i,

o .,o,~,

1

Jl

1

i

i

i,

i

100

no a

R

15

100

y esb . . . . . R .

14

I I

2

O

3

O

O

4

~O,K,.~

2

.,

.....

.......

5

.,O,r

6

0

0

...

no

R

85

ye s

R

.........72

100

no

R

38

100

yes . . . .

R

5

100

s .

..

100 ,,,,,

,,,,,,

.

100

no

4

120

yes

S

33

R = i-Bu

100

yes

S

49

10

80

yes

S

58

11

60

yes

S

61

12

40

yes

S

58

7

8

O

O

R.O.JJ,..A.~

9

61

13

s

R = Me

60

yes

14

6

R = n-Bu

60

yes

100

no

S

9c

120

yes

S

50 c

17

100

yes

S

62 c

18

80

yes

S

66 c

19

60

yes

S

74 c

20

40 . . . . . . .

yes

S

74 c

15 16

O ~

N

7

.....

a: A trace amount"0f acetic acid (2% w t / w t ) was added to the substrate. b: Pivalic acid (50% w t / w t ) was added to the substrate. c: The results reported in ref. 4.

59

.. . . . . . . .

, i,

234 T a b l e 2. Enantioface-Differentiating H y d r o g e n a t i o n of Various Alkyl 3oxoalkanoates entry

substrate

o

0

temp.

reaction

(~

time (h)

ee a

100

4

85

60

31

84

40

43

81

100

18

91

60

34

94

40

75

94

100

55

88

60

71

96

40

95

96

100

30

_C

60

67

_C

100

36

87

60

52

90

14

40

66

87

15

100

36

69(PA) b

16

60

43

71(PA) b

100

48

85

100

24

85

60

45

87

100

33

84

40

88

. .0 % e L

2

O

O

--oX..fl-./ ~0,.~~ O

O

..o.J~ O

10 11 12

O

O

13

N

O O

~

"-0~ O A c

17

O

18

O

19

-J-o

20

../

10

11

O

12

O 13

O,,

O ,,

21 )K'O'A""~ 14 60 a: All products were R excess except the product from 9. b: Pivalic acid (9 g) was added to a solution of the substrate in THF (10 ml). C: No reaction.

235 of the MNi was observed. In the reactions carried out in the absence of pivalic acid, the ee of the R configuration in excess was very low in I (n = 0), extremely increased in 2 (n = 1), decreased in 3 (n = 2), and then diminished in 4 (n = 3), whereas in 7 a slight excess of the S configuration was observed (Table I entries 1, 3, 5, 7, 15). When a l a r g e a m o u n t of pivalic acid was added in the reaction media, each substrate showed a characteristic change in stereochemistry and ee of the reaction. With the addition of pivalic acid, the ee in R excess was u n c h a n g e d in 1_, significantly decreased in 2 and 3_, while an appreciably high ee in S excess was observed in 4 as well as in _7 (Table 1, entries 2, 4, 6, 9, 17). The change in the alkyl group of the alkoxy side in 4 from/-butyl to methyl (5) or n-butyl (6) resulted in no essential change in ee (Table 1, entries 11, 13, 14). Our early study indicated that the ee in reaction 7 with pivalic acid s h o w e d significant temperature dependence [4]. The present study showed that reaction 4 with pivalic acid also resulted in a temperature dependence similar to that of Z (Table 1, entries 8 to 12). The results of a series of reactions shown in eq 3 are summarized in Table 2. All reactions proceeded to over 90% conversion and afforded almost a quantitative yield of 3 - h y d r o x y e s t e r with more than 85% ee except for the reaction of 1___00. The hydrogenation of 10 which carried a bulky t-butyl group next to the carbonyl group did not proceed (Table 2, entries 10 and 11). The elongation of the chain or branching of the alkyl group at the acyl side tended to increase the ee of reaction (Table 2, entries 1, 4, 7, and 12). Except for the reaction with 2 of which the ee showed no temperature dependence, the ee of the reaction with 8, 9_, and 1__!increased with a decrease in reaction temperature and reached to plateau at 60 ~ (Table 2, entries 1 to 9 and 12 to 14). The hydrogenation of 1__! in the presence of pivalic acid also decreased the ee same as in the case of _2 (Table 2, entries 15 and 16 and Table 1, entry 4 ). Introduction of a bulky alkoxy group in acetoacetate, 13 and ~ gave almost the same result as 2 (Table 2, entries 18 to 21).

3. D I S C U S S I O N S

In o r d e r to explain the stereochemistry of the enantioface-differentiating hydrogenation over TA-MNi, we have proposed a two point interaction (2P) mode for MAA (2) [6] and a one point interaction (1P) mode for 2-octanone (7) [7]. Figs 1 and 2 s h o w a detailed sketch of each model provided by (R,R)-TA-MRNi. These models can be simplified as shown in Figs 3 and 4, respectively. Both models have been p r o p o s e d based on the concept that the adsorption mode of TA should be the same in each case and one of the OH groups in TA located close to the Ni surface (site 1) interacts with C=O of the substrate to be hydrogenated at the immediate surface of Ni through hydrogen bonding. The 2P mode achieved when the substrate possesses a second functional group to associate with the second OH in TA (site 2) as is the case with M A A gave a product of R configuration. If the substrate has no additional functional group as is the case with 2-octanone, association of TA and the substrate becomes 1P mode. In this case, a weak steric repulsion takes place between the OH group of site 2 and the alkyl chain of the substrates and the orientation of the adsorbed substrate favors the formation of the S configuration to some extent. Pivalic

236 acid employed as an additive in the reaction system is expected to enhance the steric effect by forming a bulky repulsive fence at site 2 with the association and making the stereo-control ability of 1P effective [4]. The contribution of these two modes to the stereo-controlling ability of the E site is considered to be different. The contribution of the 2P mode directly relates to the formation of the R product, while the 1P mode contribution varies from the formation of an almost racemic product to the formation of a product with an appreciably high excess of S depending on type of substrate and the reaction conditions. Pivaric acid,

0

9 C H

site 2 site 2

o

MAA

(R,R)-TA

(R,R)-TA

site 1

H21L ~

~

H,

Nickel Surface

Nickel Surface

2-,

Figure 1. Two-point interaction Model (2P) of MAA

,r

/

Figure 2. One-point interaction Model (1 P) of 2-Octanone C

\ site 2 O ..... O ~

,r

site 2

.

..... o ' ~

'i o

"i' 0 iHi

Figure 3. Two-point model (2P)

OH

~L

i

|l lllll

H,,i

,ill

Figure 4. One-point model (1P)

Taking the above mentioned characteristics of the two modes into consideration, we introduced the concept of stereo-control in the enantio-differentiating hydrogenation of various functionalized prochiral ketones on TA-MNi based on the coexistence of 2P and 1P on the E site of the catalyst. That is the 1P function counteracts the 2P function when 1P and 2P coexist, and the relative contribution of the two modes determine the stereochemistry of the product produced in excess and also relates qualitatively to the i factor. From the study of a series of substrates, I to 4 and 7_, summarized in Table 1, the (R,R)-TA-MNi found to show quite interesting stereochemical behavior in connection with the concept of stereocontrol mentioned above. From the stereochemistry of the

237 hydrogenation product in the absence of pivalic acid (Table 1, entries, 1, 3, 5, 7, 15), it was expected that the contribution of 2P was in excess for K to 3, and the 2P contribution became m a x i m u m in 2_, while the contribution of 1P and 2P were compensated in 4_, and exclusive 1P participation in 7 resulted in a slight S excess. From the ee of each reaction product (Table 1, entries 1, 3, 5, 7, 15), participation of 2P and 1P was expected to change systematically depending on the degree of fit between the two carbonyl groups in the substrate and the two interaction sites in TA as shown in Figs 3, 5, 6, and 7, respectively. The results of the reactions in the presence of pivalic acid (Table 1, entries 4, 6, 9, 17) clearly showed that pivalic acid in the reaction media decreased the 2P participation by competing with the interaction at site 2 with the substrate and enhanced the 1P contribution and its efficiency by forming a steric fence in the proximity of site 2. It is noteworthy that, in the absence and presence of pivalic acid, the ee of the hydrogenation of 3 changes from an appreciable R excess to negligible R excess and that of the hydrogenation of 4 changes from zero to appreciable S excess, respectively. On these two substrates, the m o d e of stereocontrol was transient from 2P to 1P contribution depending on the presence or absence of pivalic acid. In our early study of the hydrogenation of 2-alkanone with pivalic acid, it was reported that the ee was significantly temperature-dependent and this behavior was characteristic of the 1P m o d e [4]. A similar t e m p e r a t u r e dependence of ee found in 4_ and 7 (Table 1, entries 8 to 12 and 16 to 20) is also additional support of the presumption that the hydrogenation _4 with pivalic acid proceeds mostly by the 1P mode. The decrease of ee with increased temperature in the reaction by the 1P mode was explained by a decrease in steric repulsion due to the decrease in the tightness of h y d r o g e n - b o n d i n g at site 1, a decrease in the associative ability of pivalic acid with site 2, and also an increase of the mobility of alkyl chain of the substrate with the increased temperature. , o"~

[

AN

site 2 ........... 0

9 .r site 2.. .... 0..~/0

/

site 2 ....."....

/

,

0

. . . .

i,o=

Figure 5. Possible 2P mode

Figure 6. Possible 2P mode

Figure 7. Possible 2P mode

for I

for 3

for 4

A series of studies with various 3-oxoalkanoates provided further information on the factor i. As listed in Table 2, all substrates give more than 85% ee in the R configuration in excess. These facts clarify that a strong contribution of 2P is common to all 3-oxoalkanoates. The other important finding is that elongation or branching of the alkyl moiety at the acyl side of a 3-oxoalkanoate tends to increase the ee of hydrogenation. An increment in ee from 2 to 9 is especially significant (Table 2, entries 2 and 8). Fig 8 and Fig 9 are schematic sketches of the S configuration favored the 1P expected in 2 and 9_, respectively. In 2_, no significant steric interaction is

238 expected between the methyl group at the acyl side and the OH of site 2, whereas in 9_, the isopropyl group at the acyl side and the OH of site 2 come into contact, making 1P contribution difficult. In this context, the participation of 1P became less in reaction of 9_ than in reaction of 2 and hence 9__gave higher ee than 2 due to the increase in the relative participation of 2P. The temperature dependence of ee found in reaction 9 suggested that reaction of 9 still involved a small portion of the 1P mode at high temperature. Results of other substrates carrying a long chain at the acyl side; 8_,11 (table 2, entries 4 to 6, 12 to 14,) are also rationalized in the same way as that of 9. Through the series of discussions described above, the various forms of stereochemical behavior of MNi could be explained rationally on the basis of our newly proposed concept. A"x. sitex2 Si~oH~C~-- CH 3

V site"i'O= o 0 Figure 8.1P mode for 2 /

Figure 9.1P mode for 9 /

The 96% ee is the highest so far achieved by the enantio-differentiating hydrogenation over an asymmetrically modified heterogeneous catalyst. Although it is difficult to separate factor i and E~ (E+N), the present results indicated that both of them are well optimized and become almost unity.

4. EXPERIMENTAL 4.1. Preparation of TA-NaBr-MRNi TA-NaBr-MRNi was prepared by the reported method [3]. RNi (W-1 type) was prepared from 1.9 g of Raney nickel alloy (Kawaken Fine Chemical Co., Ni/A1 = 42/58). To wash out the excess base and aluminum salts, a sufficient amount of deionized water was used with ultrasonic irradiation. The modifying solution was prepared by dissolving of (R,R)-tartaric acid (1 g) and NaBr (6 g to 10 g) in 100 ml of water and adjusting the pH to 3.2 with 1N NaOH aqueous solution. RNi was heated in the modifying solution at 100 ~ for 1 hour, washed with water (50 ml), methanol (50 ml, twice), and THF (10 ml). The TA-NaBr-MRNi obtained by this method was immediately used for the hydrogenation. 4.2. Preparation of the Substrates The following substrates were obtained from commercial sources, methyl pyruvate (1), methyl acetoacetate (2), methyl 4-oxopentanoate (3), and methyl 3-oxopentanoate (8). Alkyl 5-oxohexanoates (4, 5 and 6) were prepared by condensation of methyl acetoacetate and methyl acrylate followed by acidic hydrolysis, decarboxylation, and esterification [8]. Methyl 3-oxo-4-methylpentanoate

239 (9), methyl 3-oxononanoate (11), and methyl 3-oxo-lO-acetyloxydecanoate (12) were prepared from Meldrum's acid and 2-methylpropanoyl chloride, heptanoyl chloride, and 8-acetyloxyoctanoyl chloride, respectively by a reported method [9]. i-Propyl acetoacetate (13) and t-butyl acetoacetate (14) were prepared from diketene and the corresponding alcohols in the presence of triethylamine [10]. Each substrate was purified either by fractional distillation or column chromatography before use. 4.3. Hydrogenation

In an autoclave of 100 ml capacity, TA-NaBr-MRNi and a solution of 1.5 g to 10 g of the substrate in 10 ml of THF were placed. In some cases, 9 ml of pivalic acid was added, and in the other cases, 0.2 ml of acetic acid was added. Hydrogen was charged into this at ca. 107 Pa as the initial pressure. Reaction temperatures and times of the hydrogenation are listed in Tables I and 2 in the text. After cooling, the reaction mixture was filtered and purified either by distillation or column chromatography on silica gel depending on the nature of the product. Structural characterization of each purified product was carried out using NMR 0EOR GX 400 spectrometer) and IR (JASCO IR-88 spectrometer) spectra. 4.4. Determination of ee of the product

The enantiomeric excess (ee) of the hydrogenated products was determined either by polarimetry, GLC equipped with a chiral column or 1H-NMR with a chiral shift reagent. Methyl lactate and methyl 3-hydroxybutanoate, obtained from I and 2_., respectively, were analized polarimetry using a Perkin-Elmer 243B instnmlent. The reference values of [(Z]D(neat) were +8.4 ~ for (R)-methyl pyruvate and -22.95 ~ for methyl 3-hydroxybutanoate. Before GLC analysis,/-butyl 5-hydroxyhexanoate, methyl 5-hydroxyhexanoate, and n-butyl 5-hydroxyhexanoate, obtained from 4_, 5_, and 6, respectively, were converted to the pentanoyl esters, methyl 3-hydroxybutanoate was converted to the acetyl ester, and methyl 4-methyl-3hydroxybutanoate obtained from 9 was converted the ester of (+)-~-methyl-(x(trifluoromethyl)phenyl acetic acid (MTPA). The GLC analysis with a Shimazu GC 17A equipped with a column of CPChirasil DEX CB (25 m, 0.25 mm id, GL Science, Japan) was applied for the determination of the ee of the following compounds at the temperature stated in parentheses: i-butyl 5-pentanoyloxypentanoate (150 ~ methyl 5-pentanoyloxypentanoate (130 ~ n-butyl 5-pentanoyloxypentanoate (150 ~ and methyl 3acetyloxybutanoate (100 ~ the MTPA ester of 4-methyl-3-hydroxybutanoate (160 ~ obtained by the above mentioned derivation, and methyl 3-hydroxypentanoate (80 ~ methyl 3-hydroxynonanoate (140 ~ i-propyl 3-hydroxybutanoate (90 ~ and t-butyl 3-hydroxybutanoate (100 ~ obtained directly from ~, !1, 13, and respectively. In all cases, enantiomers showed completely separated peaks, and repeated analysis showed an error within 0.2%. The ee of methyl 10-acetyloxy-3hydroxydecanoate obtained from 12 was determined by 1H-NMR in the presence of (+)-Eu(hfc)3. In the cases of methyl pyruvate and methyl 3-hydroxybutanoate, the ee value was determined by polarimetry and GLC fell within the range of 0.2%.

240 ACKNOWLEDGEMENT This work was partially supported by a Grant-in-Aid for Scientific Research No. 06640697 from the Ministry of Education, Science and Culture, Japan. We also thank Toyo Kasei Kogyo Co. Ltd., who permitted us to use their facilities.

REFERENCES

1. A. Tai and T. Harada in "Tailored Metal Catalysts" (Ed. Y. Iwasawa), D. Reidel, Dordrecht, p. 265 (1986). W. M. H. Sachfler in "Catalysis in Organic Reactions" (Ed. L. Augustine), Chem. Ind., 22, 189 (1985). 2. T. Harada, M. Yamamoto, S. Onaka, M. Imaida, H. Ozaki, A. Tai, and Y. Izumi, Bull, Chem. Soc. Jpn., 54, 2323 (1981). 3. A. Tai, T. Kikukawa, T. Sugimura, Y. Inoue, S. Abe T. Osawa, and T. Harada, Bull. Chem. Soc. Jpn., 67, 2474 (1994). 4. T. Osawa, T. Harada, and A. Tai, J. Cat. 121, 7 (1990). 5. T. Harada, A. Tai, M. Yamamoto, H. Ozaki, and Y. Izumi, Proc. 7th Int. Congr. Catal. Tokyo,p 364 (1980). 6. A. Tai, T. Harada, Y. Hiraki, S. Murakami, Y. Izumi, Bull. Chem. Soc. Jpn., 56,1414 (1983). 7. T. Osawa, T. Harada, and A. Tai, J. Mol. Catal. 88, 333 (1994). 8. S. Okamoto, T. Harada, A. Tai, Bull. Chem. Soc. Jpn., 52, 2670 (1979). 9. Y. Oikawa, K. Sugano, O. Yonemitsu, J. Org. Chem., 43,2087 (1978). 10. T. Kato, T. Chita, Chem. Pharm. Bull., 23, 2263 (1975).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) l l t h International Congress on Catalysis - 40th A n n i v e r s a r y

Studies in Surface Science and Catalysis, Vol. 101 L996 Elsevier Science B.V. .

.

241

.

Enantio-differentiation over heterogeneous catalysts. The shielding effect model. J6zsefL. Margitfalvi, Mihhly Heged0s and Ern6 Tfirst Central Research Institute for Chemistry of the Hungarian Academy of Sciences 1025 Budapest, Pusztaszeri ut 59-67, Hungary This paper deals with the origin of enantio-differentiation over heterogeneous catalysts. A new model is proposed, in which the modifier provides a specific shielding effect. A prochiral molecule, due to the specific character of shielding, can adsorb onto the metal surface by its unshielded site resulting in enantio-differentiation. As emerges from computer modeling, quantum chemical and quantum mechanical calculations made on the catalytic system: cinchonidine - a--keto esters - Pt the shielding effect can be responsible both for the rate acceleration and induction of enantio-differentiation. This model is based on an earlier proposition, which suggest (i) the formation of a weak complex between the modifier and the substrate in the liquid phase and (ii) the hydrogenation of either the shielded or unshielded forms of a-keto ester over the Pt sites. The shielded form gives the optically active, while the unshielded one the racemic product. Further support for this model was obtained in kinetic experiments and kinetic modeling. I. INTRODUCTION There is a growing interest for enantioselective reactions and there is a real need to develop both homogeneous and heterogeneous catalysts with high enantioselectivity. There are only two heterogeneous catalytic reactions in which high optical yields (up to 95 %) were obtained: hydrogenation of 13-keto esters and cx-keto esters over Ni-tartrate and cinchona-Pt/Al203 catalysts, respectively [1-6]. Despite the extensive studies devoated to the above two hydrogenation reactions the exact nature of the enantio-differentiation (ED) is still not really known. The lack of knowledge of the nature of ED involved in heterogeneous catalytic hydrogenation reactions is the main reason that we are far away to design heterogeneous catalysts with the desired enantio-differentiation ability. In this work a new approach is desribed, which can help to understand ED over heterogeneous catalysts. We also hope that this approach can be used to find new modifiers for enantioselective heterogeneous catalytic reactions. The basis for this approach is the steric shielding known in organic chemistry [7,8]. A chiral template molecule can induce shielding effect (SE) in such a way that it preferentially interacts with one of the prochiral sites of the substrate. If a substrate is preferentially shielded its further reaction can take place only from its unshielded site resulting in ED.

2. EXPEREVIENTAL

2.1. Computer modeling Computer modeling was applied to investigate the ability of commonly used modifiers to create SE. In this respect the following calculations were carried out: (i) conformational

242 analysis of the modifier and substrate, (ii) molecular docking; i.e formation of the [modifier substrate] complex, (iii) adsorption of the [modifier-substrate] complex onto Pt surface. The corLformational analysis of the modifier and substrate was performed using the Hypercube: Hyperchem 3.1 programs with the MM+ forcefield. The minimum energy conformations were determined using the Polak-Ribiere (conjugate gradient) minimization algorithm, and the conformational analysis was carried out with rigid quinoline and quinuclidine parts. The molecular docking calculations were performed with BIOSYM programs. The optimization of the geometry was carried out using the Discover (cvff forcefield, VA09A minimization method) and the Ampac/Mopac (AM1 semiempirical method) modules of the InsightlI program package. These calculations were performed for the whole modifier-substrate system from different starting positions. The adsorption of the modifier-substrate complex onto Pt (111) surface was investigated using the Solids-Docking module of the InsightlI package. This module determines the conformations of the adsorbed molecules by a combined approach of high temperature molecular dynamic simulations with molecular mechanics minimization. All the calculated structures were visualized on a Silicon Graphics workstation.

2.2. Hydrogenation of methyl and ethyl pyruvate The hydrogenation of ehtyl pyurvate (EtPy) was carried out at 23 ~ in a SS autoclave equipped with an injection chamber for separate introduction of the modifier. Cinchonidine (CD) and Troger's base (TB) was used as modifiers. Different batches of EtPy, (Fluka) and Pt/Al203 catalysts (Engelhard E 4759, 5 %w Pt, Dpt = 25 %) were used. Experimental details incliding GC analysis can be found elsewhere [3,12]. The optical yield was calculated as e.e. = (~]-[S])/([R]+[S]). The e.e. values were corrected for the amount of racemic product formed in minor amount in the reactor prior to the injection of CD. 2.3. Kinetic modeling The material balance was calculated for EtPy, ethyl lactates (EtLa) and CD by solving the set of differential equation derived form the reaction scheme. Adam's method was used for the solution of the set of differential equations. The rate constants for the hydrogenation reactions are of pseudo first order. Their value depends on the intrinsic rate constant of the catalytic reaction, the hydrogen pressure, and the adsorption equilibrium constants of all components involved in the hydrogenation. It was assumed that the hydrogen pressure is constant during the kinetic run. Some of the rate constants (kr and kina) were approximated from independent kinetic measurements. The starting values of the differential equations were obtained from the first measured concentration of EtPy and EtLa. The initial ratio of the open and closed forms of the modifier ([CDclose,]o/[CDopen]o) was estimated. The estimation was based on NOESY NMR spectra.

3. THE PRINCIPLE OF CHEMICAL SHIELDING Recently it has been evidenced that a large aromatic substituent, such as naphtyl, can provide a intramolecular steric shielding for an c~-keto ester moiety [9] resulting in enantio-

243 differentiation in the hydrogenation of the 0t-keto group. No ED was observed if the naphtyl ring was substituted for a phenyl one. The ED was attributed to the SE induced by the large aromatic moiety. Similar results were observed in the enantioselective hydrogenation of ethyl pyruvate over Pt/AI203 catalyst in the presence of new types of modifiers [ 10]. In the presence of these new modifiers the ED was completely lost if the naphtyl or quinolyl ring was replaced by phenyl or pyridyl group. It should also be mentioned that in the hydrogenation of ot-keto esters over CD-Pt/AI203 catalysts the ED was partially or fully lost if the quinoline ring of the modifier was partially or fully hydrogenated [11 ]. As emerges from the these results large aromatic substituents might play an important role in the induction of ED. The similarities of above experimental results inspired us to investigate the role of SE in heterogeneous catalytic enantioselective hydrogenation reactions. In heterogeneous catalytic reaction the SE means that a given template molecule interacts with the prochiral substrate in the liquid phase in such a way that one of the prochiral sites is preferentially shielded. If the substrate is shielded then its adsorption onto the metal can take place with its unshielded site resulting in ED. An organic molecule can induce SE if it has (i) an asyrmnetry center (A), (ii) an appropriate bulky functional group 03) for weak interaction with the substrate, (iii) bulky but planar group (C) to induce the steric shielding. If the above requirements are fulfilled the modifier can form a week complex with the substrate. In the above complex the modifier should have an umbrella like conformation with high extent of 't~oncavity" as shown in Figure 1. The role of "concavity" in chemical shielding has been discussed earlier [7,8]. The shielded form of the [substrate - modifier] complex formed in the liquid phase can maintain its entity even after adsorption onto the metal surface. It should be added that based on reaction kinetic data the formation of [substrate - cinchonidine] complex in the liquid phase was suggested in our earlier and recent studies [4, 12]. There are also important requirements for the heterogeneous catalysts: (i) the catalyst should not hinder the formation of the [substrate - modifier] complex, (ii) the modifier should not adsorb irreversible onto the catalyst; (iii) the catalyst should be inactive in the transformation of the modifier into a new derivative, (iv) the catalyst should be resistant towards poisoning by modifier, substrate or product.

4. RESULTS AND DISCUSSION

4.1. Computer modeling The principles of the SE were applied for two enantioselective hydrogenation reactions: (i) hydrogenation of 13-keto esters over Ni-tartrate and (ii) hydrogenation of ct-keto esters over cinchona-PffAl203 catalysts. In this respect the tartaric acid - f3-keto ester system gave a negative result. Neither the substrate nor the modifier have bulky substituents required for SE. The first approach applied for [cinchonidine (CD) - ct-keto ester] complex was also unsuccessful. In the open conformation CD cannot provide the required steric shielding. In open form either the quinuclidine or the quinoline moiety of CD will interact with the substrate. It has already been demonstrated that the quinuclidine moiety has a crucial role both in the rate acceleration and the induction of ED [13].

244 In earlier kinetic and computer modeling [ 1, 2, 14] the open form of CD (CDopen) was used to illustrate the adsorbed [CD - ct-keto ester] complex. In this complex the quinuclidine nitrogen was involved in the interaction with the substrate directly or via a proton bridge. We have modelled the [CDopen - methyl pyruvate] complex. The result is shown in Figure 2. In this complex there is no steric hindrance to prevent the free rotation of the substrate around the quinuclidine nitrogen. Thus, in complex shown in Figure 2. there is no preferential stabilization of the substrate. In earlier computer modeling it was suggested that Pt is involved in the stabilization of the [CDopen--~-keto ester] complex, i.e. the Pt surface prevent the free rotation of the substrate, however the driving force for enantio-differentiation, i.e. for preferential adsorption of the substrate, was not discussed [ 14]. In our second approach the closed form of CD (CDclosed) was used for modeling (see Figure 1). It was found that CD in its closed conformation can provide the concave, umbrellalike form required for steric shielding. The calculated [CDclosed-methyl pyruvate] complex is shown in Figure 3. This complex (complex (R)), after subsequent hydrogenation over Pt should result in (R) -lactate ester. The conformational change of CD from open form to closed one requires the rotation of the quinuclidine ring around the C-(9)- (C4') axis, i.e. to change the torsion angle (C4')-(C9)-(C8-(N1) from 159.42 ~ to 52.32 ~ The energy map for CD was calculated by changing the torsion angles (C3')-(C4')-(C9)-(C8) and (C4')-(C9)-(C8)-(C7). The conformational analysis indicates that CD can exist at least in four different forms. The energy needed to change the conformation of CD from the open form (O1) (it is the crystallographic form) to the most stabile closed one (CI 1) is less than 5 kcal/mole. Figure 4. shows the [CDclosed-methyl pyruvate] complex, in which the substrate is rotated by 180~ The above complex upon hydrogenation will result in (S)-lactate. The major difference between complexes (R) and (S) is the mode of interaction between the loan pair of electrons of the quinuclidine nitrogen and the keto-carbonyl group. In complex (R) the "directionality" [15] of the nucleophilic attack by quinuclidine nitrogen towards the keto carbonyl group is very favourable for the interaction with the keto carbonyl group. The orbital steering theory states [16] that a proper 'reaction window" or "reaction cone" can result in perturbation of the reacting group. We suggest that this perturbation leads to a pronounced rate increase. Thus, in complex (R ) the favourable directionality promotes the perturbation of the keto carbonyl group resulting m the observed rate acceleration. Contrary to that in complex (S) the interacting groups are misaligned. Due to this misalignment no rate acceleration can be expected, i.e. the hydrogenation of (S) complex is not accelerated. Variety of ct-keto esters, such as methyl and ethyl pyruvate, methyl mandalate, dihydro-4,4 - dimethyl-2,3 furanedione were used to calculate the shielded form of [CDclosed- ct-keto ester] complexes leading to the formation of ( R ) or (S) product, respectively. The details of these results will be a subject of a subsequent paper [ 17]. As emerges from these calculations the favourable 'ttirectionality" is maintained in complexes (R), even for dihydro-4,4 - dimethyl2,3 furanedione. Monte-Carlo simulation method was used to investigate the interaction of the [CDclosedMePy] complexes with Pt (111) surface. The result shown in Figure 5 indicates that the shielded complex can maintain its entity even after adsorption. Further computer modeling indicated that there are other molecules with the ability to induce SE. In this respect Troger's bases are of particular interest. The calculated Troger's base-methyl pyruvate complex (R form) is shown in Figure.6.

245

Figure 3. The shielded form of [CDmethyl pyruvate] complex, (R) form

Figure 4. The shielded form of [CDmethyl pyruvate] complex, (S) form

4.2. Hydrogenation experiments Hydrogenation of ethyl pyruvate in the presence of cinchonidine. In our previous studies [3, 4,14] variety of experimental data were obtained, which could not be explained by existing models [1,2] proposed earlier. These results are as follows [3,4,12]: (i) the monotonic increase type behaviour of the optical yield - conversion dependencies, (ii) the complexity of the reaction kinetics, (iii) side reactions catalyzed by CD. It was also demonstrated that the enantio-differentiation can be induced if the modifier is injected into the reactor during racemic hydrogenation. Upon injection of CD into the reactor during racemic hydrogenation the rate acceleration was always instantaneous, while the optical yield vs. conversion dependencies showed a monotonic increase type behaviour as seen in Figure 7. In acetic acid the increase part of the above dependence is so fast that it hardly can be followed by our sampling technique. At low concentration of modifier the optical yield passes through a maximum. In this case the

246 decreasing part is due to the transformation of CD during the hydrogenation reaction. Parallel formation o f CD derivatives with saturated quinoline ring was evidenced by thin layer chromatography.

Figure 7. Typical optical yield - conversion dependencies. [CD]o, M: O - 6.8 x 10.6 9 - 3.40 x 10 5, 9 - 0.8 x 10 -4 + 5.0 M AcOH.

Figure 8. The shielded form o f [ C D methyl pyruvatesyn] complex, ( R ) form

We have compared, the rate acceleration effect induced either by the CD and different moieties originated from CD, i.e quinuclidine and quinoline. These experiments were carried out in ethanol. If the relative rate of racemic hydrogenation is equal to one the following relative rates has been measured: quinoline = 2, quinuclidine = 3, cinchonidine = 40. In the presence of quinoline a short induction period was needed to observe the rate acceleration. It is suggested that during this period quinoline was partly hydrogenated. Other tertiary nitrogen bases, such as triethylamine, triethylenediamine, etc. resulted also rate acceleration with relative rate = 2-4.

247 The above comparison indicates that the rate acceleration induced by CD is more pronounced than that of the other tertiary nitrogen bases. This fact also indicates that in CD a cooperative effect should exist between the quinuclidine nitrogen and the quinoline ring. The

cooperative effect is in force if the modifier is in a shielded form. One of the most interesting side reactions taking place during the enantioselective hydrogenation is the transesterification of the substrate or the reaction product. If the enantioselective hydrogenation of ethyl pyruvate was performed in methanol as a solvent the formation of methyl pyruvate and methyl lactate was observed. CD appeared to be an effective catalyst for the above transesterification reaction. The transesterification reaction can be attributed to the perturbation of the ester carbonyl group in the [CDclosed-substrate] complex. The possibility of this side reaction was predicted by earlier quantum-chemical calculations [ 18]. These results indicated that the reaction pocket in methyl pyruvate for the nucleophilic attack is situated between the two carbonyl groups, i.e. both carbonyl groups can be perturbed by a nucleophile provided both carbonyl groups have the fight "directionality". However, the fight "directionality" for both carbonyl groups can be obtained if they are in syn position. The conformational analysis of methyl pyruvate shows that it can have two conformers. In the second conformer the two carbonyls are in syn position. The anti-syn conformational change requires 3 kcal. The [CDclosed- methyl pyruvatesyn] complex ((R) form) was also calculated and shown in Figure 8. In the above complex the "directionality" of the lone pair of electrons of the quinuclidine nitrogen is advantageous for interactions with both the keto and the ester carbonyl groups. Table 1. Hydrogenation of ethyl puruvate in the presence of Troger's base No

Pressure, bar

Modifier .....

Acid added, M

Rate contant min "1

Conv. '~" %

Optical '~* yield,

1. 2. 3. 4. 5. 6. 7.

50 50 50 50 10 50 50

TB TB TB TB CD* CD*

AcOH, 5.0 AcOH 0.5 TFAc, 0.5 AcOH, 5.0 AcOH, 5.0 AcOH, 5.0

0.0041 0.0075 0.0057 0.0050 0.0045 0.0560 0.1320

10.9 12.5

0.088

21.2

0.242

11.2 12.9 97.9 99.5

0.383 0.248 0.860 0.931

Solvent: Toluene, T = 10 ~ * measured at 23 ~

TFAc- trifluor acetic acid,

** Conversions and optical yields measured after 60 minutes.

H y d r o g e n a t i o n of ethyl puruvate in the presence of Trogers base. It has been demonstrated by computer modeling that other organic molecules, such as Troger's base (TB) can induce ED in the hydrogenation of ethyl pyruvate over Pt/AI203. The results obtained in the presence of TB are summarized in Table 1. TB results in ED only in the presence of acetic acid. This modifier, contrary to the cinchona alkaloids, did not result in rate acceleration and the enantio-differentiation was moderate.The main difference between CD and TB is that in the [TB - substrate] complex there is no interaction between the tertiary nitrogen and the keto

248 carbonyl group. Consequently, there is no pronounced rate acceleration and the optical yield is relatively low. The relatively low optical yield can be attributed to the small size of the shielding group. Further experiments, including synthetic works and kinetic measurements are in progress to optimize the optical yield induced by Troger's bases. 4.3. Kinetic Modeling The shielding effect model suggest that both the rate acceleration and the induction of ED is attributed to the formation of a shielded [substrate-modifier] complex in the liquid phase (reaction (1)). In this complex CD is in closed form. The above complex, referred as [X]cl upon hydrogenation will result in (R) lactate (reaction (2)). Reactions (1) and (2) strongly resemble the corresponding steps in enzyme catalyzed reactions. The formation of complexes [X] takes place in an equilibrium reaction. The corresponding unshielded [substrate-modifier] complex, [X]closed can also be hydrogenated, however it gives racemic product similar to the free substrate, see reactions (4) and (5), respectively. The rate of reaction (2) is much higher than that of reactions (4) and (5). The equilibrium reactions for the conformational changes of CD and for the transformation of [X]close d into [X]open are also included into the reaction scheme ( reactions (6) and (7), respectively). The hydrogenation of both forms of CD is also taken into account. This reaction, in which the quinoline ring of CD is hydrogenated is responsible for the decrease of the optical yield with conversion observed at very low concentration of CD. The deactivated form of CD formed in the above reaction is referred as

CDma. The simplified reaction scheme for the enantioselective hydrogenation of ot-keto esters over cinchona-Pt/Al20 3 catalyst can be written as follows: Substrate

+

[CD]cl

<---kl/k. 1--->

[X] cl

+

H2/Pt

--k* .... >

Substrate

+

[CD]o p

<---k2/k.2--->

[X]op

+

H2/Pt

-k*r---> (R)-lactate + CDop

(4a)

[X]op

+

H2/Pt

--k'----> (S)-lactate + CDop

(4b)

Substrate

+

H2~t

-kR.... > (R)-lactate

(Sa)

Substrate

+

H2/Pt

-ks .... >

[X]c I

(1)

(R)-lactate + [CD]ci

(2)

[X]o p

(S)-lactate

(3)

kR=ks=kr

(Sb)

[CD]op

<---kcD/kcD-I--->

[CD]cl

(6)

[x]~l

<---kx/k_x....... >.

[X]op

(7)

[CD]cl

+

H2/Pt

--kina...... >

[CD]ma

(8a)

[CD]o p

+

H2/Pt

--kina...... >

[CD]ina

(8b)

In this scheme, due to the rate acceleration effect, the enantioselective hydrogenation is much faster than the two racemic hydrogenation reactions (k* > kr, k* > k'r). Please note that the rate constants for the hydrogenation reactions of are pseudo fist order, which contains in a

249 certain form the intrinsic rate constant and the adsorption equilibrium constants of all components involved in the hydrogenation. More detailed reaction network will be needed, which will take into account all of the adsorbed forms of substrate, modifier, reaction product and by-products. Two kinetic experiments with different CD concentrations were used for kinetic modeling. In this simulation all of the rate constants not involved in the hydrogenation step were not altered. The calculated and simulated kinetic curves and optical yield-conversion dependencies are shown in Figure 9a and 9b. The results of kinetic modeling indicates that the whole kinetic curve and the optical yield - conversion dependencies can be well described by a kinetic model derived from the shielding effect model.

.0

-

'

1.0

'-"- '-

i

0.8"

0.8

~, 0.6

"~, 0.6

g 0.4

=.0.4

0.2

0.2 1

0.

'

0

20

40

60

80

Time [min]

100

120

.

l

.

l

,

1

0.0~

0.0

,

'

0.2

0.4

0.6

0.8

1.0

Conversion

Figure 9. Results of kinetic modeling, a. Conversion vs. time, b. Optical yield vs. conversioia dependencies. I I - [CD]o = 6.8 x 10-6 M, 9 - 3.4 x 10-5 M, [Substrate]o = 0. 98 M., [Rlactate]o = 0.01, [S-lactate]o = 0.01 M. [CDopen]o/[CDelosed]o = 9/2. The continuous curves were obtained by kinetic modeling.

CONCLUSIONS A new model is proposed to undestand the origin of ED over heterogeneous catalyts. The model was applied for the enantioselective hydrogenation of ct-keto esters. The shielding effect model is based on steric shielding provided by a large aromatic ring. The new model can explain all of the observations, which could not be interpreted by existing models [ 1,2]. It is the first working model, which can explain (i) the instantaneous rate acceleration, (ii) the monotonic increase character of the optical yield - conversion dependencies and (iii) the appearence of the maximum in the optical yield - conversion dependencies. The shielding effect model can explain the functional behaviour of different parts of the cinchonidine molecule. This model can be used to design new modifiers for the enantioselective hydrogenation of otketoester. The computer modeling with molecular mechanics calculations appeared as a powerful tool to give qualitative explanations for the modifer-susbtrate interactions taking

250 place in the liquid phase. The next step should be to give quantitative data for the above interactions. These studies are in progress in our laboratories.

ACKNOWLEDGMENT Financial support given by OTKA (Grant No: T1801 and T4340) is greatly acknowledged. REFERENCES

1. I.M. Sutherland, A. Ibbotson, R.B. Moyes and P.B.J. Wells, J. Catal., 125, (1990) 77. 2. M. Garland and H.U. Blaser, J.Am.Chem. Soc., 112, (1990) 7048. 3. J.L. Margitfalvi, B. Minder, E. T/das, L. Botz and A. Baiker, New Frontiers in Catalysis, (Proc. 10th Int. Cong. Catal. Budapest, July 1992), Guczi, L. et al. (eds), Elsevier, Amsterdam (1993) 2471. 4. J.L. Margitfalvi, Chem. Ind. (Marcel Dekker), 62 (Catal. Org. React., Scaros, M.G., Pmnier, M.L. (eds)), (1995) 189. 5. Y. Izumi, Adv. Catal., 32 (19830 215. 6. A.Tai and T.Harada, Taylored Catalysts, Y. Iwasawa (ed.), D.Reidel, Dordrecht, (1986) 265. 7. H. Bushman, H.D. Scharf, N. Hoffmann and P. Esser, Angew. Chem. Int. Ed. Engl., 30 (1991)477. 8. G. Helmchen and R. Schiere, Angew.Chem. Int. Ed. Eng. 22 (183) 237. 9. U.Maitra and P. Mathivanan, Tetrahedron: Asymmetry, 5 (1994) 1171. 10. K.E. Simons, G. Wang, T. Heinz, T. CJiger, T. Mallat, A. Pfaltz and A. Baiker, Tetrahedron: Asymmetry, 6 (1995) 505. 11. J.L. Margitfalvi, P. Marti, A. Baiker, L. Botz and O. Sticher, Catal. Lett., 6, (1990) 281. 12. O. Schwalm, B. Minder, B., J. Weber, and A. Baiker, Catal. Lett. 23 (1994) 245. 13. J.L. Margitfalvi J. and M. Hegedus, 8th International Congress on the Relation Between Homogeneous and Heterogeneous Catalysis, Balatonfured, Hungary, Sept. 1995. N~ oral presentation; accepted for publication in J. Mol.Cat. 14. H.U. Blaser and M. Muller, stud. Surf. Sci. Catal. 59 (1991) 73. 15. F. Manger, Tetrahedron, 39 (1983) 1013. 16. D.R. Storm and D.E. Koshland, Jr., J.Am.Chem.Soc. 94 (1972) 5805. 17. E. Tfirst and J.L. Margitfalvi, to be published. 18. O. Schwalm, J. Weber, J. Margitfalvi and A. Baiker, J. Mol. Structure, 297 (1993) 285.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

251

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

R a c e m i z a t i o n of ( l S ) - ( - ) - e x o - 2 , 4 - D i d e u t e r o a p o p i n e n e o v e r Pd" E v i d e n c e for an I n t r a m o l e c u l a r 1 , 3 - D e u t e r i u m Shift Gerard V. Smith a'b, Boris Rihter a, Agnes Zsigmond b'c, Ferenc Notheisz b'c, and Mih~ly Bart6k c aDepartment of Chemistry, bMolecular Science Program, Southern Illinois University, Carbondale IL 62901, United States of America CDepartment of Organic Chemistry, J6zsef Attila University, D6m t6r 8, 6720 Szeged, Hungary

1. INTRODUCTION The hydrogenation and isomerization of alkenes can usually be described by the classical Horiuti-Polfinyi mechanism. According to that mechanism, in a deuterium atmosphere, double bond migration incorporates deuterium into the allylic position.

Scheme 1: Horiuti-Poldnyi (classical) mechanism for double bond migration. /

, i

D2

.4 ,

/!

D

a

Early experiments [1 ] using (+)-apopinene and deuterium showed, however, that in the isomerized molecules the deuterium content was very low and the isomerization was much faster than deuterium incorporation into the allylic position. Therefore it seemed probable that isomerization takes place through an intramolecular hydrogen shift. A sigmatropic 1,3hydrogen shift was suggested, in which the allylic e n d o - H shifted (top shift) [2].

Scheme 2: 1,3-top side hydrogen shift

252 Although this type of reaction is symmetry forbidden in an unadsorbed molecule, theoretical calculations showed that in a molecule adsorbed on transition metals, such a shift is allowed [3-5]. Later, other theoretical calculations suggested another type of 1,3-hydrogen shift, one in which the allylic exo-hydrogen is abstracted by the surface from an adsorbed alkene (either 1,2-diadsorbed or 7r-complexed) and the resulting x-allyl species moves over the abstracted hydrogen in such a way that it adds to the former vinylic position and causes, in effect, a stepwise intramoleeular 1,3-hydrogen shift (bottom shift) [6].

Scheme 3: ~-allyl shift mechanism for double bond migration

To distinguish between these two hydrogen shift mechanisms, (1S)-(-)-exo-2,4dideuteroapopinene was constructed as a probe molecule. A top shift of the allylic endo-H will not affect the deuterium content of the molecule and no change should occur in the hydrogen content at any position, but a bottom shift of the allylic exo-D will decrease deuterium in the vinylic position (C2) and increase deuterium in the allylic position (C4).

Scheme 4: Comparison of topside and bottom side hydrogen shifts

H

either classical or ~-allylic shift

ki

D

~

ki

(+)-2d,4d-ap~

D

topside suprafacial sigmatropicshift

2. E X P E R I M E N T A L By a ten-step route from ct-pinene [7], (1S)-(-)-exo-2,4-dideutero-apopinene was synthesized thrice with different deuterium concentrations. The hydrogen contents (19~ 15.1%, and 6.1%) of the molecules at the C2 position were determined by 200 MHz proton

Scheme 5: Numbering in apopinenes 5

4

1

2

(-)-apopinene

253 N M R with the C3 proton used as an intemal standard. Determination of hydrogen content at C4 was not possible on the 200 MHz instrument due to overlapping peaks. Double bond migration (racemization) within (1S)-(-)-exo-2,4-dideutero-apopinene was studied on Pd-black, 0.46% Pd/SiO2, 1.17% Pd/SiO2, 0.4% Pd/AI203, 1.0% Pd/A1203, and PdsoSi20 metallic glass catalysts in both deuterium and hydrogen. Hydrogenations were run on neat apopinene (except where noted) in the liquid phase at one atmosphere of hydrogen or deuterium and were stopped at different percentages to furnish several different percentages of racemization. In the recovered alkenes from these reactions, the hydrogen contents at C2 were determined. The apparatus and the methods for studying the hydrogenaton and for measuring percent dispersion (%D) of the catalysts by hydrogen chemisorption have been described earlier [8].

3. R E S U L T S AND DISCUSSION Essentially, the same results are found for reactions in deuterium and hydrogen (Table 1). As double bond migration (racemization) proceeds, the hydrogen content at C2 increases.

Table 1. The racemization of (1S)-(-)-exo-2,4-dideuteroapopineneon Pd catalysts in deuterium catalysts

racemization, %

0.46% Pd/SiO2 (60.5%D) 38.9 81.1 48.9 56.7 PdsoSi2ometallic glass 10.6 1% Pd/Al:O3 (49.0%D) 30.8 0.4% Pd/Al203 (77.5%D)~ 10. l

symbol + + + *

Pdso-Si20metallic glass

32.2 39.6 1.17% Pd/SiO2 (40%D)@ 33.0

PdsoSi2ometallic glass 62.9 0.46% Pd/SiO2 (60.5%D) 68.0 Pd-black 57.3

,I, x x

% H content in C2 original % final % 19 19 19 19 19 19 19

36.3 50 38.6 41.4 in hydrogen 11.6 in hydrogen 31.9 22.6

15.1 15.1 15.1

33.5 34.4 27.2

6.1 6.1 6.1

35.5 36.1 32.3

poisoned with carbon tetrachloride poisoned with carbon disulfide

The results are easier to see in Figure 1 where representative data are plotted. Dotted lines indicate the change in hydrogen content expected to occur at C2 for a 1,3-bottom shift in a molecule with the particular starting hydrogen content in C2. The solid lines represent the expected absence of change for a 1,3-top shift, and the symbols are those shown in Table 1.

254

60

H content on C2, % 1,3-bottom shift

50 40

..A

t

30 9

.

s

20

1 s

t

1,3-top shift

i

10 0

0

I

1

1

I

1

1

1

1

1

10

20

30

40

50

60

70

80

90

100

racemization, %

Figure 1. Racemization of (1S)-(-)-exo-2,4-dideuteroapopinene on different Pd catalysts.

Although the experiments are imperfect because the percentages of hydrogen at C4 are not known, little evidence exists for the 1,3-top shift. With the exception of the Pds0Si20 metallic glass experiment in hydrogen (fifth experiment from top in Table 1), all data fall near the lines calculated for a 1,3-bottom shift for the respective original percentages of hydrogen at C2. Earlier, with unlabeled apopinenes, we observed that double bond migration appeared to be occurring in deuterium atmospheres without deuterium incorporation into the aUylic position. This led us to postulate an intrarnolecular top shift of hydrogen because a bottom shifting hydrogen should become diluted by surface deuterium if it spends time on the surface [1]. However, these present data suggest a bottom shift is occurring without dilution from surface hydrogen or deuterium and raises several interesting questions. Is the shifting hydrogen or deuterium isolated or protected from the surface hydrogendeuteriurn pool? And how fast does the hydrogen-deuterium pool equilibrate over the surface? To partly answer these questions we consider apopinene. Apopinene is a unique molecule and not representative of most alkenes. Over Pt apopinene undergoes double bond migration approximately eight times faster than cyclohexene [9] and over Pd it undergoes double bond migration as much as seventeen times as fast as addition [Sb]. Assuming the shifting hydrogen or deuterium is not protected from the surface hydrogen-deuterium pool, these results suggest that the rapid double bond migration occurs faster than migration of surface hydrogen or deuterium. It places an upper limit on the rate of migration of surface hydrogen. Apopinene and similar molecules may furnish a unique way to measure rates of migration of hydrogen and deuterium on surfaces under working conditions of catalysts.

255 On the other hand, is the shifting hydrogen protected from incursion by surface hydrogen or deuterium because of some feature of the shift? For example, the shift could occur on the surface but inside the ring such that it is inhibited from migrating outside the reaction sphere (see Scheme 3). In such a case, outside hydrogens or deuteriums would also be inhibited from migrating into the reaction sphere and diluting the shifting hydrogen or deuterium. To answer these and other questions we are continuing to construct likely probe molecules.

4. ACKNOWLEDGEMENT Financial Support from OTKA (1885/91 and 4182/92), NSF INT-8403357, and the USHungarian Joint Fund (No. 177) is gratefully acknowledged. The authors thank Daniel Ostgard for preparing and characterizing some of the catalysts.

REFERENCES

G.V. Smith and D. Desai, Ann. N. Y. Acad. Sci., 214 (1973) 20. G.V. Smith and J. R. Swoap, J. Org. Chem., 31 (1966) 3904. F.D. Mango, Adv. Catal., 20 (1969) 291. F.D. Mango, Coord. Chem. Rev., 15 (1975) 109. A.B. Anderson, J. Chem. Phys., 63 (1975) 4430. A.B. Anderson, D. B. Kang and Y. Kim, J. Am. Chem. Soc., 106 (1984) 6597. B.D. Rihter, Ph.D. Dissertation, Southern Ill. University at Carbondale, 1985. Current address: Aldrich Chemical Co, 230, South Ember Lane, Milwaukee WI. 53233 8. a.G.V. Smith, .~. Molnar, M. M. Khan, D. Ostgard, and T. Yoshida, J. Catal., 98 (1986) 502; b. G. V. Smith, D. Ostgard, M. Bartok, and F. Notheisz, Catalysis of Organic Reactions, (P. N. Rylander, H. Greenfield, and R. L. Augustine, Eds.) Marcel Dekker, Inc., 1988, p 409. 9. G.V. Smith, J. A. Roth, D. S. Desai, and J. L. Kosco, J. Catal.,30, 79 (1973)

1. 2. 3. 4. 5. 6. 7.

This Page Intentionally Left Blank On the other hand, is the shifting

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (F_xls.) 1 l th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

257

F e / M g O catalysts for the selective h y d r o g e n a t i o n of nitriles Gary Bond* and Frank S. Stone School of Chemistry, University of Bath, Bath BA2 7AY, United Kingdom

Fe/MgO catalysts with 5 to 30 mol % Fe have been prepared by impregnation and coprecipitation. Their reducibility has been measured and a comparison made of their Fe ~ surface areas. Catalysts prepared via coprecipitation yielded larger iron areas than those via impregnation. The activity and selectivity of the reduced catalysts for the hydrogenation of propanenitrile at 20-30 bar and 473 K and of ethanenitrile at 1 bar and 508 K have been determined. The most active catalysts are those prepared by coprecipitation and they show high selectivity for primary amines. The activity for ethanenitrile hydrogenation correlates with the iron surface area.

1. INTRODUCTION The catalytic hydrogenation of nitriles is a widely-used method for the industrial production of amines. The reaction, which is catalyzed by a variety of transition metals, also has fundamental interest in that the nature of the metal is known to be a major parameter in influencing the resulting proportions of primary, secondary and tertiary amines. Platinum-based catalysts, for example, direct the hydrogenation in favour of tertiary amines, but among the catalysts most active and selective towards primary amines (often the most desired products industrially) are nickel and cobalt. These metals, and especially nickel, have been frequently studied for the hydrogenation of both low and high molecular weight nitriles il-3], but iron, adjacent to cobalt in the 3d series, has received little attention [4]. Iron, however, has merits in that in comparison with nickel and cobalt it is less prone to catalyze hydrogenolysis and deposit coke. In the industrial context it is also attractive economically. It was accordingly chosen as the catalyst metal for the present investigation. Since the hydrogenation is a metal-catalyzed reaction, it is appropriate to use an oxide support to enhance the dispersion. However, the support, like the metal, needs to be chosen with the desired selectivity in mind. The early view [5] that the selectivity in nitrile hydrogenation is determined largely by the behaviour of the partially-hydrogenated intermediate, the imine R-CH = NH, which can either accept two further hydrogens to form the primary amine or can react with an already-formed amine to start a sequence which * Present address: Department of Chemistry, University of Central Lancashire, Preston, PR1 2HE, United Kingdom

258 leads to the secondary or tertiary amine is still accepted. This latter sequence involves condensation reactions and as such is favoured by acidic conditions. Thus to enhance selectivity towards primary amines, acidic supports are best avoided [2]. Indeed, a basic oxide support is attractive for the added reason that the product amine molecules involved in the formation from the imine of the secondary and tertiary amines, being themselves basic, are less likely to remain adsorbed on the catalyst and be available for onward reaction than would be the case with an acidic oxide. With this in mind, a logical choice for the oxide support for iron in the present work was magnesium oxide, pre-eminent as a basic oxide and readily preparable in high surface area form. Supported iron catalysts are notoriously difficult to reduce [6-8] and thus a substantial fraction of the iron can be expected to remain inactive for the catalysis of hydrogenation. Particular attention has therefore been paid to the preparation of Fe/MgO catalysts by several different methods and examination of their effectiveness in producing metallic iron of adequate specific surface area after reduction in hydrogen. The activity and selectivity for primary amine formation have been determined for the hydrogenation of ethanenitrile (acetonitrile) and propanenitrile.

2. EXPERIMENTAL 2.1 Materials Fe/MgO catalysts with loadings in the range 5-30 mol % Fe were prepared by four methods (a-d) as follows: (a) Impregnation (IP). This method was similar to those used by Boudart et al. [6] and Topsoe et al. [91. A slurry of basic Mg carbonate (Fluka) in 160 ml of distilled water was heated with stirring to 340 K and 160 ml of an aqueous solution of ferric nitrate (BDH) preheated to 340 K was rapidly added. After stirring for 30 min at 340 K the impregnated slurry was filtered, washed and vacuum-dried at 360 K before being calcined in air at 773 K. Different iron loadings were obtained by varying the concentration of the Fe(NO3) 3 solution. The calcined precursor was reduced in a static system (see Sec. 2.2) in hydrogen at ca. 30 Torr for 24 h at 553 K, 24 h at 623 K and 24 h at 693 K, the water vapour produced being frozen out in a 77 K trap. The resulting catalysts are designated as IPx, where x = 100 tool fraction Fe/(Fe + Mg). (b) Carbonate coprecipitation (CCP). Solutions of Fe(NO3) 3 and Mg(NO3) 2 (BDH), each 1M, were mixed in appropriate ratios and heated to 360 K. 1M (NH4)2CO 3 solution was added dropwise until precipitation was complete. After washing and filtering, the precipitate was dried at 373 K and calcined at 773 K. This precursor was reduced as for (a) above. These catalysts are designated CCPx, where x is the mol % Fe as before. (c) Hydroxide coprecipitation (OH). A mixture of Fe(NO3) 3 and Mg(NO3) 2 solutions as in (b) was in this case added dropwise into a stirred vessel containing ammonia solution at 303 K. The pH was maintained constant at pH 11 by further addition of ammonia. The resulting precipitate was washed, dried and calcined as in (b) above, and was reduced as in (a). These catalysts are designated OHx, where x denotes the mol % Fe as previously.

259 (d) Hydrotalcite method (HT).

100 ml 1M Fe(NO3) 3 and 300 ml 1M Mg(NO3)2 was

pumped simultaneously with 400 ml of 0.125M (NH4)2CO3 into a stirred reaction vessel at a rate of 70 ml h -1. The temperature of the vessel was kept constant at 323 K and the pH was maintained constant at 9.8 by the addition of NH4OH. The resulting precipitate of magnesium iron hydrotalcite (pyroaurite, nominally Mg6Fe2(OH)16CO3) was dried, calcined and reduced as in (c) above. This catalyst is designated HT. 2.2 Catalyst characterization The reduction procedure described above was carried out in a static system in order to facilitate determination of the extent of reduction in situ. The reduction was monitored gravimetrically using a microbalance (CI Electronics MK II) and volumetrically by measuring the decrease in hydrogen pressure. Total surface areas were determined by the BET method using nitrogen at 77 K. Carbon monoxide chemisorption was used to estimate the surface area of metallic iron after reduction. The quantity of CO chemisorbed was determined [6] by taking the difference between the volumes adsorbed in two isotherms at 195 K where there had been an intervening evacuation for at least 30 min to remove the physical adsorption. Whilst aware of its arbitrariness, we have followed earlier workers [6,10,11] in assuming a stoichiometry of Fe:CO = 2.1 to estimate and compare the surface areas of metallic iron in our catalysts. As a second index for this comparison we used reactive N20 adsorption, N20(g) ~ N2(g) + O(ads), the method widely applied for supported copper [12]. However, in view of the greater reactivity of iron, measurements were made at ambient temperature and p = 20 Torr, using a static system. Bulk characterization of calcined precursors and reduced catalysts was carried out by X-ray diffractometry using Cu Kcx radiation. Reduced catalysts were first passivated by exposure to N20 as described above. Line-broadening analysis was carried out on the Fe(110) reflection to obtain the iron particle size. Overlap with the MgO(200) reflection limited its usefulness to the more highly-loaded catalysts. 2.3 Catalysis Hydrogenation tests were carried out using an autoclave reactor (ICI, Wilton, UK) in the case of propanenitrile and a flow reactor operating at atmospheric pressure for ethanenitrile. (a) Autoclave reactor experiments. The procedure was as follows. 0.5 g of passivated catalyst was activated by reduction in H 2 at 573 K and after transfer via a nitrogen-purged glove box was installed in the reactor with 10 ml propanenitrile. The reactor was pressurized with H 2 and heated to 473 K. The subsequent fall in pressure from a maximum of about 30 bar (20% nitrile : 80% hydrogen) as hydrogenation proceeded was monitored for 20 hours. The liquid product was subjected to GLC analysis. (b) Flow reactor experiments. For a more precise comparison of the activity and selectivity of the Fe/MgO catalysts, the hydrogenation of ethanenitrile was investigated. A laboratory flow reactor operating at 1 bar pressure was employed. A sample (0.5 g) of passivated catalyst was loaded into the reactor and after activation (as above) the temperature was lowered to 508 K (unless otherwise indicated) and the H 2 flow (at a chosen

260 value in the range 10-50 ml min -1) was diverted through a saturator containing ethanenitrile held normally at 273 K before passing over the catalyst. After achieving steady state conditions the exit gas was sampled and analyzed by GLC. Space velocity, temperature and the respective concentrations of CH3CN and H 2 in the feed were separately varied in order to obtain the rate of of reaction, activation energy and reaction orders in ethanenitrile and hydrogen for individual catalysts.

3. RESULTS 3.1 Characterization of catalyst precursors Four catalysts with nominal compositions of 5, 10, 20 and 30 mol % Fe were prepared by each of the methods (a), (b) and (c) described in Sec. 2.1, together with one by method (d). Actual compositions were close to the nominal ones (see later). in the OH series, two phases were detectable by XRD in the dried precipitate. One was a phase with the pyroaurite structure, carbonate having presumably arisen from atmospheric CO 2, and the other brucite, Mg(OH)2 , to which pyroaurite is closely related structurally. For both the CCP and IP series, the only structure identifiable in the dried precipitate was that of magnesium hydroxycarbonate. X-ray analysis of the calcined precursors showed MgO together with y-Fe20 3 in the case of the OH series and HT, but ~-Fe20 3 with the CCP and IP series. MgFe20 4 spinel was also detectable in some cases. 3.2 Characterization of the reduced catalysts Iron was present as Fe 3+ in the calcined precursors. For all the catalysts the reduction procedure described in Sec. 2.1 resulted in incomplete reduction of the Fe 3+ to metallic iron. This is in agreement with the findings of previous authors [6,11 ]. The individual percentage reductions of Fe 3+ to Fe ~ as determined by the separate gravimetric and volumetric measurements (See. 2.2), are shown in Table 1. The values are calculated on the assumption that all the Fe 3+ is reduced to F e 2 + prior to the onset of reduction to Fe ~ There is good agreement between the two methods. Table 1 also records the actual Fe/(Fe + Mg) ratio in the catalysts as determined by atomic absorption spectroscopy (AAS) on the calcined precursors. We present next the data which relate to the surface area of the reduced catalysts. All the catalysts had a total (BET) surface area in the range 50-100 m2g -1, the lower values being for those with the highest content of iron, but the more important catalyst characterization parameter for the present work is the surface area of the metallic iron, studied here by CO chemisorption and N20 reaction. The CO chemisorption was evaluated using the range of the adsorption isotherm between 75 and 150 Torr, where the difference in volume adsorbed between the first (chemisorption and physical adsorption) and the second (physical adsorption) isotherms was constant. Expressing this volume as a number of molecules and assuming 2 Fe o atoms per chemisorbed CO, followed by conversion to surface area on the basis of 1.2 x 1019 atoms Fe per m 2 (an average for the (110), (100) and (111) faces ofbcc Fe), we obtained the Fe ~ surface areas shown in Table 2. Also shown in this table are

261 values derived from N20 reaction, where the amount of N20 decomposed was determined gravimetrically with the microbalance and where a stoichiometry of 2.5 molecules N20 reacted per surface Fe~ was assumed. This value supposes a lesser oxidation depth than the "approximately 4 monolayers of metallic iron oxidized upon prolonged exposure to N20 at room temperature (and 280 Tort)" reported by Vogler et al. [ 13], but we used a pressure of only 20 Torr to drive the reaction. Fe:N20 stoichiometry of 1:2.5 gives good agreement with the CO data. However, more significant is the agreement of the variation in Fe ~ surface area within the group of 13 catalysts, taking the data from CO chemisorption on the one hand and comparing them against the data from N20 reaction on the other. Table 1 Effect of loading and method of preparation on the reducibility of iron in Fe/MgO catalysts Reduction to Fe ~ (%) Catalyst

Fe/(Fe + Mg) by AAS analysis

Gravimetric determination

Volumetric determination

IP5 IP10 IP20 IP30

4.9 10.0 19.8 29.9

24 32 37 53

24 31 35 49

CCP5 CCP10 CCP20 CCP30

4.8 10.0 18.9 30.8

20 26 59 69

22 25 54 65

OH5 OH10 OH20 OH30

4.8 9.5 18.5 29.2

27 39 54 57

26 37 51 54

HT

22.6

42

40

For the catalysts with highest iron loadings, the average size of the iron crystallites (d) was obtained from the line-broadening of the Fe(ll0) reflection, the instrumental broadening having been determined with a sample of sintered iron powder (Koch-Light). The results (Table 3) show that the OH catalysts have the smallest particle sizes. An estimate of the average particle size can also be made [6] from the relation d(nm) = 0.85/D, where D is the dispersion, the ratio of surface Fe ~ atoms to total Fe ~ atoms, the

262 Table 2 Iron surface areas from CO chemisorption and N20 reaction Fe surface area (m2g-leat) Catalyst by CO chemisorption

by N20 reaction

IP5 IP10 IP20 IP30

1.5 2.1 3.4 5.6

1.3 2.5 3.1 5.5

CCP5 CCP10 CCP20 CCP30

2.9 4.4 5.0 5.3

3.1 4.7 5.1 5.7

OH5 OH10 OH20 OH30

5.0 9.5 10.1 7.2

5.1 9.8 10.2 7.0

7.5

7.4

HT

Table 3 Iron particle size from X-ray line broadening and percentage-reduction/CO-chemisorption

Iron particle size (nm) Catalyst X-ray line broadening

reduction and CO chemisorption

IP20 IP30

47 48

28 33

CCP20 CCP30

45 46

29 47

OH20 OH30

39 36

13 27

263 former being given by the CO chemisorption results and the latter by the loading and percentage reduction data. The d values obtained in this way are also shown in Table 3. 3.3 Hydrogenation of propanenitrile The catalysts were studied for their activity and selectivity in propanenitrile hydrogenation in simple batch experiments at 473 K using the autoclave reactor. A rough comparison of activity was made by taking as a measure of the rate of hydrogenation the slope of the tangent to the curve of H 2 uptake versus time at p = 25 bar. The results showed clearly that the activity of the catalysts prepared by impregnation (IP) was less than that of those prepared by coprecipitation. The selectivity was obtained by carrying out GLC analysis of the products after 20 hours of reaction. It was found that all the catalysts prepared by coprecipitation gave selectivities to the primary amine in excess of 60%, whereas the IP catalysts were again different, giving a selectivity of less than 60% whatever the loading. An informative overview of the activity and selectivity results is obtained by presenting them as a matrix (Fig. 1), which serves also to show the general trend of higher selectivity being associated with higher activity (the arrow at OH10 in Fig.1 is to indicate that it has an off-scale rate, 0.2 bar rain-l). The liquid products other than the primary amine were the secondary and tertiary amines, the latter in very small yield (--1%), 1-propylamino-propene (the most abundant by-product) and 1,1-dipropylamino-propene (~1%). The gas remaining after cooling the reaction mixture was sampled with a gas syringe in a few cases (reactions using OH catalysts) to test whether any hydrogenolysis had occurred. The results indicated that about 4% of the nitrile which had reacted had undergone hydrogenolysis: GC analysis showed that the principal gaseous products from this hydrogenolysis were propane (50%), propene (35 %) and ethane (10%). 0.04

100

8O" ~--

OH5 OH3 0 O t ccP20 ~

. 9

HT 9CCP30

l

OH10

-~" 0 . 0 3

OH20

.,.

E "7 -- 0 . 0 2

. . . . . . C C P 10 .--IP30 > --IP20 o 9

60":e

OH3

0

E

40-

=: 0 . 0 1

20 9 IP5 0

0

9

........ I ......... t 0.05 O. 10

l 0.15

Rate (bar min-1) Fig. 1. C2HsCN hydrogenation at 473 K. Selectivity to C2H5CH2NH 2 vs. activity.

--

0

............

1. . . . . .

5

I

10

Fe ~ area (m2g- 1 ) Fig.2. CH3CN hydrogenation at 508 K. Activity vs. Fe~ surface area.

264 3.4 Hydrogenation of ethanenitrile The aim of the experiments with ethanenitrile (CH3CN) was to study hydrogenation activity more quantitatively than was possible with the autoclave reactor. Catalysts IP30, OH5, OH20, OH30 and HT were selected for investigation. Steady-state conversion to ethylamine was measured at 508 K and atmospheric pressure for various space velocities, the hydrogen flow rate being varied from 10 to 50 ml rain -1. Plots of conversion vs. reciprocal space velocity were linear. The respective gradients, which are measures of the reaction rates for hydrogenation, are plotted in Fig.2 as a function of the iron surface areas from Table 2. The rates correlate with the surface areas: note particularly the low rate for the IP catalyst, in spite of its high loading. The selectivity for hydrogenation to the primary amine was high in all cases, with only minor amounts of secondary amine and no tertiary amine. There was only slight evidence of hydrogenolysis. The activation energy of the reaction, determined for catalysts OH20 and HT over the range of temperature from 450 to 510 K, using a flow rate of 20 ml rain -1 for H 2 and with the CH3CN saturator at 273 K, was 59 kJ mol -] in both cases. The order of reaction with respect to ethanenitrile was measured for catalyst OH5 at 508 K and for HT at 483 K by maintaining the temperature of the saturator at different temperatures between 273 K and 295 K at 20 ml rain -] H 2 flow rate. From plots of log rate vs log [CH3CN ] the order derived was -0.25 + 0.08 for OH5 and 0.28 + 0.04 for HT. The order of reaction for H 2, obtained by diluting the flow of H 2 with He whilst keeping the saturator at 273 K, was measured for HT at 483 K and found to be 1.9 _4- 0.3.

4. DISCUSSION 4.1 Reducibility and metal surface area The characterization results show that the method of preparation of Fe/MgO exerts a significant influence on the amount of iron which can be reduced to Fe ~ (Table 1). The percentage reduction increases with the loading in each series (IP, CCP, OH), but the coprecipitation with NH4OH, which gives brucite-like structure to the precipitate and y-Fe.20 3 on calcination (the OH series) provides more easily reduced precursors than those obtained via precipitation of carbonate-rich hydroxycarbonate. This difference is even more strongly reflected in the results on iron surface areas (Table 2) where the OH series catalysts (and HT) show consistently higher values. As already indicated, the difficulty of reducing supported iron in hydrogen is well-known [6, 8,11 ]. It probably arises from a combination of causes, the two most important of which are a strong interaction with the support 16,8] and reoxidation or inhibition by water vapour in the pores of the oxide [14]. With MgO as support, there is undoubtedly a strong tendency for iron, especially at the F e 2 + stage of reduction, to be present at least in part as FeO-MgO (FexMgl_xO) solid solution [6,8]. This need not be deleterious to the ultimate formation of finely-divided iron, provided the method of preparation has led to a solid solution in which the Fe 2+ ions are well-distributed. The iron particles are limited in size

265 by the extraction and reduction of Fe 2+ ions being restricted to a relatively small depth of solid solution, as in the analogous case of reduction of high surface area NiO-MgO and CoO-MgO [15]. We suggest that this is working to greatest effect with the OH series of catalysts. The average size of the iron crystallites evaluated using the d = 0.85/D formula for OH5 is 3.8 nm as compared with 11.4 nm for IP5. In the IP series, the Fe 2 + ions in the FeO-MgO may be clustered, as noted in previous work [6]. The presence of some very finely-divided iron also in the catalysts with higher loadings is strongly indicated by the comparison between XRD-derived and dispersion-derived sizes in Table 3. The disparity, also evident in previous work [6,11], arises because very fine particles tend to be 'lost' in line-broadening analysis but fully contribute in chemisorption. The OH catalysts again show up favourably, exhibiting the lowest sizes by line-broadening analysis. Significantly, the above disparity is greatest with OH20, suggestive of a high proportion of very small Fe particles. At 30 mol % Fe, the support is less effective in dispersing the iron due to the high loading. m

4.2 Fe-catalyzed hydrogenation of nitriles The results show that iron supported on MgO is an effective catalyst for the selective hydrogenation of nitriles to the primary amines. However, the selectivity is impaired if the catalyst has low activity (Fig. 1). The inference from the experiments with propanenitrile is that the imine intermediate, which one may expect to be strongly chemisorbed and activated on iron, needs to react rapidly with adsorbed hydrogen if condensation reactions to form secondary and tertiary amines are to be avoided. The availability of adsorbed hydrogen may well be enhanced by high dispersion and a plentiful supply of finely-divided iron, conditions which are manifestly weakest in catalysts IP5, IP10 and IP20, and especially so in IP5 (Table 2), and it is these catalysts which are consistently lowest in activity. Catalysts of the CCP series, on the other hand, perform well, but in general the most effective are those of the OH series. These are the catalysts with the highest iron surface areas. Verhaak et al. [2] have shown convincingly that the reactions leading to the secondary and tertiary amines involve the support. The by-product 1-propylaminopropene which we observed is presumably formed via attack of propylamine on propylimine, followed by elimination of NH 3, as shown below: +H 2 CH3-CH2-C-~N ~

CH3-CH2-CH=NH / + NH2-CH2-CH2-CH3 NH 2 ,b I

CH3-CH2-CH2-N-CH-CH2-CH 3 H / - NH3 !

CHB-CH2-CH2-NH-CH =CH-CH 3

266 1-Propylaminopropene is thus on a logical route to dipropylamine, which can form directly by hydrogenation of its C = C bond. This condensation-elimination reaction may indeed involve the support [2] or at least the metal-support interface. The kinetic results with ethanenitrile support the above conclusions. The plot of Fig.2 shows that the hydrogenation to the primary amine is metal-catalyzed with a rate which is proportional to the metal surface area. The fact that the plot does not extrapolate to the origin can be interpreted as indicating that some of the iron measured by exposure to CO or N20 is not accessible to the reactant nitrile. The activation energy (59 kJ mol-1), which is the same for OH20 and HT, is similar to the value of 52-55 kJ mo1-1 reported by Verhaak for the hydrogenation of CH3CN on Ni and Co [16]. The near-zero order in CH3CN confirms strong chemisorption of the nitrile and the imine on the iron surface. The authors thank the Science and Engineering Research Council and also ICI Chemicals and Polymers Ltd for their support of this work.

REFERENCES 1. J. Volf and J. Pa~ek, Stud.Surf.Sci.Catal., 27 (1983) 105. 2. M.J.F.M. Verhaak, A.J. van Dillen and J.W. Geus, Catal.Lett., 26 (1994) 37. 3. F. Medina, P. Salagre, J.E. Sueiras and J.L.G. Fierro, Appl.Catal., A, 92 (1993) 131; 99 (1993) 115. 4. Ullmann's Encyclopaedia of Industrial Chemistry, Vol.A2, VCH, Weinheim, 1985. 5. J. yon Braun, G. Blessing and F. Zobel, Chem.Ber., 36 (1923) 1988. 6. M. Boudart, A. Delbouille, J.A. Dumesic, S. Khammouma and H. Topsae, J.Catal., 37 (1975) 486. 7. A.J.H.M. Kock, H.M. Fortuin and J.W. Geus, J. Catal., 96 (1985) 261. 8. D.E. Stobbe, F.R. van Buren, A.W. Stobbe-Kreemers, A.J. van Dillen and J.W. Geus, J. Chem. Soe., Faraday Trans., 87 (1991) 1631. 9. H. Topsae, J.A. Dumesic, E.G. Derouane, B.S. Clausen, S. Morup, J. Villadsen and N. Topsae, Stud.Surf.Sci.Catal., 3 (1979) 365. 10. H.J. Jung, M.A. Vannice, L.N. Mulay, R.M. Stanfield and W.N. Delgass, J.Catal., 76 (1982) 208. 11. S. Mousty, B.S. Clausen, E.G. Derouane and H. Topsae, Stud.Surf.Sci.Catal., 16 (1983) 385. 12. J.J.F. Scholten and J.A. Konvalinka, Trans.Faraday Soc., 65 (1969) 2465. 13. G.L. Vogler, X.Z. Jiang, J.A. Dumesic and R.J. Marion, J.Catal., 89 (1984) 116. 14. A. Baranski, M. ~agan, A. Pattek, A. Reitzer, L.J. Christiansen and H. Topsae, Stud.Surf.Sci.Catal., 3 (1979) 353. 15. J.G. Highfield, A. Bossi and F.S. Stone, Stud.Surf.Sci.Catal., 16 (1983) 181. 16. M.J.F.M. Verha~, Thesis, University of Utrecht, 1992. ISBN 90-393-0243-X.

J.W. Hightower,W.N. Delgass,E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996ElsevierScience B.V. All rights reserved.

267

Selective synthesis of ethylenediamine from ethanolamine over modified H-mordenite catalyst K. Segawa a, S. Mlzuno, " a M. S u g .m r aa, and S. Nakata b aDepartment of Chemistry, Faculty of Science and Technology, Sophia University, 7-1 Kioi-cho Chiyoda-ku, Tokyo 102, Japan* bR & D Center, Chiyoda Corp., 13 Moriya-cho, Kanagawa-ku, Yokohama 221, Japan The synthesis of ethylenediamine (EDA) from ethanolamine (EA) with ammonia over acidic types of zeolite catalyst was investigated. Among the zeolites tested in this study, the protonic form of mordenite catalyst that was treated with EDTA (H-EDTA-MOR) showed the highest activity and selectivity for the formation of EA: at 603 K, W/F=200 g h tool", and NH3/EA=50. The reaction proved to be highly selective for EA over H-EDTA-MOR, with small amounts of ethyleneimine (EI) and piperazine (PA) derivatives as the side products. IR spectroscopic data provide evidence that the protonated EI is the chemical intermediate for the reaction. The reaction for the formation of EDA from EA and ammonia required stronger acidic sites in the mordenite channels for higher yield and selectivity. I. I N T R O D U C T I O N Ethylenediamine (EDA) is used in a wide variety of applications in industry. The applications for EDA include uses as chelating reagents, surfactants, fabric softeners, lubricating oil additives, fungicides, insecticides, and resinous polymers. EDA is made by amminolysis of ethylene dichloride with ammonia [ E D C process], or by reductive amination of ethanolamine (EA) with hydrogen and ammonia fMEA process] on a commercial basis. However, the E D C process h a s numerous drawbacks: the production of by-product sodium chloride and the cost of corrosion-resistant equipment for such a process [1]. On the other hand, the M E A p r o c e s s includes high pressure reactions (10-20 MPa) over transition metal catalysts, and shows lower selectivity for EDA [2]. In view of the wide utility of EDA, there has been substantial work done in the preparation of EDA t h a t involves the reaction of ethylene oxide with ammonia, EA with ammonia, or ethylene glycol with ammonia. Certain zeolite catalysts have been studied in the field ofEA amination to suppress the formation of bulkier by-products, and generate enhanced yields of EDC [3]. The development of a selective catalyst is to reduce utility cost and/or to expand the plant capacity. The process requires a solid acid catalyst; objectives of the process have been to maximize conversion of the

268 organic substrate with ammonia to form EDA and to maximize selectivity. In this study, as an alternative process for EDA synthesis, acid-catalyzed amination of EA under atmospheric pressure, has been studied; such a process is much more environmentally benign than the present industrial processes. A s i ~ i f i cant advantage afforded by the treatment of EDTA in sodium form of mordenite followed by ion-exchange to protonic form is that the process is extremely effective in producing EDA in high conversion while at the same time being highly selective to EDA. This is contrast to many of the subsequent processes [EDC process and MEA process] which sacrificed selectivity for EDA in favor of conversion.

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

Preparation of catalysts. Zeolites (JRC-Z) and silica-alumina (JRC-SAL-2, Si/Al=5.3) samples were supplied by the Catalysis Society of J a p a n (JRC: J a p a n Reference Catalysts). Three different types of zeolites were studied: Na-FAU (faujasite, JRC-Z-Y5.6, Si/Al=2.8), Na-MOR (mordenite: JRC-HM10, Si/Al=5.0), and Na-lVIFI (ZSM5, JRC-Z5-25, Si/Al=12.5). K-LTL (Linde type L, HSZ500KOA, Si/AI=3.0) and K-CHA (K-chabazite, Si/AI=2), were supplied by TOSOH and by Air Products and Chemicals, Inc., respectively. Some Na-MOR samples were treated with an aqueous solution of H4EDTA (ethylenediamine-tetraacetic acid) under reflux at 383 K for 4 h. After cooling and filtration, the samples were washed with distilled water and dried at 373 K, and were then calcined at 773 K (EDTA-MOR, Si/A1=5.2-11.2). Add-type zeolites were prepared by ion-exchange of Na-form or K-form of zeolite with aqueous solution of NH4NO3; ion-exchanged samples were dried at 373 K for 24 h and then calcined in a furnace at a constant temperature increase (1 K rain i ) from 373 K to 773 K and kept at 773 K for 5 h. Catalytic reactions. The reaction was carried out at 543-643 K by using a flow reaction system with a mixture ofEA, NH3, and N2 in the ratio of 1/50/25 at atmospheric pressure. The flow rate of the mixture gas was 76 cm 3 rain "1. Prior to the reaction, the catalyst was calcined at 773 K under 02 flow for 2 b_ The reaction products were analyzed by an on-line gas chromatograph (FID) which was equipped with a 30-m capillary column (TC1701). Adsorption measurements. Chemisorption of base molecules (NH 3 and EA) on acidic zeolites was confirmed by IR spectroscopy and high-temperature microcalorimetry. A vacuum-tight IR cell with KBr windows was designed to fit an infrared spectrometer (270-30, Hitachi) and to be attached to a vacuum system (10 -4 Pa). The cell was arranged such that the zeolite wafer could be lowered into slots between the optical windows, and withdrawn upward by the action of a magnet into the heated portion for the pretreatment and adsorption of NHs and EA. After evacuation at 773 K for lh, the zeolite sample was cooled to 373 K before adsorption of the base molecules to be studied. IR spectra were obtained at room temperature. High-temperature micro-calorimetry of NH3 on zeolite catalyst was obtained at 473 K by the calorimeter (HAC-450G, Tokyo Rikou). Each sample (1.5 g) was charged to the calorimeter, and it was evacuated at 673 K for 4 h. NH3 (15 mmol g-i) portions were admitted dose by dose at 473 K.

269 3. R E S U L T S AND D I S C U S S I O N The catalytic activity and selectivity of EDA synthesis from EA and N H 3 over various zeolite catalysts are shown in Table 1. Among the various solid adds, the protonic form of mordenite (H-MOR) catalyst and the protonic form of mordenite catalyst that was treated with EDTA (H-EDTA-MOR), showed the higher selectivity for EDA. Over H-EDTA-MOR; the selectivity of EDA was about 72 % at 96 % of EA conversion with the presence of a n excess a m o u n t of ammonia (NHz/EA=50). Small amounts of EI and piperadine derivatives (PA) were formed as by-products. PA derivatives are mainly piperazine, amirtoethylpiperadine, and 1 , 4 - d i a z a b i c y c l o - o ~ e (DABCO). When the reaction t a k e s place over some other solid acid catalyst, such as amorphous silica-alumina, 100 % of EA converted to EI, PA derivatives, and other oligomers including aminoethylaminoethanol (Others: polyamines). On H-CHA (chabazite), H-FAU (faujasite-Y), and H-LTL (L-type), only a small a m o u n t of EDA was formed. The major product was EI and PA, or other polyamine oligomers were formed. On HMFI (ZSM5) catalyst, the major products were PA and other higher polyamines.

Table 1 Catalytic Activities and Selectivities of EDA Synthesis* over Various Zeolite Catalysts Catalyst*** SiO2-AI203 H-CHA H-FAU H-LTL H-MOR H-EDTA-MOR H-MFI

Pore Size

Conversion

Selectivity**/%

Si/A1

/nm

/%

EDA

EI

PA

Others

5.3 2.2 2.8 3.0 5.0 6.9 12.5

--0.38 0.74 0.71 0.70 0.70 0.54

I00 4 6 15 42 96 23

1 4 13 23 77 72 36

7 69 76 35 7 0 2

49 13 6 21 9 18 31

43 14 5 21 7 10 31

*Reaction Conditions: Temperature=603 K, W/F=200 g h mol"1, NH3fEA=50 ** EDA: ethylenediamine, EI: ethyleneimine, PA:piperazine derivatives *** CHA: chabazite, FAU: faujasite (Y), LTL: Linde type L, MOR: mordenite, EDTA-MOR: Namordenite was treated with EDTA and then ion-exchanged to H-form, MFI: ZSM5.

The results (Table 1) suggest that the selective synthesis of EDA from EA and NH 3 requires stronger acidic sites in the mordenite channels to suppress the formation of bulkier PA derivatives and other polyamines (Others in Table 1). The t r e a t m e n t by EDTA on Na-MOR and then ion-exchange to protonic form gives significant improvements for EDA synthesis. Figure 1 shows the catalytic activities and selectivities for EDA synthesis over H-MOR (Figure 1A) and HEDTA-MOR (Figure 1B), as a function of time on stream. Over H-MOR catalyst, the catalytic activity was much lower than t h a t of H-EDTA-MOR catalyst, and some deactivation occurred with increasing time on stream. On the other hand,

270 even if the catalytic activity over H-EDTA-MOR was higher than t h a t of other catalysts, no significant deactivation occurred aIter 3 hours on stream under the whole range of the reaction conditions tested in this study. The selectivity to EDA over H-EDTA-MOR is much higher than t h a t over H-MOR.

"

100

I

I

I

A 80

O

r'-!

60

-

" 40 .9

-

(D

I

I

- 100

EDA

-

._>

I

r-I

~

~

L~

-

I

.

B i.

i

I

I

i

_ Conversion

I

__~_

_

80

I-!

m

W

!--1

LJ

EDA 60

(,D e~

" 0

20

"0.

PA

_i

Others M

El 0

1

2

40

Conversion

-

20 ~

I

3

m

I

4

m

P~~iI~II~

m

I

I

5

6

Others . =

0 7

. 0

1

9 O

_ '~J

_

;,.,,--

~

_A

A

_A

_A

2

3

4

5

6

Time on s t r e a m / h

Time on s t r e a m / h

Figure 1. Catalytic activities and selectivities for EDA synthesis over (A) H-MOR (Si/AI=5.0) and (B) H-EDTA-MOR (Si/AI=6.1). Reaction conditions: temperature=583 K, NH3~A=60, PEA-1.2 kPa.

100

80

Or)

-

I

D

.

.

I

!

- 0 ' "

Conversion

QJb

•

- _

60

r

._o

40

-

o O

==.Others ~ --

e-.

20

-

lib /'%

-

.,,.,,

~.

PA

%./

v

B A,

0 4.0

6.0

8.0

, 10.0

~_ 12.0

Si/AI

Figure 2. The effect of dealumination by EDTA over MOR (Si/AI=5.0) for ethylenediamine synthesis. Reaction conditions: temperature=603 K, NH3/EA=50, PEA=I.4 kPa.

271 Figure 2 shows the catalytic activity and selectivities after treatments by EDTA, as a function of Si/A1 ratios. The treatment of zeolite with H4EDTA has been shown to dealuminate the zeolite framework. The bulk ratio of (Si/A1) in the zeolite was determined by XRF analysis. The reaction of EA with excess amounts of NH 3 over H-MOR provided higher selectivities for EDA (61%) but, activities are low (53 % conversion). In an attempt to improve activity and selectivity by modification of catalyst structure, moderately dealuminated H-MOR catalysts were prepared by EDTA treatments. When the Si/A1 ratios become above 6 from Si/AI=5.0 (H-MOR) of parent mordenite, a large enhancement of the catalytic activities was observed (87-93 % conversion). The selectivity of EDA was not significantly changed over H-EDTA-MOR in comparison with H-MOR, except over higher dealuminated catalysts (Si/AI=ll.2). Due to the channel structure of MOR, EDA is formed over the protonic sites that are located inside the main channels of MOR, and cyclization to PA and polyamines (higher oligomers) might be formed non-selectively on the external surfaces of MOR crystals.

160 L 140

I

I

I .........

.................................. t

t

f

(SI/AI)

120

O H-MOR (5.0)

r-

.g

g

0r "0 0 -r

---! ..........

~00 80

.......... ..........

i! ...... ......

..........

! .......

oo .......... i ........... 40

....

0.0

0.5

i ..........

O H-EDTA-MOR

(5.9)

9 H-EDTA-MOR

(6.1)

A H-EDTA-MOR

(8.0)

i .......... I t'

i

i

1.0

1.5

2.0

2.5

NH3 adsorbed/mmol g-1

Figure 3. High-temperature micro-calorimetry ofNH 3 on H-MOR (Si/AI=5.0) and H-EDTA- MOR (Si/A1=5.8-8.0): NH 3 adsorbed at 473 K. H-MOR has larger numbers of stronger acidic sites than those of other zeolites [5]. Figure 3 shows high-temperature micro-calorimetry of NH 3 over HMOR (Si/AI=5.0) and H-EDTA-MOR (Si/A1=5.8-8.0). We preferred the temperature at 473 K for calorimetry measurements to collect precisely the information of the stronger acid sites. All micro-calorimetric curves decreased with increasing coverage of NH 3 on zeolites. The results suggest that the acid strength and acid amount of H-MOR are higher than the H-EDTA-MOR. The MOR samples after being dealuminated by EDTA (H-EDTA-MOR) show similar acid site distributions for the range of Si/A1 ratios from 5.8 to 8.0.

272 100. -

.

I

'

I

'

A 80

w

9...@--"

-zl00 !

,...e""'"

-

80

,

40

'

. .

,,'

40

~-

o 20 rO

0

-

U _

0

100

200

,

_A 300

O

,

A

400

Contact time (W/F) /g h tool-1

PAl '

-

hers

20

~

Oon,,or ,on

~

~ 0 500 0

~

P 20

A 40

I

l 60

80

100

NH3/EA

Figure 4. EDA synthesis on H-EDTA-MOR (Si/Al=6.1) as a function of contact time (A) and as a function of partial pressure of NH 3. Reaction conditions: temperature=583 K, NH3/EA=0-80, PEA=I.4 kPa.

Deeba and coworkers [3] reported that the H-MOR dealuminated by acid leaching showed higher selectivity for EDA at lower conversions: about 60 % selectivity at 30 % conversion (NH3~A=16). However, at higher conversions, selectivity for EDA was not as high. In this study, if the reaction conditions included longer contact time (W/F=500 g h tool-Z), conversion exceeds about 95 % of EA with 80 % selectivity to EDA. The time courses of EDA synthesis over HEDTA-MOR catalyst at 583 K (PEA=I.4 kPa, NH3/EA=50) are shown in Figure 4A. The initial product of reaction was EI, and the formation of EDA followed. However, the selectivity of EDA did not exceed 85 % at higher conversion region. The selectivity of PA derivatives and higher polyamines increased with increasing contact time. The catalytic performance over H-EDTA-MOR as a function of partial pressure of N H 3 is also shown in Figure 4B. The activity rises with increasing partial pressure of NH3. Importantly, the formation of higher polyamines and EI is extremely retarded at higher partial pressure of NH3, whereas E D A is the major product. The results suggest that the intramolecular condensation of E A occurred at the initialstage of reaction to produce El intermediate, then adsorbed El was activated by the stronger protonic acid sites to produce EDA. The activation of El over the stronger acidic sitesis the rate-determining step for this reaction; such adsorption species m a y be converted to E D A with excess amounts of N H3. Intermolecular condensation of E A to form El, P A derivatives,and higher polyamines m a y occur at weaker acidic sites,which m a y be located on the external surfaces of mordenite crystals. W h e n E A adsorption occurred under N 2 flow at 583 K on H - E D T A - M O R catalyst without N H 3, El, higher polyamines, and P A were formed (Stage A in Figure 5). But, when an excess amount of N H 3 (NH3/EA=50) was introduced in

273 the reaction system at Stage B, EDA formed selectivity increased as did the conversion. At Stage C (same reaction conditions at Stage A), the formation of EDA was suppressed completely, and EI, higher polyamines, and PA reappeared again.

100

>,

._>

-

80

9

I

A

'

'~

El

"~ 60 (D =o

-

t,.

O

0

-0"

20

-

0

EDA

El

-"

-

"@ . . . .

0,,

Conversion -O ....

e"

'

Conv ers ion

e...

._o 40 O3

i

c

-

O

'

P A' -

e.-

~

Conversion ~

Others

2

4

" --O. . . .

6

@--

8

Time on stream/11

Figure 5. Transient responses for EDA synthesis over ; H-EDTA-MOR 1 (Si/AI=6.1). Reaction conditions: temperature=583 K, W/F=200 g h mol'-, PEA=0.95 kPa, NH3/EA=0 at Stage A and C, NH3/EA=50 at Stage B.

The adsorption studies of EA or ammonia on H-EDTA-MOR (Si/Al=6.1) by IR spectroscopy (Figure 6) suggested that the reaction may proceed through ammonio-ion of EA over protonic acid sites to produce an EI intermediate. When HEDTA-MOR was exposed to 0.3 kPa of EA and evacuated at 473 K (Figure 6A), NH3 T deformation bands build up at 1597 cm* and 1497 cm " together with CH2 deformation band at 1460 cm". Those deformation bands are attributed to the -1 presence of protonated primary amines over the catalysts. At 1372 cm and 1324 cm "1 wave number regions (Figure 6A), OH deformation bands are observed. We previously reported that the major acidic sites were BrCnsted sites on HMOR, which are mainly located in the main channels [5]. The IR spectrum (Figure 6A) suggests that EA is protonated and adsorbed as ammonio-ion of EA (NHs*CH2CH2OH)+ on H-EDTA-MOR, and not adsorbed as an oxonium-ion (NH2CH2CH2OH2-). When the ammonio-ion on H-EDTA-MOR is heated and evacuated at higher temperature (Figure 6B-6D), the adsorption species were changed to secondary amines. The deformation bands of ammonio-ion are shifted to lower wave number: NH2 § deformation bands build up at 1600 cm 1, together with CH2 N§ deformation band at 1442 cm "1. The intensities of OH deformation bands (1370 -1 cm , 1324 cm "1) are decreased with increasing evacuation temperature. The

274 results suggest t h a t the ammonio-ion transformed to protonated EI over the catalyst surface. The intr~molecular condensation of EA occurred at the initial stage of reaction to produce EI. Then EI was activated by the stronger protonic acid sites to produce EDA. The secondary protonated amines are strongly held on the surfaces of H-EDTA-MOR without the presence of ammonia.

1600

1442

o

1600

L

/

.

A, I .... 1750

I 1500

I I 1300 1750

! 1500

I 1300

Wavenumbers/cm 1

Figure 6. IR spectra of adsorbed EA and NH3 on H-EDTA-MOR (Si/AI=6.1): (A) H-EDTA-MOR exposed 0.3 kPa of EA at 473 K and evacuated at 473 K, (B) evacuated at 523 K, (C) evacuated at 573 K, (D) evacuated at 623 K, (E) H-EDTA-MOR exposed 0.3 kPa of Nil3 at 473 K and evacuated at 473 K, (F) after recording of spectrum E, the sample was exposed to 0.3 kPa of EA at 473 K and evacuated at 473 K, (G) evacuated at 523 K, (H) evacuated at 573 K.

When H-EDTA-MOR was exposed to 0.3 kPa of NH3 at 473 K and evacuated at 473 K (Figure 6E), only protonated ammonia (NH4 +) was observed. The NH4 + deformation band builds up at 1443 cm -1 together with a small amount of coordinated NH3 bonds (deformation) t ha t attached to Lewis acid sites at 1624 cm -1.

275 The major acidic sites on H-MOR are Brcnsted sites determined by pyridine adsorption studies: above 80 % of acidic sites are Brcnsted sites and the rest are Lewis acid sites [4, 5]. A f a r adsorption of NH3, 0.3 kPa of EA are admitted on HEDTA-MOR at 473 K (Figure 6F); adsorbed NH3 is easily replaced by EA to produce defo _rma_fion bands of NH3 § (1597 cm -x, 1497 cm-X), CH2 (1460 cm-1). This spectrum is quite the same as the spectrum in Figure 6A. The results suggest that adsorption of EA is much stronger than that of NH3. When adsorbed EA is heated up to 573 K (Figure 6G-6H), the spectra are almost the same as the spectra in Figure 6B and 6C.

HO~NH

2 '---

L

EA

H20

HN \

~NH2

/

NH

\

PA

ll-H2~ N H

El

H+

~ N+ ~ H2

/

]

l EA,EOA -~ NH3

H2N ~

NH2

EDA

Scheme 1. Reaction pathway of EA synthesis from EDA and NH3 over H-EDTAMOR.

4. CONCLUSION The synthesis of EDA from EA with NH3 over acidic types of zeolite catalyst was investigated. Among the zeolites tested in this study, H-EDTA-MOR was the best catalyst. The reaction proved to be highly selective for EA over H-EDTAMOR, with small amounts of EI and PA derivatives as the side products. The catalytic activity for EDA synthesis rises with increasing partial pressure of NH3. The initial product of reaction was EI, and the formation of EDA followed. The reaction pathways for the formation of EDA from EA and NH3 are summarized in Scheme 1. The results suggest that the formation of EDA required stronger acidic sites in the mordenite channels with excess amounts of NH3. The mordenite channels may retard the formation of bulkier PA derivatives and other polyamines. The reactions of EA proceed through ammonio-ions by the addition of protons of H-EDTA-MOR; then intramolecular condensation of ammonio-ions of EA occurred to produce an EI intermediate over the active sites of the catalyst. EI is protonated by the stronger acidic sites with excess amounts of N H 3 to

276 produce EDA. Protonated EI intermediates are quite stable over H-MOR because the presence of stronger acidic sites. The excess amounts of NH s m a y be required to remove the protonated EI to form EDA. This must be the main reason why the dealuminated MOR (H-EDTA-MOR) enhances the catalytic activity for EDA synthesis from EA and NH 3, since the acid strength of H-EDTAMOR is slightly lower than that of H-MOR (see Figure 3), to give relatively easier desorption of protonated EI intermediate from the surfaces to form EDA. REFERENCES

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

S. Kumoi, N. Kubota, and T. Hiroi, Kagaku Keizai, 11 (1985) 56. Ninters, J. R., U.S. Pat., 4 404 405 (1983). Deeba, M., Ford, M. F., and Johnson T. A., European Patent 252 424 (January 13, 1988). M. Deeba, M. E. Ford, T. A. Johnson, and J. E. Premecz, J. Mol. Catal., 60 (1990) 11. K. Segawa, M. Sakaguchi, and Y. Kurusu, Stud. Surlf. Sci. Catal., 36 (1988) 579. K. Segawa and H. Tachibana, J. Catal., 131 (1991).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

277

Competitive Reaction Pathways in Propane Ammoxidation over V-Sb-Oxide Catalysts: an IR and Flow Reactor Study G. Centia, b and F. Marchi ~ a Dip. Chim.Industriale e dei Materiali, University of Bologna, V.le Risorgimento 4, 40136 Bologna, Italy* b Dip. Chimica Industriale, University of Messina, Contrada Papardo, 98100 Messina, Italy The surface t r a n s f o r m a t i o n s of propylene, allyl alcohol and acrylic acid in the presence or absence of NH3 over V-antimonate catalysts were studied by IR spectroscopy. The results show the existence of various possible p a t h w a y s of surface t r a n s f o r m a t i o n in the m e c h a n i s m of propane ammoxidation, depending on the reaction condition and the surface coverage with chemisorbed NH3. A surface reaction n e t w o r k is proposed and used to explain the catalytic behavior observed in flow reactor conditions. 1. I N T R O D U C T I O N V-Sb-oxide based catalysts show interesting catalytic properties in the direct synthesis of acrylonitrile from propane [1,2], a new alternative option to the commercial process s t a r t i n g from propylene. However, further i m p r o v e m e n t of the selectivity to acrylonitrile would strengthen interest in the process. Optimization of the behavior of Sb-V-oxide catalysts requires a thorough analysis of the relationship between structural/surface characteristics and catalytic properties. Various studies have been reported on the analysis of this relationship [3-8] and on the reaction kinetics [9,10], but little attention has been given to the study of the surface reactivity of V-Sb-oxide in the transformation of possible intermediates and on the identification of the surface mechanism of reaction. 2. E X P E R I M E N T A L V-Sb-oxide samples with Sb:V ratios of 1.0 and 3.0 were prepared by the following precipitation-deposition method: VC13 is dissolved in a 0.1N aqueous HC1 solution (A) and SbC15 is dissolved in a 3-5N HC1 aqueous solution (B). The dropwise addition of A to B leads to the formation of a white precipitate (mainly Sb-hydroxide) and a d a r k blue solution due to VO 2§ ions. VO 2§ is then precipitated over the Sb-hydroxide by adding dropwise a concentred aqueous solution of "Fax: +39-51-644.3680;e-mail: [email protected]

278 ammonia up to basic pH. The precipitate is filtered, washed three times with distilled water and then dried at 140~ overnight. The resulting solid is then calcined in a flow of air up to 600~ (3 h) using a constant rate of increase in temperature (50~ The surface areas of the samples after calcination are 17 and 10 m2-g-1 (Sb:V= 1.0 and 3.0, respectively). Prior to r u n n i n g the catalytic tests the samples prepared by this precipitation-deposition method were activated at 500~ (3h) in a flow of propane, a m m o n i a and air. Details on the characterization of these samples have been reported elsewhere [8,11]. The catalytic tests were carried out using an a p p a r a t u s composed of (i) a continuous fixed-bed stainless steel reactor, (ii) a section for the preparation of feed by mixing already calibrated mixtures of the single components in helium, (iii) a system for on-line gas-chromatographic (GC) analysis (two GCs equipped with a flame ionization and thermoconducibility detector, respectively) and (iv) a section for the analysis of NH3 conversion and HCN formation by absorption in appropriate solutions and titration[6,9,10]. Tests were made using 2-4 g of sample with particle dimensions in the 0.1-0.2 m m range and diluted in a 1:5 ratio using an inert support. The axial t e m p e r a t u r e profile was monitored by thermocouples inserted into the catalytic bed. Tests were made using the following feedstock: 7.5% propane and a O2/C3 and NH3/C3 ratio of 1.56 and 1.60, respectively. The gas-space hourly velocity (GHSV) was 2800 h -1. Fourier-transform infrared (IR) spectra (resolution 2 cm -1) were recorded with a Perkin Elmer 1750 i n s t r u m e n t in a quartz cell connected to grease-free evacuation and gas m a n i p u l a t i o n lines. The self-supporting disk technique was used. Before recording the spectra, the samples were treated with 02 at 450~ (lh), then cooled down to r.t. before evacuating the 02. The sample was then evacuated at 400~ Evacuation at higher t e m p e r a t u r e s lead to a drastic cut off of IR trasparency. All reactants were purified prior to the adsorption experiments. Due to the better resolution of the spectra, only results for Sb:V=I.0 are reported here, however the IR data for Sb:V=3.0 were not significantly different.

3. R E S U L T S 3.1 C a t a l y t i c T e s t s The catalytic behavior in propane ammoxidation of Sb:V=I.0 and 3.0 is summarized in Fig. 1. The tests were carried out using a propane concentration of about 8% and oxygen as the limiting reactant, because these experimental condit-ions agree with those indicated as preferable in the patent literature [12] and from the analysis of the reaction kinetics [9,10]. For both catalysts, as the reaction t e m p e r a t u r e increases the the selectivity to acrylonitrile (ACN) passes t h r o u g h a m a x i m u m at about 480~ Maximum selectivity is about 30% and 60% for Sb=l.0 and 3.0, respectively. Increasing the temperature, decreases the selectivity to propylene (C3=) and increases t h a t to carbon oxides (COx). It may be noted t h a t the lower selectivity to ACN in Sb:V= 1.0 is

279

not necessarily associated with a higher formation of carbon oxides, but r a t h e r with a higher side conversion of NH3 to Nz "~which leads to a higher consumption of O2. Sb:V=I.0 is thus less selective to ACN, but more selective to C3=. It is interesting also to note t h a t the formation of acetonitrile does not follow t h a t of ACN, but instead selectivity to AcCN decreases as the temperature increases. The AcCN/ACN ratio thus is m a x i m u m at low temperatures and decreases ~continuously. In Sb:V=I.0 the ~=AcCN/ACN ratio passes from ~0.60 at 440~ to 0.18 at 500~ whereas the ratio is about ~ over Sb:V=3.0 (0.23 and 0.08, respectively).

3.2 Infrared study

Fig. 1 Catalytic behavior in propane ammoxidation of Sb:V = 1.0 and 3.0 (bottom and top, re- Propylene: The IR spectra of praspectively). Symbols: C conversion, S selectivity, pylene in contact at r.t. with C3 propane, C3= propylene, ACN acrylonitrile, Sb:V=I.0 and evacuation at proAcCN acetonitrile, gressively higher t e m p e r a t u r e s are reported in Fig. 2. Well evident bands are noted at 1660, 1468 and 1455 (double), 1388 and 1372 (double), 1328, 1255, 1170, 1132 and 1101 cm -1, the last very intense. The same bands are observed by adsorption of isopropyl alcohol and may be assigned to a mixture of a isopropoxylate species (1101 cm-1; vc-o ) and coordinated acetone (intense band at 1660 cm-1; vc=o); the latter assignment is confirmed by the analysis of the spectrum of acetone adsorbed over the same catalyst. Analogous species were detected by adsorption of propylene over vanadiatitania catalysts [13,14]. Additional weak bands at 1640 and 1070 cm -1 are observed in the spectrum of propylene in contact with the catalyst (Fig. 2a), but disappear by evacuation at r.t. (Fig. 2b). They can be assigned to vc--c and vc-o, respectively, of propylene adsorbed in the form of allyl alcoholate. With increasing evacuation t e m p e r a t u r e (Fig. 2c,d) the bands associated with isopropoxylate progressively decrease and disappear by evacuation at 120~ Likewise, the bands of coordinated acetone increase in relative intensity up to about 160~ and then decrease at higher temperatures. The spectrum of the sample evacuated at 160~ (Fig. 2e) is dominated by an intense band near 1540 cm -1 plus other bands at 1438, 1660, 1420 and 1370 cm-1; the latter three bands

280

are due to acetone which has not been completely transformed. At an evacuation temperature of 200~ (Fig. 2f) the spectrum is dominated by two intense bands at 1532 and 1448 cm -1, typical of carboxylate species. Other weak bands at 1640 and 1375 cm -1 (vcc and 5CH, respectively) indicate the presence of an acrylate species and at 1360 cm -1 (5CH3) of an acetate. The very weak bands at 1860 and 1795 cm -1 suggest the possible formation of a cyclic anhydride (vas and vs of C=O in cyclic anhydrides, respectively), reasonably by dimerization of propylene and oxidative attack to form maleic anhydride.

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Propylene thus readily reacts at r.t with V-Sb-oxide to form an isopropoxylate species which then may develop along two possible pathways: (i) the first route (attack from surface hydroxyl groups) gives products of carbon chain degradation (acetone plus CH4) ! ~ '~ e a! * and (ii) the second route occurring at higher ii. temperatures gives rise first to an allyl alcoholate species which then transforms to acrolein and acrylate. The second routes b "' J" t ~ -i ~'~ thus involves nucleophilic oxidative attack by lattice oxygen. Propylene + NH3: The results of propylene and ammonia coadsorption experiments are summarized in Fig. 3. Two series of tests were made. In the first series of tests (procedure A) propylene is put in contact at r.t. with the catalyst for 5 min, then removed by evacuation at r.t.; later ammonia 1800 1600 1400 1200 o m ~ is put in contact for 2 min at r.t. and then Fig. 2 IR spectra of 60 torr propylene removed by evacuation. Only species chemiin contact at r.t with Sb:V=I (a), subsorbed on the catalyst thus remain. Spectra sequent evacuation at r.t. (b) and folwere then recorded at increasing temperalowing evacuations at increasing temperatures: (c) 80, (d) 120, (e) 160 and tures under vacuum. In the second series of experiments (procedure B) propylene and (f) 200~ ammonia together are put in contact with the catalyst and the spectra recorded at increasing temperatures of contact with the mixture of C3H6+NH3. In the first experiments the change in coadsorbed species is studied, whereas in the second type of tests the effect of the presence of propylene and ammonia also in the gas phase is analyzed. ,"

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282 face acidity characteristics of V-Sb-oxide catalysts, thus modifying surface reactivity.

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Allyl alcohol: Allyl alcohol in contact with VSb-oxide catalyst at r.t. (Fig. 4a) gives rise to a spectrum characteristic of allyl alcoholate (1645 cm-l: vc=c; 1422 cm-Z: scissoring =CH2; 1445, 1342, 1360 and 1104 cm -1 s c i s s o r i n g CH2-O-, wagging CH2 and vc-o). A progressive increase in the i n t e n s i t y of the band at 1185 cm -1 with increasing t e m p e r a t u r e of evacuation indicates t h a t acrolein, not present initially, forms progressively with a parallel decrease in the relative intensity of the bands of allyl alcoholate. At 200~ the bands of acrolein (1641 cm -1 vc=o, 1625 cm -1 vc=c, 1430 cm -1 scissoring CH2, 1372 and 1282 cm -1 5CH, 1188 cm -1 vcc) disappear with the formation of new b a n d s clearly evident in the spectrum after evacuation at 250~ (Fig. 4g). The l a t t e r s p e c t r u m is characterized by bands at 1636, 1530, 1437, 1375 and 1278 cm 1 which correspond to those observed for acrylate species. The intense band at 1145 cm -1 which develops at higher t e m p e r a t u r e s is instead due to the reduction of the catalyst and formation of oxygen vacancies [5]. The adsorption of allyl alcohol t h u s readily gives rise to an allyl alcoholate species at room t e m p e r a t u r e . At higher t e m p e r a t u r e s this species first t r a n s f o r m s to chemisorbed acrolein and t h e n to an acrylate species. Allyl alcohol + NH3: The coadsorption exp e r i m e n t s were m a d e by first p u t t i n g the catalyst into contact with NH3 at 200~ followed by removal of gaseous NH3 at r.t. and then p u t t i n g the c a t a l y s t with the chemisorbed a m m o n i a into contact w i t h the allyl alcohol at r.t. Using procedures for the coadsorption tests like those used for propylene + NH3, less resolved spectra were obtained. The more r e m a r k a b l e differences in comparison with the analogous spectra for propylene + NH3 (Fig. 3) or acrylic acid + NH3 coadsorption (see below) are t h a t (i) the band in-

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dicating the formation of a nitrile species (2280 cm-Z) falls at a slightly higher frequency w i t h respect to the former case, due to interaction with Lewis sites, and occurs at lower temp e r a t u r e s (150-200~ i n s t e a d of 250300~ and (ii) the formation of nitrile species also occurs in the absence of a m m o n i a in the gas phase. Bands also are observed at 1550 and 1670 cm -1 which agree w i t h the formation of an imine or oxime species by reaction of acrolein with coordinated NH3. This i n t e r m e d i a t e species t r a n s f o r m s to nitrile species at lower temperatures t h a n those for the i n t e r m e d i a t e species formed in co-adsorption tests of acrylic acid and a m m o n i a .

Acrylic acid: The IR spectra obtained

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by adsorption of acrylic acid at r.t. and evacuation at increasing temp e r a t u r e s are shown in Fig. 6. Adsorption and evacuation at r.t. gives rise to a spectrum with bands at 1636 cm -1 (vc=c), 1494 cm 1 (Vas -CO2-), 1436 cm -1 (scissoring CH2 overlap to vs -CO2-), 1376 cm 1 (SCH), 1277 cm 1 (vcc) and 1070 cm -1 (rocking CH2) indicating the rapid formation of an acrylate species. Weak bands are also observed at 1660 and 1600 cm -1 which m a y be a t t r i b u t e d to weakly bonded acid. These bands d i s a p p e a r by evacuation above r.t.. The acrylate species is stable up to an evacuation t e m p e r a t u r e of about 200~ and then decreases in relative intensity at higher t e m p e r a t u r e s . Already in the spectrum obtained by evacuation at 200~ (Fig. 6c) a b r o a d e n i n g of the band at 1494 cm -1 and a shoulder at 1350 cm -1 indicate partial degradation to an acetate species.

Acrylic acid + NH3: Two series of experiments were made. In the first series

284 (procedure A) ammonia is first adsorbed at r.t. followed by removal of gas phase ammonia by evacuation. Later the catalyst with chemisorbed ammonia is put in contact at r.t. with the acrylic acid. Spectra are then recorded at -...: : . . - ..~ increasing temperatures of evacuation. In the second series of experiments (procedure B) acrylic acid is first put into contact with the !." .~ catalyst followed by evacuation. Then the i. "..'/ "" catalyst with chemisorbed acrylic acid is put q-.. into contact with NH3 at increasing temperatures. Analogously to the case of propylene and C ..: a aI 1". ammonia coadsorption, using procedure A only ........._ :I,, , chemisorbed species are formed, whereas in I I I | ii I procedure B ammonia is also present in the gas phase. The two procedures give rise to different results. In both cases acrylic acid, present in the p,( ,I form of acrylate, readily reacts with ammonia / I!' 'I at r.t. forming a species characterized by an inq | i| ~'. I ,,, II I " /ills :l tense band at 1535 cm -z indicating the formaJr I |1 ~ i! 9t i tion of an amide. With increasing reaction b ,:: :; temperature (100~ however, in the case of procedure A the band at 1535 cm -1 shii~s to 1495 cm -1 and a weak band forms at 1720 cm -1. The latter band is characteristic of undissociated and weakly coordinated acrylic acid. This indicates that at 100~ amide dissociates with formation of the free acid. When ammonia is instead present in the gas phase (procedure B), the amide species undergoes transformation to 1800 1600 1400 1 2 0 0 r "1 acrylonitrile with a maximum in the intensity Fig. 6 IR spectra of 1 torr acrylic of the vcs band at 2220 cm -1 at an evacuation acid in contact (5 min) with Sb:V=I temperature of about 300~ and evacuation at r.t (a), and folCoordinated acrylic acid and ammonia thus lowing evacuations at 100 (b) and react faster at r.t. to form acrylamide, but in 200~ (c). the absence of ammonia which inhibits the reactivity of the BrOnsted sites, the amide dissociates at higher temperatures with formation of the free acrylic acid. When the reactivity of the BrOnsted sites is blocked by their transformation to ammonium ions, the amide m a y be dehydrogenated to form the acrylonitrile. ~

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Co-adsorption experiments show a complex role of the nature and concentration of chemisorbed ammonia species. Ammonia is not only one of the reactants for the synthesis of acrylonitrile, but also reaction with Br@nsted sites inhibits their reactivity. In particular, IR experiments show that two pathways of reaction are possible from chemisorbed propylene: (i) to acetone via isopropoxylate intermediate or (ii) to acrolein via allyl alcoholate intermediate. The first reaction occurs preferentially at lower temperatures and in the presence of hydroxyl groups. When their reactivity is blocked by the faster reaction with ammonia, the second pathway of reaction becomes preferential. The first pathway of reaction is responsible for a degradative pathway, because acetone further transform to an acetate species with carbon chain breakage. Ammonia as NH4 + reacts faster with acrylate species (formed by transformation of the acrolein intermediate) to give an acrylamide intermediate. At higher temperatures the amide may be transformed to acrylonitrile, but when BrCnsted sites are present, the amide may be hydrolyzed to reform ammonia and free, weakly bonded, acrylic acid. The latter easily decarboxylate forming carbon oxides. Ammonia also reacts with the acrolein intermediate, via the formation of an imine or possibly oxime intermediate which transforms faster to the acrylonitrile than to the acrylamide intermediate. This p a t h w a y of reaction occurs at lower temperatures in comparison to that involving an acrylate intermediate, but its relative importance depends on the competitive reaction of the acrolein intermediate with the ammonia species and with catalyst lattice oxygens. NH3 coordinated on Lewis sites also inhibits the activation of propane differently from that absorbed on Brr sites. The scheme shown in Fig. 7 summarizes the proposed surface reaction network in propane ammoxidation.

286 Although the reaction scheme of Fig. 7 is based mainly on the IR data, it also can be used to discuss various aspects of the surface reactivity found in flow reactor studies (Fig. 1). The higher acetonitrile to acrylonitrile ratio at lower temp e r a t u r e s agrees well with t h a t expected from the effect of the reaction temperature on the relative rates of propylene transformation via isopropoxylate or allyl alcoholate intermediates. The considerable effect of the rate of side a m m o n i a oxidation to N2 on the selectivity to acrylonitrile (compare Fig. 1A and 1B) agrees well with the effect of a m m o n i a on the modification of reaction p a t h w a y s discussed above, in addJt|on to its role as coreactant. The optimization of V-Sb-oxide catalysts thus requires the tailoring of various aspects of the surface reactivity: (i) increasing the rate of the p a t h w a y via allyl alcoholate versus t h a t via isopropoxylate, (ii) increasing the rate of reaction of ammonia with the acrolein intermediate versus conversion of the latter to acrylate and (iii) reducing the rate of conversion of chemisorbed a m m o n i a to N2, the latter a factor which also has a considerable effect on the first two points. The financial support from the Ministero Pubblica Istruzione (60%) is gratefully acknowledged.

REFERENCES 1. Y. Moro-oka, W. Ueda, in Catalysis - Vol. 11, The Royal Society of Chemistry, Cambridge U.K. 1994, p. 223. 2. G. Centi, R.K. Grasselli, F. Trifirb, Catal. Today, 13 (1992) 661. 3. R. Nilsson, T. Lindblad, A. Andersson, C. Song, S. Hansen, New Developments in Selective Oxidation II, V. Cortes Corberan and S. Vic Bellon Eds., Elsevier Pub.: Amsterdam 1994, p. 281. 4. A. Andersson, S.L.T. Andersson, G. Centi, R.K. Grasselli, M. Sanati, F. Trifirb, Appl. Catal. A, 113 (1994) 43. 5. G. Centi, S. Perathoner, Appl. Catal., 124 (1995) 317. 6. G. Centi, R.K. Grasselli, E. Patan~, F. Trifir6, New Developments in Selective Oxidation, G. Centi, F. Trifirb Eds., Elsevier Pub.: Amsterdam 1990, p. 515. 7. G. Centi, E. Foresti, F. Guarneri, New Developments in Selective Oxidation II, V. Cortes Corberan, S. Vic Bellon Eds., Elsevier Pub.: Amsterdam 1994, p. 281. 8. G. Centi, S. Perathoner, Preparation of Catalysts VI, G. Poncelet et al. Eds., Elsevier Science Pub.: Amsterdam 1995, p. 59. 9. R. Catani, G. Centi, F. Trifirb, R.K. Grasselli, Ind. Eng. Chem. Res., 31 (1992) 107. 10. A. Andersson, S.L.T. Andersson, G. Centi, R.K. Grasselli, M. Sanati, F. Trifirb, New Frontiers in Catalysis, L. Guczi et al. Eds., Elsevier Pub: Amsterdam 1993, p. 691. 11. G. Centi, P. Mazzoli, Catal. Today (issue on Fundamentals of Oxide Catalysts), in press. 12. M.A. Toft, J.F. Brazdil, L.C. Glaeser, U.S. Patent, 4,784,979 (1988). 13. V. Sanchez Escribano, G. Busca, V. Lorenzelli, J. Phys. Chem., 94 (1990) 8939. 14. G. Busca, G. Ramis, V. Lorenzelli, J. Molec. Catal., 55 (1989) 1.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

287

'Seeing' The Active Site in Catalysis. STM and Molecular Beam Studies of Surface Reactions Michael Bowker

Reading Catalysis Centre, Department of Chemistry, University of Reading, Whiteknights Park, Reading RG6 6AD. and IRC in Surface Science, University of Liverpool.

1. ABSTRACT The 'Holy Grail' of catalysis has been to identify what Taylor described as the 'active site' that is, that ensemble of atoms which is responsible for the surface reactions involved in catalytic turnover. With the advent of atomically resolving techniques such as scanning tunnelling microscopy it is now possible to identify reaction centres on planar surfaces. This gives a greater insight also into reaction kinetics and mechanisms in catalysis. In this paper two examples of such work are described, namely CO oxidation on a Rh(110) crystal and methanol selective oxidation to formaldehyde on Cu(110).

2. INTRODUCTION Since early in this century the concept of the 'active site' in catalysis [1] has been a focus of attention in this area of chemistry. This was proposed to be that ensemble of surface atoms/reactants which is responsible for the crucial surface reaction step involved in a catalytic conversion. Since those days much work has been done in the area, which cites the concept of the active site. However, no such ensemble has been positively identified due to the lack of availability of techniques which could image such a structure, which is of atomic dimensions. However, in more recent times science has made rapid strides in this direction. It is now possible to use EXAFS in situ during a catalytic reaction to examine the average coordination of metal atoms in the small particles which often exist in precious metal catalysts [2]. High resolution transmission electron microscopy has evolved to the level of atomic resolution, but can only be used ex-situ, or in situ with moderate pressures when special cells are fitted [3].

288 The most recent innovation in this field is scanning tunnelling microscopy, which has the capability of atomic resolution. In the work reported here two surface reactions are examined using this technique. These reactions are of relevance to automobile catalysis (CO oxidation on Rh) and methanol oxidation/synthesis on Cu. It is proposed that active sites are imaged in these reactions and that these active sites can indeexl be extremely dilute on the surface.

3. RESULTS AND DISCUSSION Oxygen adsorption on Rh results in reconstruction of the surface layer and an STM image of an oxygen-dosed surface is shown in fig. 1. The image shows bands of bright regions separated by dark lines. LEED has shown a multiplicity of structures on Rh(110), depending on the exact oxygen coverage. The important common feature of these is a missing row of Rh atoms oriented in the { 110} direction. Different oxygen coverages then produce different numbers of Rh atoms within the bright bands, which are mixed structures of Rh and O; the c(2x6) structure has 2 rows of Rh, while the c(2x8) has 3 [4-7]. Due to a staggering of the atoms in adjacent bands the periodicity in the unit cell is doubled, stretching over two bands in the [001] direction [6,7]. If these structures are exposed to CO, the reaction takes place in quite a surprising manner - the surface shows one dimensional reactivity and fig. 2 is an STM image of the same area before and after such a reaction. The reaction appears to initiate at step edges or defects on the surface and then proceeds in the {011 } direction to eat away at the Rh-O bands. Furthermore, certain structures are more reactive than others; the c(2x8) bands react faster than the c(2x6). It appears, then, that the active sites for CO oxidation are located at the end of these bands of oxidised Rh and are very specific in nature. This explains earlier results in which a low initial reactivity to CO2 formation was observed using molecular beam measurements [8]. The reaction probability was low at high oxygen coverages but increased as the oxygen-coverage diminished, until finally decreasing again as the oxygen becomes dilute on the surface. An example of this is shown in fig. 3a. If we turn to quite a different system, namely methanol oxidation on Cu(110), a similar reactivity pattern can be observed. Fig. 3b compares the reaction of methanol with preadsorbed oxygen at two oxygen coverages - 0.25 and 0.5 monolayers, the latter being saturation of the p(2x 1) structure on the surface. The oxygen saturated layer shows very low reactivity initially, which increases with time, goes through a maximum, and decreases to zero when the oxygen is all used up. When a Cu(110) surface is dosed with a partial layer of oxygen, the absorbate forms long, thin, one-dimensional islands. This structure has been observed by many other workers [9-11] and is quite characteristic of this adsorption system, with oxygen reacting with surface Cu atoms producing an added row structure of alternating oxygen and Cu atoms (a p(2xl) structure and with a surface atomic density of only half that of the original surface. This is thought to be formed by the diffusion of Cu atoms away from step edges to bind to growing CuO chains [ 10,11]. When this surface is reacted with methanol, then methoxy is formed as demonstrated by a variety of techniques, including molecular beam methods [12], TPD [12,13], IRAS [14] and UPS [15]. Fig. 4a shows a layer of oxygen which has partially reacted and formed

289

Figure 2. Showing the beginning of CO reaction with the surface along the long Rh-O islands, which are between the missing rows, as bright streaks.

290

30

E

v

20-

_J Z

~

d w

.

10-

r

<

0

1 O0

.

rIME(s.)

200

300

Figure 3a. CO 2 evolution from the reaction of a molecular beam of CO with oxygen predosed onto Rh(110) to a coverage of 0.7 monolayers at a crystal temperature of 540 K, showing low reactivity at high oxygen coverage.

0.2

OXYGEN

r. IA PRECOVERAGE 0.25ML

S

0.5ML

0.1

O0

'

0

-

-

9

!

1000

TIME(s.)

2O00

Figure 3b. The sticking probability of methanol on a Cu(110) surface predosed with half a monolayer of oxygen. There is an induction period to adsorption taking place and formaldehyde is evolved coincidentally with the sticking. Adsorption temperature of 353 K.

291 methoxy. The system is clearly phase separated, with no methoxy units within the remaining oxygen islands, but islands of methoxy exist between them. The methoxy forms an ordered p(5x2) glideline structure which can be seen in LEED [12]. More indicative of the active site for reaction is the series of images a-d given in fig.4. These show that, as reaction proceeds, the long oxygen islands generally do not diminish in width, but shrink in length. This proves that, again in this case, the active site for methoxy formation (from which formaldehyde is produced, by dehydrogenation) is at the end of the oxygen islands. In this case the ends are terminations of the Cu-O islands in the {001} direction. A model can then be drawn for this reaction, as shown in fig.5. The active sites are very few, in comparison with the total oxygen sites available, and a pseudo steady state is quickly produced with OH groups at the end of these chains. In the molecular beam experiments with 1/4 monolayer of oxygen adsorbed, for which the initial reactivity of methanol is quite high ( - 0 . 1 5 ) , even though the active oxygens are so dilute there is an immediate maximum high, steady state evolution of water, until the oxygen is used up at which point H20 evolution ceases. The STM results explain why there is no induction time to H20 evolution [12], since conversion of the few active sites to OH can occur very quickly, with the first few methanol molecules which arrive at the surface. The sequence of the reaction is then as follows. Methanol adsorbs onto the surface in a weakly-held state in which it can diffuse over relatively long distances to find active sites for reaction. The evidence for this is the fast H20 production and the sticking coefficient of - 0.15 indicating that, although the active sites are only few in number (an estimate, based on several STM images is - 5% of all the oxygen atoms for 1/4 monolayer dosed) the methanol has a > 15 % chance of finding them. This precursor state mobility and range of diffusion has been clearly demonstrated for alcohol reactions previously by us [16,17]. Immediately at the start of reaction these few active sites are converted to OH which, represent the steady state termination of these islands. More methanol then reacts at these sites to produce water. The next step in the reaction is not so certain, but clearly a new active site has to be created with an active oxygen. At steady state this would be achieved by oxygen from the gas phase, but in these transient conditions that cannot happen. Instead, it is likely that the terminal Cu atom, exposed by water formation, diffuses away from the site, leaving an exposed oxygen and reforming the active ensemble. Good evidence for this is that upon methoxy adsorption and reaction nearby steps change their shape and expand [18], presumably due to the addition of individual copper atoms. What happens to the methoxy formed by this process is strongly temperature dependent. At low temperature (up to - 340K) it is stable on the surface and forms the beautiful structures shown in fig.2. Since the active oxygen is used in such reactions then the methoxy must (i) not block the active site at its formation or (ii) diffuses away from the active site. Our evidence indicates the latter to be the case since methoxy is present at sites away from the oxygen islands. Above approximately 340 K the methoxy is unstable and decomposes to yield formaldehyde and hydrogen in the gas phase. Above approximately 400 K, the stoichiometry of the reaction changes to

CH3OH, + O~ --,H2CO, + H2Og

292 that i s , the methoxy is now so unstable that as soon as it ( and a hydroxyl group) is formed the methyl group loses a hydrogen, which reacts with the hydroxyl to produce water. Whether this is a direct reaction (that is, both hydrogens from the same methoxy form water) or an indirect reaction (that is hydrogen diffusion on the surface before reaction with hydroxyl) is uncertain at present.

Figure 4. Sequential images of the clean off of oxygen from the same area of Cu(110), by methanol; 20 nm x 20 nm images. These clearly show the 'shrinking' of the long, thin oxygen p(2xl) islands in one direction, by reaction of the oxygen at the short ends of the islands. This is where the 'active' oxygen is located.

293

Figure 5. Cartoon models of the reaction of methanol with oxygen on Cu(110). 1; A methanol molecule arrives from the gas phase onto the surface with islands of p(2xl) CuO (the open circles represent oxygen, cross-hatched are Cu). 2,3; Methanol diffuses on the surface in a weakly bound molecular state and reacts with a terminal oxygen atom, which deprotonates the molecule in 4 to form a terminal hydroxy group and a methoxy group. Another molecule can react with this to produce water, which desorbs (5-7). Panel 8 shows decomposition of the methoxy to produce a hydrogen atom (small filled circle) and formaldehyde (large filled circle), which desorbs in panel 9. The active site lost in panel 6 is proposed to be regenerated by the diffusion of the terminal Cu atom away from the island in panel 7.

294 These are just two examples of specific reactions, but other groups have recently demonstrated similar types of anisotropic reactivity on metal surfaces. For instance, these include the work of Besenbacher et al. for ammonia oxidation on Ni(110) [19], Crew and Madix for CO oxidation on Cu(110) [20], both of these using STM, and by Roberts et al. for ammonia radiation on Cu(110) using XPS [21]. The argument can be put forward that none of these findings relate to the active site involved in the steady state reactions occurring at high pressure and temperature during industrial catalysis. It is important to note in this respect that although these are transient and different conditions, surface science has identified the major intermediates involved in many industrial catalytic process (for example, formate in methanol synthesis [22] and nitrogen atoms in ammonia synthesis [23]). The sites responsible for various steps in the catalytic reaction can also be identified by STM and their importance for catalysis can be inferred from a knowledge of the important intermediates. Further, as discussed below, in the future STM might be applicable more directly to the high pressure and temperature conditions of industrial catalysis.

4. CONCLUSIONS In this paper the utility of STM for studying surface reactions has been demonstrated. It is proposed that the technique makes it possible to identify the 'active site' involved in surface reactions and that they can indeed be a minority species on the surface. We can go further and propose that it is possible, in the future, that STM could be used as an in-situ technique for studying catalytic reactions at the atomic level. However, several caveats must be placed against this possibility. Firstly, and most importantly, catalysis is a dynamic phenomenon. That is, generally, many turnover events occur at each site each second at high pressures and this may be beyond the time resolution of STM. In the same vein it has been demonstrated that diffusion plays an important role in surface reactions, diffusion of both the adsorbate and of surface metal atoms. This will be occurring very rapidly at the high temperatures where catalysis is usually carried out. Furthermore, STM produces a high electric field at the surface which can perturb the course of the reaction, if it is carried out du_ring the reaction process itself. Nevertheless, it is clear that STM is an extremely useful tool for studying surface reactions and gives us additional, atomic level information about the nanoscale mechanism. It is also the case that in the near future forays closer to real catalytic conditions of high temperature and pressure will be made, though its use as a truly in-situ, non invasive technique is rather further from establishment.

REFERENCES .

2. o

H.S. Taylor, Proc.Roy.Soc. A108 (1925) 105. See, for example, R.W. Joyner in "Elementary Reaction Steps in Heterogeneous Catalysis", eds. R.W.Joyner and R. Van Santen (Kluwer, 1993) p.249. See, for example, E. Bithell, R. Doole, M. Goringe, M. Allen and M. Bowker, Phys. Stat. Solid. 146 (1994) 461.

295

o

5. 6. o

o

9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23.

M. Bowker, Q. Guo and R.W. Joyner, Surf.Sci., 253 (1991) 33. E. Schwartz, J. Lenz, H. Wohlgemuth and K. Christmann, Vacuum 41 (1990) 167. F. Leibsle, P. Murray, S. Francis, G. Thornton and M. Bowker, Nature 363 (1993) 706. V. Dhanak, K. Prince, R. Rosei, P. Murray, F. Leibsle, M. Bowker and G. Thornton, Phys.Rev.B. 49(1994)5585. M. Bowker, Q. Guo and R.W. Joyner, Surf.Sci., 280 (1993) 50. F. Chua, Y. Kuk and P. Silverman, Phys.Rev.Letts., 63(1989) 386. D. Coulman, J. Wintterlin, R. Behm and G. Ertl, Phys.Rev.Lett., 64 (1990) 1761. F. Jensen, F. Besenbacher, E. Laegsgaard and I. Stensgaard, Phys.Rev.B41 (1990) 10233. S. Francis, F. Leibsle, S. Haq, N. Xiang and M. Bowker, Surf.Sci., 315 (1994) 284. I. Wachs and R.J. Madix, J.Catal. 53(1978) 208. M.A. Chesters and E. McCash, Spectrochim. Acta 43A (1987) 1625. M. Bowker and R.J. Madix, Surf. Sci., 95 (1980) 190. M. Bowker and P. Pudney, Catal.Letts., 6 (1990) 13. M. Bowker, Surface Rev. and Letts., 1 (1994) 569. F. Leibsle, S. Francis, S.Haq, X. Ning and M. Bowker, Phys.Rev.Letts., 72 (1994) 2569. L. Ruan, I Stensgaard, F. Laegsgaard and F. Besenbacher, Surface Sci., 314 (1994) L873. W. Crew and R.J. Madix, Surface Sci., 319 (1994) L34. A. Carley, P. Davies, M. Roberts and D. Vincent, Topics in Catal. 1 (1994) 35. M. Bowker, R. Hadden, H. Houghton, J. Hyland and K.C. Waugh, J.Catal. 109 (1988) 263. D. Strongin and G.A. Somorjai, in "Catalytic Ammonia Synthesis" ed. J.R. Jennings (Plenum, N.Y. 1991) 133.

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J.W. Hightower, W.N. Delgass, E. lglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

297

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

C a t a l y t i c F o r m a t i o n of C a r b o n - C a r b o n B o n d s in U l t r a h i g h V a c u u m : C y c l o t r i m e r i z a t i o n of A l k y n e s on R e d u c e d TiO2 S u r f a c e s tL G. Pierce, V. S. Lusvardi, and M. & Barteau Center for Catalytic Science and Technology, Department of Chemical Engineering, University of Delaware, Newark, DE 19716 1. ABSTRACT The steady-state reaction of a molecular beam of methylacetylene impinging on a reduced TiO2(001) crystal was examined under UHV (10 -9 to 10 -8 mbar) conditions. Cyclotrimerization to form trimethylbenzene took place at steady-state. The reaction is truly catalytic; more than 20 turnovers could be obtained without deactivation of the single crystal catalyst. Experiments examining the cyclotrimerization rate as a function of pressure showed that, in the region from 3.0 • 10 -9 to 2.4 x 10 -8 mbar, the reaction was approximately first order in alkyne partial pressure. Temperature variation experiments showed a small decrease in rate as temperature was increased, indicating a slightly negative overall activation energy between 400 K and 535 K. A four-step kinetic model of the cyclotrimerization reaction was constructed using information derived from TPD experiments. These results demonstrate that steady-state catalytic studies of carbon-carbon bond forming reactions, characterized by an overall decrease in number of moles, can be performed on single crystal metal oxides under UHV conditions. 2. I N T R O D U C T I O N We have pre~ously demonstrated the stoichiometric cyclotrimerization of a variety of C2--C6 alkynes on reduced TiO2(001) surfaces in UHV using TPD [1,2]. A key result was the identification (on the basis of XPS [3]) of the surface site requirement for this reaction as individual Ti cations in the +2 oxidation state. This surface chemistry appears to be directly analogous to the oligomerization and trimerization of alkynes by various-low valent organometallic complexes in homogeneous catalysis [4-7]. Recognizing those parallels, we have proposed [1,8] an analogous mechanism for alkyne cyclization on reduced TiO2 surfaces which involves the formation of metallacyclopentadiene intermediates by oxidation of surface Ti +2 centers:

Ti+2 /

C

~

+ ~

, i

I

Ti +2

Ti +2

298 Since this scheme regenerates the original coordinatively u n s a t u r a t e d Wi+2 centers upon desorption of the aromatic, it could, in principle, represent a catalytic cycle for heterogeneous alkyne cyclization. The p r e s e n t study reports a test of t h a t h y p o t h e s i s - - t h e feasibility of catalytic cyclotrimerizationmon a reduced TiO2 surface in IfHV. There are relatively few examples of C-C bond formation on solid surfaces under UHV conditions. There are virtually no examples of catalytic C-C bond formation under such conditions. Perhaps the closest precedent for the present studies on reduced TiO2 can be found in the studies of Lambert et al. on single crystal Pd surfaces. Early UHV studies demonstrated t h a t acetylene could be trimerized to benzene on the P d ( l l l ) surface in both TPD and m o d u l a t e d molecular beam experiments [9,10]. Subsequent studies by the same group and others [11,12] demonstrated that this reaction could be catalyzed at atmospheric pressure both by palladium single crystals and supported palladium catalysts. While it is not clear t h a t catalysis was achieved in UHV, these and subsequent studies have provided valuable insights into the mechanism of this reaction as catalyzed by metals, including spectroscopic evidence for the hypothesized metallacyclopentadiene intermediates [10,13,14]. As generally holds for surface science studies, the n u m b e r of s t e a d y - s t a t e experiments examining the catalytic activity of metal oxide single crystal surfaces is quite small relative to t h a t of comparable studies of metallic surfaces [16,17]. Only a handful of reports concerning the catalytic behavior of single crystal metal oxide surfaces exists in the l i t e r a t u r e . K u n g et al. s t u d i e d the catalytic decomposition of 2-propanol [18] and m e t h a n o l [19] over ZnO surfaces to investigate the structure sensitivity of this reaction on the Zn-polar (0001) and O-polar (0001) surfaces. A series of papers published recently by I w a s a w a et al. has examined the decomposition of formic acid on TiO2 (110) surfaces [20-22], focusing on the dependence of dehydration and dehydrogenation selectivities on pressure and temperature. It is i m p o r t a n t to note t h a t these studies deal with decomposition reactions. In contrast, we report the first example of the catalytic assembly of carbon-carbon bonds on an oxide surface under UHV conditions. m

EXPERIMENTAL All experiments were carried out in a VG ESCALAB i n s t r u m e n t described previously [23,24]. Volatile products were detected w i t h UTI 100 C m a s s spectrometer. The ionizer of the mass spectrometer was enclosed in a quartz envelope with a 5 mm aperture for sampling gases from the vacuum. The TiO2 (001) surface was cleaned and reduced by cycles of ion b o m b a r d m e n t as previously described [3]. The distribution of t i t a n i u m oxidation s t a t e s was determined from curve fitting the Ti(2p3/2) envelope in x-ray photoelectron spectra [3]. After surface preparation, reaction e x p e r i m e n t s were conducted in e i t h e r the TPD or steady state mode. TPD experiments have been described [ 1]. XPS spectra were also obtained following a saturation exposure of the sample using the same procedure as t h a t for the TPD experiments. After pump down, the crystal was placed under the Mg X-ray source and the Ti(2p), O(ls), and C(ls) regions were scanned. For steady-state experiments a dosing needle was aligned perpendicular to the axis of the mass spectrometer. It was used to direct a steady b e a m of methylacetylene (Linde, 95%) at the crystal surface when the sample was placed at the aperture of the mass spectrometer. Steady state reaction experiments were

299 carried out with the sample normal rotated slightly (ca. 15 ~) from the axis of the mass spectrometer; the orientation of the doser was thus about 15 ~ from the plane of the surface. In steady-state experiments, a steady flux of r e a c t a n t was first established with the sample retracted from the doser; the sample was then inserted into the beam and the reaction products were monitored as a function of time using the mass spectrometer. RESULTS S t e a d y - s t a t e molecular beam studies of the reaction of m e t h y l a c e t y l e n e on reduced TiO2 (001) surfaces were u n d e r t a k e n to determine w h e t h e r this reaction could be p e r f o r m e d c a t a l y t i c a l l y u n d e r UHV conditions. A r e p r e s e n t a t i v e experiment is presented in Figure 1. Prior to each experiment, the surface was s p u t t e r e d a n d annealed to a t e m p e r a t u r e between 400 K and 550 K; surfaces p r e p a r e d in this m a n n e r have the highest fraction of Ti(+2) sites (ca. 30% of all surface cations) of any surface we have been able to create by initial sputtering [3]. Thus these are the surfaces most active for cyclotrimerization in TPD experiments [1]. S t e a d y - s t a t e production of t r i m e t h y l b e n z e n e (as indicated by the m/e 105 signal detected by the mass spectrometer) was characterized by behavior typical of more "traditional" catalysts: a j u m p in activity upon initial exposure of the crystal to the molecular beam, followed by a decay to a lower, constant level of activity over a longer time scale. Experiments of up to 6 hours in duration showed

350 m 0rjr ~A C

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250 200 2 c t 15o ',_ I t T i O f molecular beam

100

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Time (minutes) F i g u r e 1. T r i m e t h y l b e n z e n e production from reduced TiO2 (001) surfaces. Production of t r i m e t h y l b e n z e n e quantified by monitoring m/e 105 using m a s s spectroscopy. Surface annealed to 530 K prior to experiment (P = 8x10 -9 mbar, T = 400K).

300 t h a t the crystal retained a steady level of activity for the production of trimethylbenzene; no experiment was able to completely deactivate the catalyst. To insure the activity for production of trimethylbenzene was indeed due to the catalyst (and not the r e s u l t of an i m p u r i t y in t h e m e t h y l a c e t y l e n e beam, for example), experiments with surfaces previously shown to be inactive for the cyclotrimerization chemistry [1] were undertaken. Identical procedures on a fully oxidized surface (produced by prior annealing to 850 K [3]) resulted in no steady-state production of trimethylbenzene, as expected. These experiments on fully oxidized surfaces demonstrate that Ti(+4) cations are inactive for alkyne cyclotrimerization The r a t e of trimethylbenzene production was quantified by t a k i n g the area between the baseline obtained when the crystal was moved out of the molecular beam and the signal during steady-state reaction (Figure 1). Quantification of the amount of trimethylbenzene produced with respect to the number of surface sites is necessary to determine conclusively whether this reaction is indeed catalytic (i.e., whether more than one turnover occurs per surface site). Methylacetylene TPD experiments were run before the steady-state experiments were initiated; as discussed previously, these indicated t h a t 87% of the surface Ti(+2) cations produced trimethylbenzene [1]. Using the m/e 105 peak area from the TPD experiment, the mass spectrometer signal corresponding to one t u r n o v e r of the Ti(+2) sites on the reduced surface could be established, and the turnover frequency calculated from the steady-state data. By this measure, the reduced TiO2 (001) surface is a successful cyclotrimerization catalyst; for the m o s t l e n g t h y experiment, over 20 turnovers of the catalyst were achieved. This method of calculating turnover frequency implicitly assumes t h a t only Ti(+2) sites are active for cyclotrimerization during the TPD and steady-state experiments, and that all Ti(+2) sites are active. It is conceivable that only a small number of Ti(+2) sites are active for methylacetylene a d s o r p t i o n and cyclotrimerization. To test this assumption, C(ls) XPS spectra to quantify coverages following methylacetylene adsorption on the reduced TiO2 (001) surface were obtained with the same experimental apparatus, along with similar experiments involving benzaldehyde adsorption [25]. Since the saturation coverage of benzaldehyde on the reduced surface at 300 K is known from previous work (0.625 m o n o l a y e r s ) [26], methylacetylene coverages can be calibrated based on relative C(ls) peak areas, and compared with the Ti(+2) population from the Ti (2p3/2) spectrum. These experiments indicated that, within experimental error (10-20%), all Ti(+2) cations were active for adsorption and reaction of methylacetylene in TPD experiments. Experiments examining the variation in trimethylbenzene production over a decade of pressure (3.0• -9 to 2.4• -8 mbar) were undertaken. Since the area derived from the mass spectrometer signal is proportional to reaction rate, a logarithmic plot of this quantity vs. the pressure should yield the reaction order as the slope (Figure 2). The fitted slope of 0.89 suggests t h a t the rate is essentially first-order over the p r e s s u r e r a n g e of our e x p e r i m e n t s . The v a r i a t i o n in trimethylbenzene production with respect to temperature was also examined. The production of trimethylbenzene was weakly dependent on the surface temperature and actually decreased slightly as the t e m p e r a t u r e of the crystal was increased between 400 K and 535 K. The a p p a r e n t activation energy was -4 kcal/mol. It is interesting to note that the turnover frequency observed for our single crystal catalyst (1-2 x 10-3 sec-1) is at the lower end of the range typically observed for commercially viable catalytic reactions at higher pressures (10 -3 to 1 sec -1) [16].

301 100 m

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Figure 2. Logarithmic plot of the trimethylbenzene production dependence on chamber pressure. Surface annealed to 500 K prior to experiment (P = 3.0x10 -9 to 2.4x10 -8 mbar, T=400K). The observation of negative apparent activation energy can most simply be interpreted in terms of the competition between the adsorption and desorption of methylacetylene on the surface. This qualitative explanation is illustrated in Figure 3, where the steady-state production of trimethylbenzene is compared with the TPD trace of methylacetylene. The fall off in steady state cyclotrimerization rate matches the tail of the desorption spectrum and illustrates the role of reactant desorption at higher temperatures controlling the availability of alkyne monomers and thus the overall cyclotrimerization rate in this temperature/pressure regime. Relative changes in the population of titanium cations in different oxidation states have been previously observed upon adsorption of oxygenates, such as formic acid and benzaldehyde [26,27]. TPD results have shown t h a t trimethylbenzene production from methylacetylene at 400 K was desorption limited [1], therefore if this reaction proceeds to completion at lower t e m p e r a t u r e s the oxidation state of the surface should be unchanged upon adsorption of methylacetylene. The Ti(2p) region of the XPS spectrum of a surface sputtered with Ar ions for 1 h, followed by annealing at 575 K for 30 rain, is compared with that of the same surface after saturation adsorption of methylacetylene in Figure 4. The Ti(2p) region was unchanged upon adsorption of methylacetylene at 265 K and subsequent flashing of the surface to 575 K also had no effect. The production of trimethylbenzene from methylacetylene did n o t produce a net change of the surface oxidation state, consistent with the catalytic cycle suggested above. The C(ls) spectra of trimethylbenzene and methylacetylene on the reduced TiO2 surface are shown in Figure 5. The spectrum of trimethylbenzene, curve a, contains a peak centered at 284.7 eV with an extended tail at higher binding energy.

302

Figure 4. XPS Ti(2p) core level after methylacetylene adsorption at 265 K on the 575 K-annealed surface.

303

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Binding Energy Figure 5. XPS C(ls) core level after a) trimethylbenzene and b) methylacetylene adsorption at 265 K on the 575 K annealed surface: c) surface in b) flashed to 575K. The C(ls) envelope can be assigned to a ring carbon peak at 284.7 and a methyl group peak at 286 eV; both peaks had a fwhm of 1.7 eV. The C(ls) area for the ring carbon p e a k was 2.5 times g r e a t e r t h a n the area for the methyl group, in a g r e e m e n t with the expected ratio of 2. Methylacetylene adsorption produced a broad symmetric peak centered at 284.7 with a fwhm of 3.0 eV. Flashing the surface to 575 K, curve c, essentially removed all carbon-containing species from the surface. The C(ls) envelope following methylacetylene adsorption may be assigned to three peaks, two of which are due to trimethylbenzene, and the other to unreacted methylacetylene, as deduced from TPD results. Fixing the relative areas, as m e a s u r e d from curve a, to produce the C(ls) envelope for trimethylbenzene, a single peak remains at 283.8 eV with a fwhm of 1.7 eV for the unreacted monomer, as shown in Figure 5, curve b. The C(ls) area of the trimethylbenzene peaks was four times of the monomer, consistent with the yields obtained previously from TPD experiments [1].

DISCUSSION The cyclotrimerization of alkynes to aromatic compounds, observed to occur efficiently on reduced TiO2 (001) surfaces during TPD experiments, can also be carried out as a genuinely catalytic reaction at low pressures. A kinetic model of the c y c l o t r i m e r i z a t i o n reaction describing the p r e s s u r e a n d t e m p e r a t u r e dependence of the behavior observed was constructed (Scheme 1).

304

(1)

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(2)

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(4)

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D:

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Ti(IV)

S: Ti(ll)

Scheme 1. Proposed kinetic model for the cyclotrimerization of alkynes to aromatic compounds on reduced TiO2 (001) surfaces. Acetylene is used as the representative alkyne for clarity. This model comprises three steps to describe the addition of each alkyne, followed by desorption of the final product. The resulting rate equation is: rate =

k]k2k3k4 A3 CI A3 +C2 A2 +C3A

C 1 = klk2k 3 C 2 = klk2k 4 + klk3k 4 + k2k3k 4

C 3 = k_]k3k 4 A pre-exponential factor and activation energy for each rate constant m u s t be established. All forward rate constants involving alkyne adsorption (kl, k2, and k3) are assumed to have equal pre-exponential factors specified by the collision limit (assuming a sticking coefficient of one). All adsorption steps are assumed to be nonactivated. Both desorption constants (k-1 and k4) are a s s u m e d to have preexponential factors equal to 1013 sec -1, as expected from t r a n s i t i o n - s t a t e theory [28]. Both desorption activation energies (26.1 kcal/mol for methyl acetylene and 25.3 kcal/mol for trimethylbenzene) were derived from TPD results [1]. Based on the above rate equation and parameters, the kinetic model predicts the behavior shown in Figure 6 with no fitting of the rate constants. Although the shapes of the experiment and model curves are similar, the m e a s u r e d rate increases at a lower pressure than is predicted by the model. Better agreement between experiment and the model can be generated by offsetting the data, here done simply by multiplying the experimental pressure by five. This disagreement can be understood considering two opposing factors not incorporated into the model. The first is the multiplying effect of the amount of methylacetylene i m p i n g i n g on the surface caused by introducing it through a needle. As compared to the chamber

305 0.005

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-9

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0.001

L_

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-

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Alkyne Pressure (mbar) Figure 6. Comparison of kinetic model predictions and experimental rates from the p r e s s u r e v a r i a t i o n e x p e r i m e n t (400 K). The d a s h e d line r e p r e s e n t s the experimental data offset by multiplying the pressure by a factor of five. pressure, the enhancement due to the concentrating effect of the dosing needle is usually taken to be between 10 and 100. On the other hand, sticking coefficients of less t h a n unity would shift the curve in the other direction, and values of 0.01 to 0.1 are not uncommon. The amount of shifting of this curve would be the product of these two effects, and it is quite reasonable that the combined result would be of the order needed to reconcile model and experiment. It is i m p o r t a n t to note, however, t h a t the turnover frequencies generated from experimental data and the kinetic model, although independently derived, were in excellent agreement. This illustrates t h a t low pressure reaction studies can successfully probe kinetics on surfaces in regions not accessible by other means. 3. C O N C L U S I O N S The cyclotrimerization of alkynes can be successfully carried out at steady-state catalytic conditions on reduced TiO2 (001) surfaces previously shown to be active for this reaction during TPD experiments. As a catalyst, the reduced surface is effective, producing over 20 turnovers without undergoing deactivation. Having established the viability of the cyclotrimerization reaction as a genuine catalytic reaction, experiments examining the pressure and temperature dependence on the cyclotrimerization rate were undertaken. At the pressures examined (10 -8 mbar), the rate is approximately first order with respect to pressure. The rate decreases slightly with increasing temperature, implying a small negative overall activation

306 energy. Utilizing these results, as well as information from methylacetylene TPD experiments, a simple kinetic model was constructed. The cyclotrimerization of methylacetylene over reduced TiO2(OOD surfaces is the first example of catalytic carbon-carbon bond formation on an oxide single crystal surface under UHV" conditions. 4. A C K N O W L E D G E M E N T

We are grateful for the support of the National Science Foundation (Grant CTS 9410965) for this research. 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.

Pierce, K. G.; Barteau, M. A. J. Phys. Chem. 1994, 98, 3882. Pierce, tL G.; Barteau, M. A. Surf. Sci. 1995, 326, L473. Idriss, H.; Barteau, M. A. Catal. Lett. 1994, 26, 123. Sonagashira, tL; Hagihara, N. Bull. Chem. Soc. Jpn. 1966, 39, 1178. Wakatsuki, Y.; Kiramitsu, T.; Yamazaki, H. Tetrahedron Lett. 1974, 4549. Yamazaki, H.; Hagihara, N. J. Organomet. Chem. 1967, 7, P22. Collman, J. P.; Kang, J. W.; Little, W. F.; Sullivan, M. F. Inorg. Chem. 1968, 7, 1298. Pierce, tL G.; Barteau, M. A. J. Mol. Catal. 1994, 94, 389. Tysoe, W. T.; Nyberg, G. L.; Lambert, R. M. J. Chem. Soc., Chem. Commun. 1983, 623. Tysoe, W. T.; Nyberg, G. L.; Lambert, R. M. Surf. Sci. 1983, 135, 128. Rucker, T. G.; Logan, M. A.; Gentle, T. M.; Muetterties, E. L.; Somorjai, G. A. J. Phys. Chem. 1986, 90, 2703. Ormerod, R. M.; Lambert, R. M. J. Chem. Soc., Chem. Commun. 1990, 1421. Ormerod, R. M.; Lambert, R. M. J. Phys. Chem. 1992, 96, 814. Patterson, C. H.; Mundenar, J. M.; Timbrell, P. Y.; Gellman,/L M.; Lambert, R. M. Surf. Sci. 1989, 208, 93. Lee, A. F.; Baddeley, C. J.; Hardacre, C.; Ormerod, R. M.; Lambert, R. M.; Schmid, G.; West, H. J. Phys. Chem. 1985, 99, 6096. Campbell, C. T. Adv. Catal. 1989, 36, 1. Rodriguez, J. A.; Goodman, D. W. S u ~ Sci. Rep. 1991, 14, 1. Berlowitz, P.; Kung, H. H. J. Am. Chem. Soc. 1986, 108, 3532. Vest, M. A.; Lui, tL C.; Kung, H. H. J. Catal. 1989, 120, 231. Onishi, H.; Aruga, T.; Iwasawa, Y. Shokubai 1992, 34,428. Onishi, H.; Aruga, T.; Iwasawa, Y. J. Am. Chem. Soc. 1993, 115, 10460. Onishi, H.; Aruga, T.; Iwasawa, Y. J. Catal. 1994, 146, 557. Idriss, H.; Kim, tL S.; Barteau, M. A. S u ~ Sci. 1992, 262, 113. Peng, X. D.; Barteau, M. A. Langmuir 1989, 5, 1051. Lusvardi, V. S.; Pierce, tL G.; Barteau, M. A., unpublished results. Idriss, H.; Pierce, tL G.; Barteau, M. A. J. Am. Chem. Soc. 1994, 116, 3063. Idriss, H.; Lusvardi, V.S.; Barteau, M. A. Surf. Sci. (in press). Dumesic, J. A.;Rudd, D. F.; Aparicio, L. M.; Rekoske, J. E.; Trevino, A. A. The Microkinetics of Heterogeneous Catalysis, American Chemical Society, Washington, DC 1993; p. 40.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996Elsevier Science B.V. All rights reserved.

307

A n e w a p p r o a c h to u n d e r s t a n d i n g the R o c h o w p r o c e s s : s y n t h e s i s of m e t h y l c h l o r o s i l a n e s f r o m C H 3 + C1 m o n o l a y e r s on Cu3Si in v a c u u m D.-H. Sun, A. B. Gurevich, L. J. Kaufman and B. E. Bent Department of Chemistry, Columbia University, New York, NY 10027 A. P. Wright and B. M. Naasz,

Dow Coming Corporation, Midland, MI 48686

1. INTRODUCTION The Rochow Process refers to the synthesis of liquid methylchlorosilanes from metallurgical silicon powder and gaseous methyl chloride: CH3C1 + Si --) (CH3)xSiCl~x The industrial process, typically conducted in fluid beds operating in the temperature range of 540 - 585 K, forms the basis of the several billion dollar per year silicone industry [2]. A remarkable feature of this process is that the addition of catalytic amounts of copper together with trace promoters such as tin, zinc, and aluminum [1, 2], is required for selective formation (>90%) of dimethyldichlorosilane in preference to other methylchlorosilanes [1]. While it is generally accepted that the dominant bulk solid phase present at reacting surface regions of Si particles is Cu3Si [3, 4], the nature of the surface active site, the details of the surface reactions, and the role of promoters in this process have been difficult to determine. In this paper, we disclose a new approach for applying vacuum surface analysis techniques to understand the Rochow Process. Prior attempts to study the molecular details of this reaction in vacuum on polycrystalline Cu3Si have been hampered by desorption of physisorbed CH3C1 before the temperature can be raised high enough to effect dissociative chemisorption [5]. In the present study, the sluggish C-C1 bond scission step is circumvented by adsorbing methyl radicals and chlorine separately onto cold Cu3Si surfaces. We find, as described below, that these methyl + chlorine monolayers are active in forming methylchlorosilanes. Furthermore, studies of samples with and without promoters show changes in activity and selectivity which parallel those found over real catalysts, and the results are beginning to show how these additives influence the catalytic process.

308 2. EXPERIMENTAL The Cu3Si samples used in this study were 1 a n 2 wafers, 2 mm thick, which were cut from alloy bars prepared at Dow Coming Corporation. Two sets of samples have been studied: pure Cu3Si (12.23 wt% silicon) and promoted Cu3Si samples which contained 15.88 wt% silicon (23% which is free silicon and 77% which is contained in the Cu3Si alloy phase) along with tin (0.01 wt%), aluminum (0.15 wt%) and zinc (0.50 wt%) [2]. Analysis of the promoted samples by scanning Auger, scanning electron microscopy, and polarized light optical microscopy showed Cu3Si grains (- 100 ~tm) embedded in a silicon matrix. Reactor studies on promoted samples which were ground to < 200 mesh showed that these promoted materials w e r e c a t a l y t i c a l l y active w i t h a m a x i m u m selectivity of 84% to dimethyldichlorosilane. Monolayer adsorption and reaction studies were performed in an ultra-high vacuum (UHV) chamber equipped with capabilities for Auger electron spectroscopy (AES) and temperature programmed reaction (TPR) studies [6]. The Cu3Si alloy wafers were mounted on a molybdenum button heater that could be resistively heated to 800 K and which could be cooled by liquid nitrogen to 100 K. The surface temperature was measured by a chromel-alumel thermocouple junction that was either spot-welded directly on the alloy surface or spot-welded to a V-shaped tantalum strip wedged into a slot on the side of the alloy. The difference between the two means of temperature measurement was less than 15 K at 200 K as evidenced by the r e p r o d u c i b i l i t y of the m o l e c u l a r d e s o r p t i o n t e m p e r a t u r e for dimethyldichlorosilane. The surfaces were cleaned by argon ion bombardment (30 mA and 3 kV) at temperatures between 100 and 500 K. As documented in prior studies [7], the Cu/Si Auger peak ratio decreases by a factor of two for surface temperatures above 700 K, and the results here show that there are also significant changes in the extent of surface segregation near room temperature. The change in surface composition near room temperature has a particularly significant effect on some of the surface reactions reported in this paper. Methyl radicals were generated in the gas phase by pyrolysis of azomethane [8] and adsorbed onto Cu3Si surfaces held at 180 K. Previous studies have identified methyl groups on copper surfaces by dosing methyl radicals from a similar source [9]. In the studies here, some nitrogen was also detected on the surface after methyl radical dosing. On the basis of the N(379 eV) : C(272 eV) Auger peak ratio in physisorbed azomethane multilayers (which is indicative of 1 : 1 C to N stoichiometry), the N : C atomic ratio on the surface after dosing is ~- 0.25. Chlorine atoms were generated on the surface by the dissociative adsorption of C12. The surface reactions were investigated by a combination of AES and TPR studies, with authentic samples of the expected products being used to calibrate the signals from the two techniques. The Auger spectra were obtained with a single pass cylindrical mirror analyzer operated with a modulation amplitude of 4 eV; the surface heating rate in the TPR studies was 2 K/s.

309 3. RESULTS AND DISCUSSION In this section, we will first demonstrate the formation of methylchlorosilanes from CH3 + C1 monolayers on Cu3Si surfaces. The effects of promoters and the effect of surface segregation on the reaction rate and selectivity are discussed subsequently. 3.1. Methylchlorosilane formation from CH3 + C1 m o n o l a y e r s After dosing methyl radicals and chlorine molecules onto Cu3Si samples which were cooled to -- 180 K, mass spectrometry was used to identify the gas phase reaction products upon heating. The silane products have been identified by monitoring their characteristic ions, which include SIC14+ (m/e=168), CH3SiC13 + (m/e=148), SIC13+ (m/e=133), (CH3)2SIC12+ (m/e=128), CH3SiC12 + (m/e=113), (CH3)2SiCI + (m/e=93), Si(CH3)3 + (m/e=73). All of these ions are detected. On the other hand, no CH3C1 + (m/e=53) or Sill4 + (m/e=32) are observed. The relative product yields depend on the CH3 to C1 ratio on the surface. In the studies reported here, this ratio has been adjusted to I : 1 (consistent with the CH3 : C1 stoichiometry in CH3C1) on the basis of a C1(181 eV) : C(272 eV) Auger peak ratio of 6.5 which is the same as that m e a s u r e d for p h y s i s o r b e d monolayers of dimethyldichlorosilane. Monolayer coverages of CH3 + C] having I : 1 stoichiometry were obtained by a 20 L exposure from the methyl radical source (approximately saturation coverage) followed by a 9.5 L dose of C12. The dominant gas phase products from these CH3 + C1 monolayers are methylchlorosilanes [(CH3)3SiC1, (CH3)2SIC12, and CH3SiC13] along with some trimethylsilane (see Section 3.3). Figure 1 s h o w s TPR spectra of the methylchlorosilane products after adsorbing, at 180 K, methyls and chlorine on a pure Cu3Si sample which was prepared by ion bombardment at 330 K. In these spectra, the relative mass spectrometer sensitivities have been taken into account (via spectra of standard samples), and contributions from cracking fragments have been accounted for so that the heights of the TPR curves in Figure I reflect the relative rates at which the indicated products are formed. The evolution of methylchlorosilanes between 450 and 600 K is consistent with the 550 - 600 K typical for the catalytic Rochow Process [3]. It is also reasonably consistent with the evolution of methylchlorosilanes at 500 - 750 K reported by Frank and Falconer for a temperature p r o g r a m m e d reaction study of the monolayer remaining on a Cu3Si surface after catalytic formation of methylchlorosilanes from CH3C1 at higher pressures [5]. Both of these observations suggest that the monolayer formed by methyl and chlorine adsorption on pure Cu3Si is similar to that present on active catalysts. For reference, methylchlorosilanes bond quite weakly to the surface and desorb at 180 - 220 K. It can thus be concluded that the rate-determining step in the evolution of methylchlorosilanes at 450 - 600 K is a surface reaction rather than product desorption. By combining the AES and TPR results, it is found that (20 + 10)% of the adsorbed methyl groups (assuming that all adsorbed carbon exists as methyl groups at the 180 K adsorption temperature) form methylchlorosilanes, -5% form trimethylsilane, and the remaining 50 - 75% decompose on the surface to deposit

310

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300

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350

400

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450

I

I

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500

550

600

650

TEMPERATURE (K) Figure 1. Relative rates of methyl chlorosilane formation from a methyl + chlorine monolayer on a Cu3Si surface. The Cu3Si surface was prepared by ion bombardment at 330 K, and a I : 1 ratio of methyl groups and chlorine atoms were reacted in these studies.

carbon. It should be emphasized, however, that carbon deposition likely occurs after the surface CH3 and C1 coverages have b e g u n to decrease as a result of methylchlorosilane formation. By contrast, in steady state catalytic processes, the surface coverages remain constant (probably at or near monolayer saturation). As a result, one should exercise caution in comparing directly the relative yields from these monlayer adsorption studies with catalytic selecfivities. However, trends in rate and selectivity as a function of promoters and surface composition in the monolayer studies are probably relevant for the catalytic process.

311 The reactions of CH3 radicals and C12 alone with Cu3Si have also been investigated. On pure Cu3Si, the dominant silane product from CH3 adsorption is SiH(CH3)3 and the temperature at which the surface is sputtered prior to methyl adsorption has a dramatic effect on the reaction rate (see section 3.3). The C12 reaction gives SIC14 evolution, and the reaction t e m p e r a t u r e is close to that for methylchlorosilane formation. 3.2. Effect of promoters on reaction rates and selectivities

The products from the reaction of CH3 + C1 monolayers on the promoted Cu3Si surface are the same as those for pure Cu3Si, but both the absolute rates and the selectivities are significantly different. In experiments analogous to those described in section 3.1, methylchlorosilanes are evolved from the p r o m o t e d Cu3Si surface between 300 and 450 K. This temperature is 200 K lower than that from the pure Cu3Si surface. This 200 K difference in reaction temperature corresponds to a difference of six orders of magnitude in rate (if the rates are extrapolated to a common reaction temperature of 500 K assuming standard and equivalent pre exponential factors for the reactions on these two surfaces [10]). The reaction selectivity, i.e. the percentage of each methylchlorosilane relative to the total yield of methylchlorosilanes, has been determined from the TPR peak areas of the major cracking fragments using the relative mass spectrometer sensitivity factors determined from studies of authentic samples [11]. These selectivities are compared with those for a typical commercial catalyst [12] in Figure 2. As shown, the 84% selectivity to (CH3)2SIC12 on promoted Cu3Si is close to the commercial value of 90%. On pure Cu3Si, however, the selectivity to (CH3)2SIC12 is only 58%. This decrease in selectivity is not unexpected for a Cu3Si sample without promoters on the basis of catalytic results [1]. For reference, it should be noted that the equilibrium concentration of dimethyldichlorosilane in a methylchlorosilane mixture containing equal numbers of methyl and chlorine is 88% [13] as a result of reversible reactions such as that shown below: 2(CH3)2SIC12 ~

(CH3)SiCI 3 + (CH3)3SiCI

The addition of promoters (and excess Si) to Cu3Si also has a significant effect on the reactivity of pure C1 monolayers formed on this surface by the adsorption of C12. We find that on the promoted samples, C1 monolayers produce SIC13 rather than SIC14. The identification of SIC13 as the chlorosilane product is based on the detection of SIC13+ and the absence of SIC14+ or higher chlorosilane ions in mass spectrometric studies. A second effect of promotion is to lower the chlorosilane evolution temperature. Specifically, on promoted samples the SIC13+ peak m a x i m u m in TPR studies is 390 K vs. --500 K for u n p r o m o t e d samples. Thus, the rates of both chlorosilane and methylchlorosilane formation are dramatically enhanced on the promoted vs. unpromoted samples. By contrast, the rate of trimethylsilane formation when methyls alone are adsorbed on the surface is relatively unaffected by the presence or absence of promoters. In this regard, it should be noted that Falconer and c o w o r k e r s h a v e p r e v i o u s l y s u g g e s t e d that the active surface sites for methylchlorosilane formation involve surface Si-C1 species [14].

312

Figure 2. Comparison of the relative yields of methylchlorosilanes from methyl + chlorine monolayers on Cu3Si and doped Cu3Si samples with typical yields for the catalytic Rochow process.

3.3. Effect of surface composition on reaction activity In all of the studies described above, the Cu3Si samples were prepared by ion bombardment at 330 K followed by cooling of the surface to 180 K before adsorbing the methyl radicals and chlorine. AES studies as well as ion scattering results in the literature [7, 15] show that this procedure produces a surface that is enriched in silicon compared with the Cu3Si bulk stoichiometry. We have found that surfaces with less Si enrichment (possibly even copper enriched relative to the bulk stoichiometry) can be prepared by ion bombardment at temperatures below 300 K. Specifically, Cu(60 eV)/Si(92 eV) Auger peak ratios of 1.2 - 1.7 compared with a ratio of 0.5 at 400 K can be obtained by sputtering at 180 K. In the case of unpromoted Cu3Si surfaces, the effect of this copper enrichment on methylchlorosilane formation appears to be relatively minor. The majority of methylchlorosilanes are evolved at 400 - 650 K as on the unpromoted surface. There is, however, a small yield of methylchlorosilanes with a peak temperature of -370 K. By contrast, for trimethylsilane formation from pure methyl monolayers, copper enrichment by low temperature sputtering shifts the dominant product peak from

313 580 K to 280 K. This dramatic effect is shown in Figure 3 for the case of samples sputtered at 330 K vs. 160 K, and in these studies the evolution of trimethylsilane is monitored via m / e = 73 as shown. This 580 to 280 K decrease in reaction t e m p e r a t u r e for trimethylsilane formation corresponds to a decrease in the reaction activation energy from -33 k c a l / m o l to -16 kcal/mol. Alternatively, if the rates of these processes could be measured at a common temperature, they would differ by more than 5 orders of magnitude. The fact that trimethylsilane is evolved at either 580 K or 280 K and not at temperatures in between suggests that there are distinctly different active sites for forming this product. The ratio of these active sites is a function of the temperature at which the surface is ion bombarded, and the transition from high to low temperature product evolution correlates directly with a factor of 1.5 to 2 increase in the Cu/Si

Sputtered at 330 K

Sputtered at 160 K

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200

300

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Figure 3. Temperature-programmed reaction spectra monitoring m / e = 73 (a cracking fragment of trimethylsilane) after adsorbing a saturation coverage of methyl groups on Cu3Si surfaces prepared by ion bombardment at 160 K (dotted curve) and 330 K (solid

curve).

314 Auger peak ratio. Although the effects of differing surface roughness for these different s p u t t e r i n g temperatures remain to be addressed, the 580 K peak temperature correlates both with the extent of silicon enrichment and also with the temperature to which methyl groups are found to be stable on a Si(100) surface [16]. Furthermore, the 280 K temperature for trimethylsilane formation on the copperenriched surfaces is consistent with the finding the methyl groups are thermally stable to -400 K on copper surfaces [9]. Finally, it should also be noted that early studies [5] have suggested that an important catalytic role for copper is to break up the Si lattice thereby making the Si more available to CH3 and C1 since Cu-Si bonds are weaker than Si-Si bonds. The results here demonstrating the enhanced reactivity for formation of trimethylsilane on copper-enriched surfaces are consistent with this interpretation. 4. CONCLUSIONS The studies presented here show that if a monolayer of methyl groups and chlorine atoms is produced on the surface of a Cu3Si wafer (free of oxygen, carbon and other impurities) then methylchlorosilanes can be formed in 10-20% yield by heating the monolayer to temperatures above 450 K. For a 1:1 ratio of methyls to chlorine, the selectivity for formation of dimethyldichlorosilane is --60%. The ratedetermining step in the evolution of these methylchlorosilanes is a surface-mediated coupling reaction as opposed to product desorption. Both the rate and the selectivity of the process are significantly increased if the Cu3Si samples contain excess Si as well as Sn, Zn, and A1. These promotional effects are consistent with those found for similar additives in the catalytic Rochow process for synthesizing methylchlorosilanes from methyl chloride and silicon. The monolayer studies here suggest that the increase in rate reflects an effect of the promoters on the surface chlorine as opposed to the surface methyl groups. By contrast, it is found that changes in the surface C u / S i ratio have a more significant effect on the reactivity of the surface methyl groups than the surface chlorine. Studies are in progress to further characterize the active sites for these reactions, and we believe that combination of a controlled monolayer approach and UHV surface analysis techniques will lead to even more definitive results in planned future studies. 5. ACKNOWLEDGMENTS Financial support from Dow Coming Corporation and from the American Chemical Society (Grant #CHE 93-18625) is gratefully acknowledged. BEB also gratefully acknowledges support from the Camille and Henry Dreyfus Foundation in the form of a Teacher-Scholar Award.

315

REFERENCES

1. 2. 3. 4. 5.

6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16.

K.M. Lewis and D. G. Rethwisch (eds.), Catalyzed Direct Reactions Of Silicon, Studies in Organic Chemistry 49, Elsevier, Amsterdam, 1993 and articles therein. W.J. Ward, A. Ritzer, K. M. Carroll and J. W. Flock, J. Catal., 100 (1986) 240. K J. H. Voorhoeve, Organosilanes: Precursors to Silicones, Elsevier, Amsterdam, 1967. T.C. Frank, K. B. Kester and J. L. Falconer, J. Catal., 91 (1985) 44. K.M. Lewis, D. McLeod, B. Kanner, J. L. Falconer and T. C. Frank, in: Catalyzed Direct Reactions of Silicon, edited by K. M. Lewis and D. G. Rethwisch, Elsevier, Amsterdam, 1993, 333, and references there in. C.-M. Chiang, T. H. Wentzlaff and B. E. Bent, J. Phys. Chem., 96 (1992) 1836. T.C. Frank and J. L. Falconer, Appl. Surf. Sci, 14 (1982-83) 359. G.H. Smudde, Jr., X. D. Peng, IL Viswanathan and P. C. Stair, J. Vac. Sci. Technol., A 9 (3) (1991). C.-M. Chiang and B. E. Bent, Surf. Sci., 279 (1992) 79. P. A. Redhead, Vacuum, 12 (1962) 203. D.-H. Sun, L. J. Kaufman, B. E. Bent, A. P. Wright, and B. M. Naasz, manuscript in preparation. B. Kanner and K. M. Lewis, in: Catalyzed Direct Reactions of Silicon, edited by K. M. Lewis and D. G. Rethwisch, Elsevier, Amsterdam, 1993, 1. D. R. Weyenberg, L. G. Mahone and W. H. Atwell, Annals of the New York Academy of Sciences, 159 (1969) 38. T. C. Frank and J. L. Falconer, Langmuir, 1 (1985) 104. Y. Samson, J. L. Rousset, G. Bergeret, B. Tardy and J. G. Bertolini, Appl. Surf. Sci., 72 (1993) 373. H. Gutleben, S. R. Lucas, C. C. Cheng, W. J. Choyke and J. T. Yates, Jr., Surf. Sci., 257 (1991) 146.

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Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

317

Ruthenium as Catalyst for Ammonia Synthesis M. Muhler*, E Rosowski, O. Hinrichsen, A. Hornung and G. Ertl Fritz-Haber-Institut der Max-Planck-GeseUschaft Faradayweg 4-6, D- 14195 Berlin (Dahlem), Germany Five Ru-based catalysts were prepared to study the effect of the support and the role of the alkali promoter in NH8 synthesis: Ru/AlzOa, Ru/MgO, Cs-Ru/A1203 and Cs(K)-Ru/MgO. The catalysts were characterized by N2 physisorption, H2 chemisorption and XPS. The absence of chlorine- and sulphur containing compounds turned out to be important for the preparation of highly active catalysts. Power law expressions were derived from conversion measurements at atmospheric pressure and at 20 bar. For all catalysts, the reaction order for H2 was found to be negative suggesting that a Pr% / PrI2 ratio in the feed gas higher than 1 / 3 would be favourable for industrial NHs synthesis at high pressure. The microkinetic analysis of the temperature-programmed desorption and adsorption of N2 and of the kinetics of isotopic exchange demonstrated the enhancing influence of the Cs promoter on the rate of N2 dissociation and recombination. XPS measurements after thorough reduction revealed a shift of the Ru 3d5/2 peak to lower binding energy by about 1 eV in the presence of Cs suggesting an electronic promoter effect. 1. Introduction

Alkali-promoted Ru-based catalysts are expected to become the second generation NHs synthesis catalysts [ 1]. In 1992 the 600 ton/day Ocelot Ammonia Plant started to produce NHa with promoted Ru catalysts supported on carbon based on the Kellogg Advanced Ammonia Process (KAAP) [2]. The Ru-based catalysts permit milder operating conditions compared with the magnetite-based systems, such as low synthesis pressure (70 - 105 bars compared with 150 - 300 bars) and lower synthesis temperatures, while maintaining higher conversion than a conventional system [3]. In spite of the industrial importance, relatively few studies in the catalytic literature deal with the kinetics of NH3 synthesis over supported Ru catalysts [4-8]. These studies focus on the inhibiting influences of PH2 and PNH8 on the rate of NH3 formation. Alkali promotion was found to decrease the inhibition by PNH8 significantly thus causing an increase in the inhibiting ' effect of PH2 [5-7]. The reaction order for N2 was found to be essentially unity indicating that the dissociative chemisorption of N2 is the rate-determining step (rds) in the overall mechanism. On multiply promoted Fe catalysts, H2 has a positive reaction order due to high coverages of adsorbed atomic nitrogen (N-.) [9]. In our laboratory a systematic study is in progress aiming at a detailed understanding of the "Corresponding author

318 catalytic phenomena involved in the synthesis of ammonia on ruthenium. The following catalysts were prepared from Rua(CO)x2 and high-purity supports: Ruthenium supported on MgO (Ru/MgO) and AlzOa (Ru/A12Oa), potassium promoted Ru/MgO (K-Ru/MgO) and cesium promoted Ru/MgO (Cs-Ru/MgO) and Ru/AIzOa (Cs-Ru/AlzOa). The catalysts were characterized by N2 physisorption (BET area), H2 chemisorption and X-ray photoelectron spectroscopy (XPS). This study presents kinetic data obtained with a microreactor set-up both at atmospheric pressure and at high pressures up to 50 bar as a function of temperature and of the partial pressures from which pov~er-law expressions and apparent activation energies are derived. An additional microreactor set-up equipped with a calibrated mass spectrometer was used for the isotopic exchange reaction (IER) 2aN 2 + a~ 2 = 2 29N2 and the transient kinetic experiments. The transient experiments comprised the temperature-programmed desorption (TPD) of N2 and H2. Furthermore, the interaction of N2 with Ru surfaces was monitored by means of temperature-programmed adsorption (TPA) using a dilute mixture of N2 in He. The kinetic data set is intended to serve as basis for a detailed microkinetic analysis of NH3 synthesis kinetics [ 10] following the concepts by Dumesic et al. [ 11].

2. Experimental The catalysts were prepared from high purity A12Os (99.99%, Johnson Matthey) or MgO (Puratronic, 99.996% metals basis, Johnson Matthey) and Rua(CO)I2 (Johnson Matthey) by wet impregnation in a rotary evaporator and subsequent heating in high vacuum following the procedures in ref. [12-14]. Details of the preparation are given in ref. [15]. The achieved metal loading was 5 wt. % Ru. The cesium-promoted catalysts Cs-Ru/MgO and Cs-Ru/A12Oz were obtained by impregnating the Ru/MgO or Ru/Al2Oa catalysts subsequent to heating in vacuum to 723 K with an aqueous solution of CsNOa (99.99 %, Strem). For the preparation of K-Ru/MgO, an aqueous solution of KNOa (99.997 %, Johnson Matthey) was used. The atomic ratios were Cs(K) / Ru - 1 / 1 for Cs(K)-Ru/MgO and Cs / Ru - 3 / 1 for Cs-Ru/A12Os. The Ru metal area was determined by volumetric H2 chemisorption in the quartz U-tube of an Autosorb 1-C set-up (Quantachrome) following the procedure described in ref. [ 16]. Prior to chemisorption, the catalysts were activated by passing 80 Nml/min high-purity synthesis gas (PN2 / Pn2 -" 1 / 3) from a connected feed system through the U-tube and heating to 673 K for alkali-promoted catalysts or to 773 K for alkali-free catalysts with a heating rate of 1 K/rain. The BET area was measured by static N2 physisorption in the same set-up. The kinetic experiments were carried out in an all stainless steel microreactor system with three gas lines which could be operated at pressures up to 100 bar. The gases used had the following purities: Ar 99.9993%, N2 99.9993%, H2 99.9993%. The feed gas was further purified by means of a self-designed guard reactor [17]. Gas analysis was performed by a non-dispersive infrared detector (BINOS, Fisher-Rosemount) which was calibrated by using a reference gas mixture (Linde). 138 mg of the 250/zm-800/zm sieve fraction were used for the kinetic experiments resulting in bed heights of less than 15 mm which prevented limitations by heat or mass transport. The reduction was carried out in synthesis gas using 40 Nml/min with a heating ramp of 1 K/min up to 673 K. The spectroscopic investigations were carried out in a modified LHS 12 MCD system. For the XPS measurements (Mg Kct 1253.6 eV, 240 W power) a fixed analyser pass energy of 108 eV

319 was used resulting in a resolution of 1.1 eV FWHM of the Ag 3d5/2 peak. The binding energy scale was calibrated using EB(AU 4fr/2) - 84.0 eV. The samples were activated in a directly attached preparation chamber (base pressure < 10-Smbar) from which the sample could be transferred into the UHV analysis chamber (base pressure 1.10 - l ~ mbar) within 1 min. The reduction was carried out in 1000 mbar synthesis gas by heating with 2 K/min to 673 K followed by 3 h NHa synthesis at this temperature. The synthesis gas mixture was replaced several times during the reduction. Charging was corrected using Mg 2s at 88.1 eV as internal standard [14]. Quantitative data analysis was performed by subtracting stepped backgrounds and using empirical cross sections [ 18]. 3. Results and Discussion

As expected, 3,-A12Os (BET area 110 mS/g) turned out to be the more stable support with a higher surface area than MgO (BET area 52 mS/g). The BET area of Ru/AlsOa was found to be 104 m2/g after NHa synthesis at 773 K which decreased significantly to 70 mS/g as a result of cesium impregnation. After NHa synthesis at 773 K, the specific area of Ru/MgO was observed to be 25 mS/g compared with 52 mS/g found for the MgO support. Cesium impregnation caused a further decrease in specific area to 23 m2/g. Table 1 Results of the H2 chemisorption measurements after NHa synthesis based on H/Ru = 1/1. NH3 synthesis was run at 773 K with Ru/MgO and Ru/AlsOa, and at 673 K with all alkali-promoted catalysts. The mean particle size was calculated assuming spherical particles. Catalyst H2 monolayer Metal area Dispersion Particle size / ~mol Hs/g / mS/g /% /nm Ru/MgO 130 12.9 53 1.9 Ru/AI203 118 11.7 48 2.1 Cs-Ru/MgO 69 6.8 28 3.6 Cs-Ru/AlsOs 100 9.9 41 2.5

The H2 chemisorption results are summarized in table 1. The Hs monolayer capacities were used to derive Ru metal dispersions and mean particle sizes assuming spherical particles. On both MgO and AlsOa, the impregnation with Rua(CO)12 resulted in mean particle sizes of about 2 nm after NHa synthesis at 773 K. It is remarkable that about the same Ru metal areas were obtained on MgO and A12Oa in spite of the largely differing BET areas of the supports. For the Cs-Ru/MgO catalyst, the amount of chemisorbed hydrogen was found to be reduced by about a factor of two. XRD measurements and TEM images revealed that the decrease in metal area is indeed due to sintering of the Ru metal particles [15]. The Ru/AlsOa catalyst was not significantly affected by the impregnation with CsNOa as shown by the increase in particle size from 2.1 nm to 2.5 nm derived from H2 chemisorption. The results of the conversion measurements at atmospheric pressure using 138 mg catalyst are shown in fig. 1. The following sequence in catalytic activity was observed: Cs-Ru/MgO > Ru/MgO > Cs-Ru/AlsOa > Ru/AlsOa. It is noteworthy that the catalytic activity of the Cs-Ru/MgO catalyst exceeds significantly the catalytic activity of a multiply promoted iron

320 catalyst (trace D in fig. 1A). A recent kinetic study provides evidence that the MgO support acts as alkaline earth promoter creating promoted sites at the interface [19]. The influence of the akali promoter is shown in fig. lB. The traces were obtained with somewhat deactivated catalysts after several weeks of NHa synthesis. Cesium ttmaed out to be a better promoter than potassium in agreement with the results obtained by Aika et al. [14]. The same authors furthermore suggest that Cs promotion is less efficient on A12Os since CsOH interacts mainly with the acidic support whereas on the basic MgO support more CsOH should be in contact with the Ru metal particles [ 14].

0.8

0 >

'I 0 R u / A I 2 0 3 - 0 Ru / MgO t a C s - R u / A i 2 0 3 " n + KNO3 tt & Ru / MgO & + CsNO3 <> Fe catalyst

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%%%

w

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Temperature / K

Figure 1. NHa concentration in the reactor effluent gas using a total flow of 40 Nml/min with

PN2 / PI-I2 - 1 / 3 at atmospheric pressure. Traces A-E in fig.lA (from bottom to top) were obtained with Ru/AlzOa, Cs-Ru/AlzOa, Ru/MgO, a multiply promoted iron-based catalyst, and Cs-Ru/MgO. The corresponding NI-Ia equilibrium concentration is displayed as dashed line. Traces A-C in fig.1B (from bottom to top) were obtained with Ru/MgO, K-Ru/MgO, and Cs-Ru/MgO.

It is known that chlorine acts as severe poison for NHa synthesis [20,21]. Hence recent kinetic studies used chlorine-free Ru precursors like Rus(COh2 [8,22] or Ru(NO)(NOs)a [7]. In addition to chlorine, the presence of sulphur was found to poison Ru catalysts. Fig. 2A demonstrates that both poisons may originate from the Ru precursor. The binding energies for the C1 2p peak and of the S 2p peak observed for Ru prepared form RuO2 are typical for chloride and sulfide anions, respectively [23]. Ru prepared from Rua(CO)x2 was found to have a significantly higher purity. As shown in fig. 2B, sulphur and chlorine impurities can also originate from the support. The XPS data of MgO with a purity of 98 % reveal the presence

321 of chloride and sulphate anions which are essentially absent in MgO with a purity of 99.999 %. Hence it is mandatory to use high-purity Ru precursors and supports to prepare poison-free catalysts.

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d

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I 198.6 e V RU3(00)12 , .... , . . , //L... 200 195 170

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.

.

.

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.

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.

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Binding Energy / eV

Figure 2. XPS C1 2p and S 2p data obtained with RuO2 and Rua(CO)x2 after reduction in synthesis gas at 1 bar and 673 K and transfer in UHV (Fig. 2A, upper half) and with two different MgO supports with purities of 98% and 99.999%, respectively (fig. 2B, lower half).

The quantitative XPS results obtained after reduction in synthesis gas are summarized in table 2. The observed ratios of O / Mg = 1.1 / 1 and O / A1 = 3.4 / 2 are in reasonable agreement with the stoichiometric ratios. The Ru concentration determined by XPS is somewhat higher for Ru/MgO than for Ru/A1203 in agreement with the results obtained with Hz chemisorption. Cs impregnation leads to a stronger decrease in the amount of Ru observed by XPS for Cs-Ru/MgO than for Cs-Ru/A12Oz due to the sintering of the Ru metal particles on MgO. The decrease of the Ru / support ratio observed for both catalysts may also be due to the shielding of the

322 Table 2 Surface composition determined by XPS after reduction in synthesis gas at 1 bar at 673 K and transfer in UHV. Catalyst Cs 3d O ls Ru 3p Mg 2s A1 2s Ru/MgO 50.4 3.3 46.4 Ru/A12Os 61.9 2.1 36.0 Cs-Ru/MgO 5.7 57.9 1.2 35.2 Cs-Ru/Al2Os 6.3 58.8 1.6 33.2

underlying Ru metal particle by a CsOH overlayer. Although a Cs / Ru ratio of 3 / 1 was used for the preparation of Cs-Ru/Al~Os, the ratio observed by XPS is similar to the ratio found for Cs-Ru/MgO. The latter catalyst was prepared with a Cs / Ru ratio of 1 / 1. This result indicates that the excess of Cs used for Cs-Ru/Al2Os has reacted with the bulk forming ternary phases with A12Os. Table 3 Power law exponents r - l~m,-~tt8 .PaN2.P~2 and apparent activation energies as a function of the total pressure determined in the given temperature range. The accuracy of the determination of the power law exponents and of the apparent activation energy is about q-0.1 and -4- 5 kJ/mol, respectively. Catalyst Pressure Temperature range c~(NHa) j3(N2) 7(H2) E,, / bar /K / kJ/mol Ru/MgO 1 513 - 603 -0.3 0.8 -0.3 69 20 573 - 663 -0.3 1.0 -0.5 78 Ru/AlzOs 1 593 - 663 -0.4 0.9 -0.1 70 20 573 - 688 -0.5 0.9 -0.3 76 Cs-Ru/MgO 1 498- 570 0.0 0.7 -0.7 96 20 550 - 630 0.0 0.8 -0.9 109 Cs-Ru/AlzO3 1 543 - 608 0.0 0.7 -0.6 103 20 573 - 663 0.0 0.9 -0.6 101

The results of the conversion measurements are summarized in table 3. The power law exponents and the apparent activation energies were derived following the analysis given in ref. [5]. The reaction orders of NHa and the reaction orders of N~ and H~ were determined by varying the synthesis gas flow between 40 Nml/min and 160 Nml/min and by varying the N2 / H2 ratio between 3 / 1 and 1 / 3 using a total flow of 120 Nml/min, respectively. Both determinations were carded out in the temperature range specified in table 3 ensuring the measurements to be in the kineticaUy controlled regime far from equilibrium. From the data shown in table 3, it is evident that the effect of Cs promotion on the power law kinetics is twofold: First, the reaction order for NHs is changed to essentially zero, and secondly, the apparent activation energy is higher by more than 20 kJ/mol in the presence of Cs. Contrary to the results obtained by Aika et al. [5], the reaction order for H2 was negative for all catalysts investigated. The positive reaction order for H~ reported by Aika et al. [5] for

323 Ru/Al2Oa and Ru/MgO may be due to the presence of chlorine originating from RuCla used for catalyst preparation. Fig. 3A shows the effluent NHa concentration observed for Ru/MgO as a function of reaction temperature for three different PN2 / Pn2 / P A r ratios at 20 bar total pressure. It is obvious that the reaction orders for N2 and H2 have opposite signs. Fig. 3B illustrates that the reaction orders for N2 and H2 partly compensate each other in the kinetically controlled temperature regime. Hence an increase in total pressure with a constant PN2 / Pn2 = 1 / 3 ratio does not lead to a significant increase in conversion at lower temperatures. For the application of alkali-promoted Ru catalysts under industrial synthesis conditions, it is necessary to find a compromise between kinetics and thermodynamics by increasing the PN2 / PH= ratio. The optimum observed for Cs-Ru/MgO prepared from Cs2COa at 50 bar is at about PN2 / Pn2 = 40 / 60 [15]. The high NHa concentration of about 8 % obtained with 0.138 g catalyst using a total flow of 100 Nml/min clearly shows that Ru catalysts have indeed the potential to replace Fe-based catalysts in industrial synthesis [ 15].

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Figure 3. Dependence of the NHa effluent concentrations observed for Ru/MgO on the feed gas composition (fig. 3A, left figure) and on the total pressure. In fig. 3A, trace A was obtained with PN2 / Pn2 / P A r - 1 / 1 / 2, trace B with PN2 / Pn2 / P A r = 1 / 3 / 0, trace C with PN2 / Pn2 / P A r - 3 / 1 / 0, respectively, using a total flow of 120 Nml/min at 20 bar. In fig. 3B, traces A-D (from bottom to top) were obtained at 1 bar, 9 bar, 20 bar and 50 bar, respectively, using a total flow of 40 Nml/min with PN~ / Pn2 - 1 / 3.

The reaction orders for N2 observed for all catalysts were close to 1.0 indicating that the

324 dissociative chemisorption of Nz is the rate-determining step in NHs synthesis. The kinetics of the interaction of Nz with Ru/MgO, Ru/AlzOa and Cs-Ru/MgO have been studied recently by performing N2 TPD and N2 TPA experiments and by determining the rate of isotopic exchange

28N2 + a~

- 2

2aN2 [ 2 4 ] .

Table 4 Rate constants ki - Ai- exp(-Ei / RT) for N2 + 2 , - 2 N - - , . Units of A/are ( torr- s) -~ for the forward reaction and s-1 for the reverse reaction forward rate constant reverse rate constant catalyst preexponential activation energy preexponential activation energy factor (kJ/mol) factor (kJ/tool) Cs-Ru/MgO 7.4.10 ~ 33.0 2.0.10 x~ 137.0 Ru/MgO 7.4.10 ~ 48.0 1.5.101~ 158.0 Ru/AI20s 7.4.10 ~ 60.6 1.5.101~ 158.0 .

.

.

.

.

.

Table 4 summarizes the rate constants ki - - A i 9 exp(-E//RT) for the forward and the reverse reaction derived from our microkinetic analysis of the steady-state and transient experiments with the three catalysts, i.e. Cs-Ru/MgO, Ru/MgO, and Ru/A12Os catalyst [24]. The rate constants in table 4 for Ru/A12Os should be considered as initial rate constants since it was not possible to achieve a higher coverage of N - - , than 0.25. Furthermore, it was not possible to detect TPA peaks for Ru/AlzOs within the experimental detection limit of about 20 ppm. Ru/MgO is a heterogeneous system with respect to the adsorption and desorption of N2 due to the presence of promoted active sites which dominate under NHa synthesis conditions. The rate constant of desorption given in table 4 for Ru/MgO refers to the unpromoted sites [ 19]. The N~ TPD, N2 TPA and IER results thus demonstrate the enhancing influence of the alkali promoter on the rate of N2 dissociation and recombination as expected based on the principle of microscopic reversibility. Adding alkali renders the Ru metal surfaces more uniform towards the interaction with N2. On polycrystalline Ru samples, IR measurements by Aika and Tamaru [25] revealed the influence of the alkali promoter on the stretching frequency of N2 -- * which was interpreted in the frame of a charge transfer mechanism. XPS should be the appropriate technique to detect charge transfer from the promoter to the Ru metal clusters. The Ru 3d spectrum of the Ru/MgO precursor after heating in vacuum to 723 K in order to decompose the adsorbed Ru carbonyl compounds is shown as lower trace in fig.4. The binding energy of the Ru 3ds/~ peak indicates that Ru is not yet reduced to the metallic state. Furthermore, the intensity ratio of the Ru 3da/2 and Ru 3d5/2 peaks shows that significant amounts of carbon compounds are present giving rise to overlapping C ls peaks at about 285 - 290 eV. After reduction (trace in the middle), the binding energy of the Ru 3d5/2 peak was found to be 280.0 eV indicating complete reduction to Ru metal. After reduction of the Cs-Ru/MgO catalyst, the Ru 3d speaks were observed to be shifted by 1 eV to lower binding energy (top trace in fig.4). It has to be noted that the Mg 2s peak had to be used as internal standard (EB(Mg 2s) - 88.1 eV) to correct for charging. However, the MgO bulk should not be affected by cesium impregnation. The XPS shift is influenced by many factors like the extraatomic relaxation energy which might change due to the presence of a Cs+O coadsorbate layer resulting from the decomposition of presumably CsOH as mobile

325

1 eV "~/~

XPS Ru3d

8000

r r

6OOO

r r-

-'-" C: 4000

2000

0 295

I

290

285

' x_..,

I ,

280

275

Binding Energy / eV

Figure 4. XPS Ru 3d data observed for the Ru/MgO catalysts. The Ru 3d spectra (from bottom to top) were obtained with the precursor after heating in high vacuum to 773 K , after reduction in 1 bar synthesis gas up to 773 K, and after impregnation with aqueous CsNOa solution and subsequent reduction in synthesis gas up to 673 K.

species. Since this shift was only observed after thorough reduction in the directly attached preparation chamber with rapid transfer in UHV, it seems plausible to assume that the treatment at 673 K in 750 mbar H2 (purity 99.9999 %) caused a partial reduction of the Cs+O coadsorbate layer thus creating oxygen vacancies which might serve as electron-donating adsorption sites for N2. Further studies are in progress to clarify this hypothesis. 4. Conclusions The preparation of Ru-based catalysts from high-purity supports using Rus(CO)xz followed by impregnation with aqueous Cs solution was shown to result in stable and active NHs synthesis catalysts. Cs-Ru/MgO was found to have a higher catalytic activity at atmospheric pressure than a multiply promoted Fe-based catalyst. Power law expressions were derived from conversion measurements at atmospheric pressure and at 20 bar. For all catalysts, the reaction order for H2 was found to be negative suggesting that a higher PN2 / Px2 ratio in the feed gas than 1 / 3 would be favourable for industrial NHa synthesis at high pressure. Studying the kinetics of the interaction of N2 with the Ru catalysts revealed that the Cs promoter enhances both the rate of dissociative chemisorption and the rate of recombinative desorption. Ru catalysts were found to be rather inactive for NHa synthesis without alkali

326 promotion. Ru/MgO turned out to be a heterogeneous system with respect to the adsorption and desorption of N2 due to the presence of promoted active sites which dominate under NHa synthesis conditions. Adding alkali renders the Ru metal surfaces more uniform towards the interaction with N2. XPS results provide evidence for an electronic promoter effect. REFERENCES 1. S.R. Tennison, in Catalytic Ammonia Synthesis, Plenum Press, New York, (Ed. J.R. Jennings), 1st. ed. (1991 ) 303. 2. EJ. Shires, J.R. Cassata, B.G. Mandelik, C.E van Dijk, U.S. Patent, 4479925 (1984) Oct. 30. 3. T.A. Czuppon, S.A. Knez, R.V. Schneider IN, G. Worobets, Chem. Engineering, March

4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17.

18. 19. 20. 21. 22. 23. 24. 25.

1993, presented at the 1993 AICHE Ammonia Safety Symposium, Sept. 1993, Orlando, Florida, 100, No.3 (1993) 19. ER. Holzman, W.K. Shiflett, J.A. Dumesic, J. Catal., 62 (1980) 167. K. Aika, M. Kumasaka, T. Oma, O. Kato, H. Matsuda, N. Watanabe, K. Yamazaki, A. Ozaki, T. Onishi, Appl. Catal., 28 (1986) 57. H. Bails, M. Glinski, J. Kijenski, A. Wokaun, A. Baiker, Appl. Catal., 28 (1986) 295. J.U. Nwalor, J.G. Goodwin Jr., Topics in Catal., 1 (1994) 285. E Moggi, G. Albanesi, G. Predieri, G. Spoto, Appl. Catal. A, 123 (1995) 145. L.M. Aparicio, J.A. Dumesic, Topics in Catal., 1 (1994) 233. O. Hinrichsen, E Rosowski, M. Muhler, G. Ertl, Chem. Eng. Sci., Proc. of the 14th Int. Symposium on Chem. React. Engineering, (1996) accepted. J.A. Dumesic, D.E Rudd, L.M. Aparicio, J.E. Rekoske, A.A. Trevino, The Microkinetics of Heterogeneous Catalysis, ACS Professional Reference Book, Washington, DC, (1993). H. Kn6zinger, Y. Zhao, B. Tesche, R. Barth, R. Epstein, B.C. Gates, J.E Scott, Faraday Discuss. Chem. Soc., 72 (1982) 53. E Moggi, G. Predieri, G. Albanesi, S. Papadopulos, E. Sappa, Appl. Catal., 53 (1989) L 1. K. Aika, T. Takano, S. Murata, J. Catal., 136 (1992) 126. E Rosowski, A. Hornung, O. Him'ichsen, D. Herein, M. Muhler, G. Ertl, Appl. Catal., (1996) in preparation. R.A. Dalla Betta, J. Catal., 34 (1974) 57. B. Fastrup, H.N. Nielsen, Catal. Lett., 14 (1992) 233. Practical Surface Analysis, John Wiley, Chichester, (Ed. D. Briggs, M.E Seah) , 2nd ed. (1994). E Rosowski, O. Hinrichsen, M. Muhler, G. Ertl, Catal. Lett., (1996) accepted. W.K. Shiflett, J.A. Dumesic, Ind. Eng. Chem. Fundam., 20 (1981 ) 246. S. Murata und K.-I. Aika, Appl. Catal. A, 82 (1992) 1. Y. Kadowaki, S. Murat~ K.-I. Aika, Stud. Surf. Sci. Catal., Elsevier Science Publishers, (Ed. L. Guczi, E Solymosi, E Tetenyi), 75 (1993) 2055. J.E Moulder, W.E Stickle, EE. Sobol, K.D. Bomben, Handbook of X-ray Photoelectron Spectroscopy, Perkin-Elmer, (1992). O. Hinrichsen, E Rosowski, A. Hornung, M. Muhler, G. Ertl, J. Catal., (1996) submitted. K. Aika, K. Tamaru, in Ammonia: Catalysis and Manufacture, Springer-Verlag, Berlin, (Ed. A. Nielsen), 1st ed. (1995).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

327

SURFACE-S~UCTURE-DEPENDE~ REACTION PATHWAYS OF M E T H Y L G R O U P S O N NI(100) and NI(111) S U R F A C E S Richard B. Hall, Miguel Castro*, Chang Min Kim*, Charles A. Mimst Exxon Research and Engineering, Rt. 22E, Annandale, N.J. 08801 *Department of Chemistry, University of Puerto Rico, Mayaguez, Puerto Rico *Department of Chemistry, Kyung Pook National University, Daegu-Si South Korea "tDepartment of Chemical Engineering, University of Toronto, Toronto, Canada 1. Introduction-

Reactions of hydrocarbon fragments such as methyl groups on transition metal surfaces play an important role in a number of catalytic processes, including oxidative coupling of methane, methanation of syngas, steam reforming, and partial oxidation of methane to syngas. Although important, our knowledge of the fundamental mechanisms and kinetics of their reactions is limited. Because of the significance of these reactions, there are an increasing number of investigations of the fundamental reaction steps of small alkyl fragments on well characterized surfaces. Several methods have been used for the preparation of adsorbed alkyl radicals and for the study of their reactions. Most frequently, surface-bound alkyl fragments are produced by decomposing alkyl halides on the metal surface. This approach has been quite successful in characterizing reaction pathways of C1 and C2 fragment reactions on Ni [ 1-3], Pt [ 1] and Cu [46] surfaces. However, it intrinsically involves co-adsorbed halide atoms. Significantly, the co-adsorbed halide limits the coverage (i.e. concentration) of alkyl groups that can be adsorbed on the surface. As we will show, the dominant reaction pathway, at least on nickel surfaces, is a function of surface coverage. High coverages are likely to be more representative of reactions under high pressure, commercial process conditions, and the use of alkyl halide precursors preclude studies of high coverages. Surface methyls have also been synthesized by the collision-induced dissociation of methane physisorbed on Ni(111) surfaces.[7, 8] This approach avoids the effects of coadsorbates other than hydrogen, and a number of aspects of the reaction and decomposition of CH3 and CH fragments on Ni(111) have been determined.J9] However, the method is relatively complex and best suited for study of low coverages. Recently, Stair and coworkers [10, 11] developed a method to produce gas-phase methyl radicals, and used this to study reactions of methyl groups on Pt surfaces [12] and on molybdenum oxide thin films [ 13]. In this approach, methyl radicals are produced by pyrolysis of azomethane in a tubular reactor located inside an ultrahigh vacuum chamber. This method avoids the complications of co-adsorbed halide atoms, it allows higher coverages to be reached, and it allows the study of reactions on oxide and other surfaces that do not dissociate methyl halides effectively. We have adopted this last approach. We report here results from studies of the reaction of methyl groups on two different nickel single-crystal surfaces, Ni (100) and Ni(111). Nickel is

328 of particular interest because it is used as a commercial catalyst in methanation [14], steam reforming [15], and partial oxidation of methane to syngas.[16] These processes have many fundamental reaction steps in comrr~n, and our goal is to understand these fundamental reaction steps and to develop detailed kinetic models of these processes. We report here on some of the mechanistic trends and describe an unexpected dependence of the reaction pathway on metal surface structure. A more quantitative kinetic analysis will be presented in a subsequent paper.

2. ExperimentalThe experiments were carried out in a modified Leybold-Heraeus stainless steel ultrahigh vacuum chamber which has been described in detail elsewhere [ 17]. Briefly, it has a Balzers mass spectrometer for residual gas analysis (RGA) and temperature programmed desorption (TPD) measurements, a hemispherical electron energy analyzer and electron gun for AES measurements, a Varian LEED spectrometer, and a differentially pumped ion gun for sputtering and low energy ion scattering (LEIS). Background pressures below 1 x 10-10 torr are achieved by a combination of turbo molecular and liquid nitrogen cooled titanium sublimation pumps. The single-crystal nickel sample (either (100) or (111) orientation) was mounted on a liquid nitrogen cooled sample manipulator via two tantalum wires spot-welded to the side of the crystal. The substrate temperature was measured using a chromel-alumel thermocouple spotwelded to the side of the crystal. The crystal could be cooled to 110 K and resistively heated to 1200 K. The crystal was cleaned by cycles of Ne+ ion bombardment (primary energy of 1.0 KeV) for 5 minutes and annealing to 1000 K for 10 minutes until most impurities were removed, as detected by AES. Residual carbon was removed by exposing the surface at 300 K to 1 langmuir (L) of oxygen, followed by heating to 800 K. Trace oxygen was removed by reducing with hydrogen (lx10-7 Torr) at 800 K Several cycles were required before no carbon was detected by AES. Methyl radicals were produced by pyrolysis of azomethane (CH3N2CH3). Azomethane was synthesized as described earlier [ 18]. It was purified periodically by freeze-pump cycles at 77 K, and the gas purity verified by RGA. The methyl radical source was similar to that developed by Stair and coworkers.[ 10, 11] The source was made of a quartz tube with 3 mm OD and 1 ram ID, resistive heating was supplied by means of a 0.25 mm diameter tantalum wire wrapped outside the quartz tube. The length of the heating zone was 4 era, recessed from the end of the tube by 1 era. An alumina tube around the outside of the heating zone served as a radiation shield. Azomethane was admitted to the hot tube at a pressure of l x10-8 to l x10-7 Torr via a high-vacuum precision leak valve. The pyrolysis tube was maintained at about 1200 K, adequate to decrease the major peaks in the mass spectrum of the parent azomethane at 58 and 43 ainu by at least a factor of 100. Ideally, pyrolysis of azomethane produces only methyl radicals and N2 through cleavage of the relatively weak methyl-N2 bonds. However, secondary reactions of the methyl radicals are unavoidable. These reactions produce ethane, methane and hydrogen in addition to the methyl and nitrogen. The product distribution was 16% CH3, 16% CI-14, 13% C2H6, 39% N2, and 17% H2 (by mole %), based on mass specmun intensities, corrected for sensitivity factors of the individual species. The distribution varied only slightly over the range of pressures used in this work, however it changes significantly at higher pressures. In this paper, the dose of methyl is reported as Langmuirs (L) of exposure to the total product gas pressure. Surface temperature was held at 120 K during dosing. At this temperature, methane, ethane and nitrogen are too weakly bound to adsorb on the surface. Control exper~maents confLrmed that pressures of methane, ethane and nitrogen 10 times higher than produced in the pyrolysis source do not interfere with methyl adsorption or reaction. Hydrogen, on the other hand, does adsorb during dosing of the methyl radicals. It is possible that the hydrogen concentration is comparable to the methyl concentration in experiments on the Ni(100) surface because the sticking probabilities are comparable. In experiments on the Ni(111) surface, it is

329 likely that the hydrogen concentration is less than 10% of the methyl concentration because the sticking probability of hydrogen on this surface is only around 0.01. In experiments on the surfaces with chemisorbed oxygen or multilayer oxide films, hydrogen does not adsorb. Calibration of the amount of desorbing hydrogen was carried out by comparing the hydrogen TPD peak area with that measured for a saturation dose of hydrogen, which is known to give 1 H atom per surface Ni atom.[19] This is defined as 1 monolayer (ML) coverage and corresponds a surface density of 1.9 x 1015 atoms/era2 for N i ( l l l ) and 1.6 x 1015 atoms/era2 for Ni(100). The coverages of all species are referenced to these values. Calibration of the amount of desorbing CH4 was carried out by comparing the methane TPD peak area to that measured for a saturation exposure of CO (.67 ML) scaled by the ratio of the mass spectrometer sensitivities, and UHV-chamber residence times for these two gases. Calibration of the amount of residual carbon was carried out by comparing the (2(272 eV)/Ni(848 eV) AES peak to peak ratio with the one measured after decomposition of a saturation amount of ethylene on Ni(111), which gives a surface carbon coverage of I/4 monolayer.[20] Self-consistent mass balances were achieved using these calibrations in experiments in which CH3 and CO were coadsorbed in varying amounts.

3. Results and Discussion3.1 CH3 on Ni(100) A representative temperature programmed desorption (TPD) specmma of products resulting from reactions of methyl on Ni(100) is shown in Fig. 1. The only gas-phase products detected were methane (16 and 15 ainu curves) and hydrogen (2 ainu curve). No C2 or higher hydrocarbons were observed. A portion of the adsorbed methyl groups decomposed on the surface by reactions R1 through R3, CH3 CH2 CH

--> CH2 + H(S) --> CH + H(S) --> C(S) + H(S)

(R1) (R2) (R3)

where (S) refers to surface-bound species. (Because only surface-bound hydrocarbon fragment groups are discussed here, this designation for surface species is omitted). The m o u n t of CH3 decomposed is determined by measuring the amount of residual surface carbon by Auger spectroscopy. The peak in the methane TPD curve occurs at 230 K. The appearance of gas-phase methane marks the temperature at which methane is formed on the surface because the reaction temperature is well above the desorption temperature for methane. As will be discussed below, the dominant mechanism for methane formation on Ni(100) is: CH3 + H(S) --> CI-I4(g)

(R4)

Hence, the appearance of gas-phase methane marks the occurrence of a C-H bond-breaking step. In the hydrogen TPD ~ m m a , the lower-temperature peak occurs at 355 K, characteristic of desorption of surface hydrogen from Ni(100). Hence it is defined by the desorption kinetics and does not provide information on C-H bond-breaking reactions that might occur at lower temperatures. The higher-temperature desorption peak at 390K on the other hand occurs above the temperature for hydrogen desorption and is believed to be related to decomposition of surface the most stable fragment, CH, or a dimer thereof to carbon and hydrogen. The spectra shown in fig. 1 are for a methyl surface coverage of about 0.1 ML. The results obtained at this coverage are in excellent agreement with earlier experiments involving decomposition of methyl iodide to produce surface-bound methyl groups.[ 1, 3, 21] The methyl coverage as a function of dose is shown in fig. 2. The methyl coverage is

330 1 E 3

v

~. 0.5 ec-

15 AMU _(CH_4)

m

0

i

i

i

,

,

i

i

i

100

i

I

t

.

.

.

.

.

.

300

.

.

i

.

.

.

.

.

.

500

.

.

.

i

i

.

i

.

700

Temperature, K Figure 1 Mass spectrometer (MS) intensities versus surface temperature in a representative temperature programmed desorption (TPD) profile of methyl on a clean Ni(100) surface. Total exposure was 1 L . Adsorption temperature was 105 K and heating rate was 3 K/s. Methane is produced near 225 K. Hydrogen desorbs in a wide temperature range, 260 - 420 K. taken to be the sum of the methane observed in TPD and the residual surface carbon measured after annealing to 600 K. The maximum methyl coverage is 0.42 ML (+ 10%). This is about 2 times higher than can be achieved via the decomposition of methyl iodide [ 1]. This qualitatively seems reasonable since methyl groups and I atoms have roughly the same diameters on the surface, and there is necessarily 1 iodine atom per methyl group. The uptake of methyl roughly follows langmuirian adsorption kinetics with a sticking probability per collision with the clean metal surface of about 0.1. The sticking probability determined from these experiments has an uncertainty of about a factor of 2, due primarily in the uncertainty in determining the absolute partial pressure of methyl groups coming from the pyrolysis source.

0.5'

I

I

I

I

~'~ 0.4 r

I

I

"="

~'

0.3 0>

ro -r-

0.2 0.1 0.0 0

2

4

I

I

I

I

6

8

10

12

C H3 Exposure (L)

--

Figure 2 Methyl coverage in monolayers (ML) as a function of methyl exposure, in langmuirs (L) of mixture of gases from pyrolysis source. Symbols are data showing the sum of methane formed plus residual carbon. Consistent values are obtained by summing hydrogen appearing in CH4 and H2 gas-phase products. Solid line is a guide to the eye.

331 The dependence of methane formation on methyl coverage is shown in fig. 3. The yield of methane increases as the coverage increases. At coverages below 0.05 ML, little methane is formed. Nearly all of the methyl groups decompose. At saturation coverage, about 0.14 ML of the methyl groups form methane, and roughly 0.27 ML decompose An important trend to note in fig.3 is that the peak in the TPD spectrum shifts to higher temperatures as the coverage increases. If the rate determining step is the rupture of the CH2-H bond, and this step is f'trst order, the peak in the TPD curves should occur at the same temperature. The shift to higher temperatures is at least in part due to the fact that higher temperatures are required for the CH2-H bond breaking reaction as the surface becomes more crowded. We find that preadsorbing carbon in various amounts up to 0.3 ML leads to a progressive shift to higher temperatures. The temperature shift observed with 0.3 ML of preadsorbed carbon is similar to that exhibited by curve h in fig. 3. The effects of st~ace crowding and co-adsorbed species will be reported in more detail elsewhere.

CH 4 / OH 3 / Ni(100) "S"

g :E

1

. . . . . . . . . . . . . . . . . . . . . . . . . .

1O0

i

.......................

|

200 300 Temperature, K

....

.........

400

Figure 3 CI-h TPD from Ni(100) at various initial coverages of CH3. Curve a=0.025; b-----0.06; c--0.10; d=0.18; e---0.33; f=0.37; g=0.40 ML.

3.2 CH3 on N i ( l l l ) Methane formation from the reactions of methyl groups on a Ni(111) surface at various methyl coverages is shown in fig. 4. The saturation coverage we observe for methyl on Ni(111) is 0.32 M'J., (+ 10%). Within experimental error, this is the same surface density of methyl groups as the saturation coverage found for Ni(100) (because there are more surface Ni atoms per ML on Ni(111)). As with Ni(100), we fred that methane yield increases faster than linearly in methyl coverage. This is illustrated in fig. 5. In the lower panel, the yield (total amount formed) is plotted, and in the upper panel, the selectivity (fraction of adsorbed methyls going to methane) is shown. Again we see that below about 0.05 ML that little or no methane is formed. This is consistent with experiments of Ceyer and coworkers who observe little methane formation from low coverages of methyl groups on Ni(111), even in the presence of considerable co-adsorbed hydrogen.[7] At coverages above 0.1 ML, about 15% of the methyl groups desorb as methane, increasing to about 30% at coverages of 0.3 ML.

332

_>, i-

1.0 A

E

-r0

0

0.5

o,p 2" 0.0 u ~ - ~

0.08

0.0

2

-

I

!

-

-

-I-

=~

~ 0.04 200

300

400

Temperature (K) Figure 4 CH4 TPD from Ni(111) at various initial coverages of methyl. Curve a--0; b=0.03; c=0.08; d--0.11; e=0.13; f=0.24; g=0.26; h--0.29; i----0.31 ML.

0.00 m -=r" 0.0

J

J

I

0.1

0.2

0.3

Initial Methyl Coverage (ML) Figure 5 CH4 yield (bottom plot) and selectivity (top plot) versus initial CH3 coverage on N i ( l l l ) . Yield is the absolute amount of CI-I4 formed, selectivity is the fraction of adsorbed CH3 that goes to CI-I4

The shape of the TPD spectra are different than they were for Ni(100), and the peak occurs at slightly higher temperatures, about 245 K. Also, in contrast to the trend observed on Ni(100), on Ni(111) the peak in the methane TPD curve shifts to lower temperatures as the surface coverage is increased. This trend is frequently associated with a reaction step that is higher than first order in reactant coverage. It is an important indication that the dominant reaction pathway is probably different than it is on the (100) surface. Further evidence for different reaction pathways is obtained in isotope labelling experiments, illustrated in fig.6. Here we present a comparison of the effects of preadsorbing various amounts of deuterium with methyl groups on the two metal surfaces. For a reaction network consisting of reactions R1 through R4, the following trend would be expected. At low deuterium coverages, most of the hydrogen required for methane formation must come from decomposition of some fraction of the methyl groups. As the deuterium coverages increases, it will compete with increasing effectiveness for undecomposed methyl groups. The balance between these two pathways can be determined from the relative amounts of CH4 and CH3]) in the TPD spectra. (No methane with more than 1 deuterium atom is observed, even at the highest deuterium coverages. This is clear evidence that at least reaction R1 is irreversible. It also suggests that reactions R2 and R3 are irreversible. Further evidence for this is presented below.) The results obtained for Ni(100) are consistent with this trend. As the deuterium coverage increases from 0.1 ML to 0.7 NIL, CH3D yields increase, and CI-I4 yields decrease, reflecting an increasing contribution from reaction of methyl with surface deuterium, and decreasing contribution from reaction with hydrogen supplied from other methyl groups on the surface. Note especially that the TPD curves shift in opposite directions for each product. As the deuterium coverages increases, the peak in the CH3D spectrum shifts to lower temperature and the peak in the CH4 spectrum shifts to higher temperature. This is the qualitative behavior

333

~ II

04

_>, "~~

_>, "~~

I

I

I

I

)D(ML)

r~=-----.-.~.-,~--~-~- ~ - - . L - . ~ 0 . 1 5 -

=

,.

o.o o.o

150

200

250

300

350

Temperature (K)

400

Figure 6 CI-I~D (dotted curve) and ~

150

200

250

300

350

Temperature (K)

_

400

(solid curve) formed on Ni(]00) (left hand

figure) and Ni(111) (fight hand figure) at a fixed CHa initial coverage of 0.2 ML, and various amounts of preadsorbed deuterium (indicated in the figures, units are monolayers of deuterium atoms) expected for reactions R 1 through R4. In fact, a simple kinetic model that includes only these reactions can quantitatively reproduce the dependence on deuterium coverage of both the relative yields of CHaD and CH4 and the peak desorption temperatures. The details of this model will be published separately. The results obtained for a Ni(111) surface are significantly different. Most striking is that even with a 3 fold excess of deuterium relative to methyl, very little CHaD is produced. Reaction R4 appears to be a minor channel on Ni(111). Furthermore, because the deuterium coverage is nearly invariant over the entire temperature range over which methane is produced, it is unlikely that the CH4 that is formed comes methyl reaction with surface hydrogen generated by methyl decomposition. One would expect more CHaD than CI-I4 because the D coverage is at all times equal to or greater than the H coverage. The possibility that the relatively slow reaction of CHa with D(S) might be due to an unusually large deuterium kinetic isotope effect was tested by performing experiments with preadsorbed H(S). We observe no significant increase in the amount of methane formed in the presence of preadsorbed H atoms on Ni(111), indicating that D(S) and H(S) have similar reaction rates. CH4 must therefore come from a disproportionation or other reaction of two hydrocarbon fragments, and not by a reaction that involves surface-bound hydrogen atoms. Similar experiments have been done with 0.25 ML of deuterium and varying amounts of methyl. The CH4 yield as a function of methyl coverage at fixed D coverage is very similar to that found in the absence of coadsorbed deuterium. This is not surprising in that the results illustrated in fig. 6 show that the main peak in the methane TPD spectrum is not significantly affected by coadsorbed deuterium. Results obtained with 0.25 ML of coadsorbed deuterium are very similar to those shown in figs. 4 and 5. The CH4 yield has a higher than first order dependence on coverage. This cannot be attributed to the coverage dependence of reaction R4, since it makes only a minor contribution. It is more likely to be due to a bimolecular reaction

334 between two hydrocarbon fragments, which would exhibit a greater than f'wst order dependence on coverage. The mechanism for a non-surface-mediated hydrogen-tranfer from one hydrocarbon fragment to another cannot be identified in ambiguously from the current data alone. Possible reaction pathways include: CH3 + CH3 --> [C2H6] --> CH4(g) + CH2

(disproportionation)

R5

where, [C2H6]is a short-lived complex involving one or more bonds to the metal surface, and (g) designates a gas-phase product; CH3 + CH3 --> CI-h(g) + CH2

(abstraction)

R6

i.e., a prompt reaction in which an H atom is stripped from a surface bound CH3 group by one that is not; CH2 + CH3 --> [-CH2CH3] --> CH4(g) + CH

(methylene insertion)

R7

(methyne insertion)

R8

where [-CH2CH3]is a surface bound ethyl group; CH + CH3 --> [=CHCH3] --> CH4(g) + C and, CH2 + CH2---> [C2H4] --> CH4(g) + C

R9

We discount the likelihood of reaction R5 because two metal-bonded methyl groups have no molecular orbitals readily available for reaction. These groups are sp3 hybridized with all 4 orbitals fully occupied. The lowest lying unoccupied orbital is so much higher in energy that it is inaccessible, and there are no obvious interactions that might enable a rehybridization to lower this energy. An abstraction mechanism, R6, might proceed on an available orbital basis, but seems unlikely based on the following. A free, or nearly free, methyl group desorbing from the surface has an orbital available for forming a fourth C-H bond. The abstraction of a hydrogen from CHa(g) by CH3(g) is well known; the heat of reaction is zero, and the energy barrier is only about 50 kJhnol. Since the C-H bond strength of CH3 is less than that in CH4 by about 50 kJ/mol, this energy barrier could be relatively low. However, it is also necessary to break, or nearly break, the metal-methyl bond to provide the free methyl group. The chemisorption energy of methyl on Ni(111) has been estimated to be 160 kJ/mol.[22] This would make the overall reaction endothermic by about this amount, much too high to expect it to occur at 250 IC Furthermore, if the metal-methyl bond were weak enough for this process to occur, we would expect to see at least some free methyl groups in the TPD specmun. We were unable to detect any desorption of free methyl groups in these experiments, although it has been observed from Cu [6] and NiO [23] surfaces. The CH2 and CH groups generated by reactions R1 or R2 do have orbitals readily available for reaction, and the energy barriers are likely to be quite low. We postulate reactions R7 through R9 based on analogy with known organometallic or metal surface chemistry. The formation of ethyl groups from CH3 and CH2 fragments (the first step in reaction R7) has been observed on Cu surfaces in experiments in which the fragments were created by the dissociation of coadsorbed CH3I and CH212 respectively.[5] However, the dominant reaction pathway of ethyl groups is to undergo a 13-hydrogen elimination to give ethylene, not methane.[5] We would expect to detect at least some formation of ethylene in conjunction with reaction R7.

335 Since we do not observe any evidence for ethylene formation, we rule out reaction R7. Similarly, we rule out reaction R9, which should also produce at least some ethylene. We believe reaction R8 is the most likely route by which the methane is formed on Ni(111). CH has been observed from the decomposition of methyl at low coverages on Ni(111) [9, 24], and, in the absence of methyl groups, CH can dimerize to give C2H2.[9] In the presence of methyl groups, other C-C bond forming reactions should be expected. Support for this comes from the fact that ethylidyne (CCH3) has been observed in reactions of CH3 on P t ( l l l ) under conditions where both CH and CH3 are likely to present.[12] Ethylidyne formation is unusually favorable on Pt(111), and it is reasonable to expect that related C2 adducts formed on Ni(111) might rearrange to give methane according to reaction R8, rather than ethylidyne. Lastly, there is some evidence in the shape of the TPD curves shown in fig. 6 for the formation of a short-lived reaction intermediate that is a precursor to the methane formation on Ni(111). Note that the CH3D curve drops sharply at the same time (temperanwe) that CH4 production is at a maximum. A reaction network consisting of reactions R1 through R4 is not able to reproduce this behavior. We can envision only two ways in which CH3D production might decline while the CH4 formation rate is increasing. The fast is that there are two different kinds of CH3 on the surface, a minority of CH3's (e.g. those at a special site) that get used up in reacting with surface D, but cannot form CH4, and the rest of the CH3's that can form CH4. The second is that all of the CH3 groups react to form a C2 complex from which CH4 is formed, leaving no CH3 groups to react with surface D. We believe that the latter is a more reasonable explanation, and that reaction R8 is the pathway by which this is occurring. A more detailed description of the kinetic modeling of the reactions of methyl on Ni(111) and (100) will be presented in a subsequent paper. 4.0 Conclusions: We have characterized the reaction pathways of CH3 groups as a function of coverage on Ni(100)andNi(lll)surfaces. Relatively high coverages, up to 0.4 ML, have been investigated for the first time. On both surfaces, methane and hydrogen are the only gas-phase products detected. Methyl C-H bonds begin breaking at around 220 K, and methyl decomposition occurs in parallel with methane formation over a limited temperature range. At high coverage, up to 1/3 of the methyl groups form methane, and the remainder decompose, ultimately producing surface carbon and hydrogen gas. The dependence of the methane yield on methyl coverage is similar on the two surfaces, but the mechanism by which methane formation occurs is different. On Ni(100), the dominant mechanism is the reaction of CH3 and surface hydrogen. On Ni(111), the reaction of CH3 with surface hydrogen is relatively slow, and the dominant mechanism involves a hydrogen transfer between two hydrocarbon fragment groups. Experimental results suggest that CI-h is formed from a C2 reaction intermediate (CHCH3) from the reaction of CH with CH3. References: 1. F. Zaera, Ace. Chem. Res, 25 (1992): 260. 2. S. Tjandra and F. Zaera, J. Catal., 157 (1994): 598. 3. S. Tjandra and F. Zaera, Langmuir, 8 (1992): 2090. 4. C.-M. Chiang, T.H. Wentzlaff, C.J. Jenks, and B.E. Bent, J.Vac. Sci. Technol., A10 (1992): 2185. 5. C.-M. Chiang, T.H. Wentzlaff, and B.E. Bent, J. Phys. Chem., 96 (1992): 1836. 6. J.-L. Lin and B.E. Bent, J. Vac. Sci. Technol., A10 (1992): 2202. 7. A.D. Johnson, S.P. Daley, A.L. Utz, and S.T. Ceyer, Science, 257 (1992): 223. 8. M.B. Lee, Q.Y. Yang, S.L. Tang, and S.T. Ceyer, J. Chem. Phys, 85 (1986): 1693.

336 9. Q.Y. Yang, K.J. Maynard, A~D. Johnson, and S.T. Ceyer, J. Chem. Phys., 102, no. 19 (1995): 7734. 10. X.D. Peng, R. Viswanathan, G.H.J. Smudde, and P.C. Stair, Rev. Sci. Instrum., 63 (1992): 3930. 11. G.H.J. Smudde, X.D. Peng, R. Viswanathan, and P.C. Stair, J. Vac. Sci. Technol., A9 (1991): 1885. 12. D.H. Fairbrother, X.D. Peng, R. Viswanathan, P.C. Stair, M. Trenary, and J. Fan, Surf. Sci, 285 (1993): LA55. 13. G.H. Smudde, Jr., X.D. Peng, R. Viswanathan, and P.C. Stair, J. Am. Chem. So<:., 115, no. 5 (1993): 1988. 14. B.B. Pearce, M.V. Twigg, and C. Woodward, "'Methanation ", in "Catalyst Handbook ", M.V. Twigg, Editor. 1989, Wolfe Publishing Ltd: London England. 15. D.E. Ridler and M.V. Twigg, Steam Reforming, in Catalyst Handbook, M.V. Twigg, Editor. 1989, Wolfe Publishing Ltd: London England. 16. M.I. Temkin, Industrial Heterogeneous Catalytic Reactions. Advances in Catalysis, eeL D.D. Eley, H. Pines, and P.B. Weisz. 1979: Academic Press. 17. R.B. Hall, C.A. Mires, J.H. Hardenbergh, and J.G. Chen, Structure and Reactivity of Nickel Monoxide Surface Films on Ni(lO0), in Surface Science of Catalysis, D. Dwyer and F. Hoffinarm, Editors. 1991, ACS Publishing: Washington DC. 18. R. Renaud and L.C. Leitch, Can. J. Chem., 32 (1954): 545. 19. K.D. Rendulic and A. Winkler, J. Chem. Phys., 79, no. 10 (1983): 5151. 20. B.E. Koel, J.M. White, and D.W. Goodman, Chem. Phys. Lea., 88 (1982): 236. 21. X.-L. Zhou and J.M. White, Surf. Sci., 194 (1988): 438. 22. H. Yang, J.L. Whitten, R.E. Thomas, R.A. Rudder, and R.J. Markunas, Surf. Sci., 277 (1992): L95. 23. C.M. Kim, J.G. Chen, C.A. Mires, and R.B. Hall, to be published, (1995). 24. J.G. Chen, M. Castro, C.A. Minas, and R.B. Hall, to be published, (1995).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 1 lth International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

337

Normalization by Oxygen Uptake of the Rates of Oxidative Dehydrogenation of Methanol and Ethanol S. Tanabe + , H. E. Davis, Jr. x , D. Wei and R. S. Weber* Department of Chemical Engineering, Yale University New Haven, Connecticut 06520-8286, USA

Abstract Silica-supported, V-substituted polyoxometalates containing V1, V3 and V 1 4 ensembles were characterized by the oxygen uptake protocol developed by Oyama and co-workers. Rates of methanol and ethanol oxidation catalyzed by these samples scaled well with the number of exposed vanadiums and almost as well with the number of oxygens taken up. The results confirm the weak dependence on the local concentration of redox centers of the catalytic oxidative dehydrogenations. A slight, residual variation in turnover frequencies when normalized by oxygen uptake is interpreted as an adjacency dependence in the number of oxygen vacancies created by the reduction step or, equivalently, as an adjacency dependence in the average degree of reduction of the samples. 1. I N T R O D U C T I O N Oxygen uptake has long been employed as a method for counting and differentiating among redox sites present on the surfaces of pre-reduced catalyst samples [ 1, 2]. While there have been some differences of opinion regarding the most useful way of carrying out the measurements [3-6], the appealing simplicity of the method recommends it over alternate, more complicated protocols that claim to count specific types of oxygen uptake sites [7-9]. Recently, Oyama and coworkers [ 10, 11 ] have refined the conditions required to achieve consistent, plausible site-counting in catalysts containing transition metals (M) employed for the oxidative dehydration of alcohols and other oxidation reactions. To probe further the significance of the O/M ratios obtained using Oyama's protocol, we have applied it to samples containing silica-supported, V-substituted Keggin ions in which the number of adjacent V ions was varied (Table 1). The parent tungsten compound, Na3[PW12040], is not reducible under the conditions described below and has no activity as an oxidation catalyst [ 12]. Therefore, these cluster-derived samples offer more control over the structures of the surface redox species than does a straightforward variation in the areal loading of the vanadium.

+ Department of Chemistry, Nagasaki University, 1-14 Bunkyo-machi, Nagasaki, 852, Japan x Department of Chemical Engineering, Massachusetts Institute of Technology, Cambridge, MA 02138, USA * To whom correspondence should be addressed.

338 Table 1 Polyoxometalate clusters employed as catalysts. Poly0xometalate Vanadium moieties Na4[PVW 11O40]/SIO2

no adjacent V (redox) neighbors

Na6[PV3W9040]/SiO2

3 vIO~v

Na9[PV14042]/SiO2

30 V - - - " O ~ v dimers, 12

dimers, 1

? " ~ V I / O trimer VI,-,-~ V O - . ~ V.I/O trimers, V...- ,-,~ V

plus higher order groupings

2. Experimental Preparation of the samples has been described in detail elsewhere [ 12]. Briefly, a high area, fumed silica (Cabosil HS-5) was impregnated with aqueous solutions of the sodium salts of the polyoxometalates to achieve submonolayer coverage of the silica (about 15wt% and about 15 area%). The samples were dried, pressed into wafers (10MPa) and then crushed and sieved into 40-60 mesh particles. Aliquots containing 5-24 gmol V were reduced in a tubular Pyrex reactor at 575, 600 or 625 K in flowing methanol vapor (5% methanol, 95% He) for 15 minutes. The samples were purged with He for 30 minutes and then exposed to a uniform sequence of 1.6 I.tmol pulses of 02 at 575, 600 or 625 K until no further oxygen uptake could be detected (usually 10 pulses sufficed). Each series of experiments was repeated at least once with a different ordering of the reduction/oxidation temperatures to check for hysteresis. The oxygen taken up by the samples was measured with the aid of a thermal conductivity cell in a gas chromatograph (HP-5790). The reduction temperature was limited to 625 K to preclude decomposition of the heteropolyanions, which is known to occur at 640-650 K [ 12-15]. The same samples, after a pretreatment in flowing oxygen (10%) at 625 K, were used as catalysts for the oxidative dehydrogenation of ethanol and methanol in the same reactor. The reaction mixture consisted of 02 (3 or 5%), methanol vapor (3%) or ethanol vapor (5%) and He (balance), all delivered by Tylan mass flow controllers or vaporizer flow controllers. Products were analyzed by gas chromatography. The catalysts exhibited no induction period and their activities were stable over many days and over repeated temperature cycles.

3. Results The oxygen uptake results were quite reproducible and they were not much affected by the history of the samples (+10% repeatability). Longer reduction times and longer purge times also did not change the results. Total amounts of oxygen taken up by these supported clusters depended strongly on the temperature of the reduction and oxidation steps (Figure 1). The results resembled those obtained by Oyama and Somorjai for samples of bulk V205 [11]. In all cases the rate of oxygen uptake (as sampled by the pulse sequence) fit a rate expression that was first-order with respect to the concentration of redox sites (Figure 2).

339

I

1.0

I

I

I

u

1.0

0.8

0.8 0

v

+

"." 0.6

m

t~ .,.., t~ e"

a~ 0.4

m

~ o.6 ~0.4

X

o ~0.2

0.2 0.o

$

I 580

I, 600

I 620

I 640

Temperature/K

Figure I. Oxygen uptake by supported clusters and vanadium oxide. Samples were pre-reduced and re-oxidized at the temperatures indicated on the abscissa. Silica-supported polyoxometalates: PVI (O), PV3 (A), PV14 (~). Bulk V205 (+, after [ 10])

0.0

0

2

4

6

8

10

Pulse Number Figure 2. Estimated rates of oxygen uptake by samples of silica-supported PV1 (circles), PV3 (triangles) and PV14 (diamonds). Approach to saturation was modeled as either first-order kinetics (solid line, X2=0.054) or second-order kinetics (dashed line, X2 =0.585).

Stoichiometries of the total oxygen uptake, O/V, measured at the highest reduction temperature and highest reoxidation temperature, varied between 0.24-0.40 (Table 2), similar to the value found by Oyama and Somorjai for bulk V205. Apparent activation energies for the catalytic reactions were as expected: about 80 kJ/mol for the formation of formaldehyde and 60 kJ/mol for the formation of acetaldehyde from the respective alcohols (Figure 3). The turnover rates of the samples were calculated either on the basis of the number of vanadiums (all of which could be assumed to be accessible) or by assuming that oxygen uptake counted the catalytic sites:

Nv =

mol product mol V x second

NO =

mol product mol O x second

The oxygen-based turnover rates were calculated in one of two ways. In the first method, the number of redox sites was set equal to the number of oxygen atoms taken up at 625 K by the samples (Table 3, Figure 3). This temperature corresponds closely to that recommended by Oyama and Somorjai [ 11 ] as being neither too high nor too low. In the second method, the number of redox sites was set equal to the number of oxygen atoms taken up when the temperature of the pre-reduction and re-oxidation steps of the titration equaled the temperature of the alcohol oxidation (Figure 4). The latter method gives rise to unusual Arrhenius curves for the reaction rates since it combines the temperature dependence of the titration with that of the catalytic reaction.

340 Table 2 Oxygen uptake stoichiometries of pre-reduced supported polyoxometalates Oxygen uptake (Treduction = Treoxidation = .... lamol O/~.mol V PV1 0.40 PV3 0.30 PV14 0.25 Bulk V205 a 0.50 a) Extrapolated to 625 K from data presented in reference [10]. Sample

625 K) e-/V 0.8 0.6 0.5 1.0

Table 3 Catalytic activities of the silica-supported vanadium oxides in either 3% 02 and 3% methanol or in 5% 02 and 3% ethanol. Oxygen uptake was measured at 625 K. Methanol oxidation at 625 K Ethanol oxidation at 575 K Sample

Nv/mHz a

Nv/mHz

No/mHz

Nv/mHz

PV1 2.9 2.0 5.0 PV3 3.7 2.2 8.8 PV14 4.0 2.4 10.0 1.4% V205 b . . . . 3.5% V205 b . . . . 5.6% V205 b . . . . 7.7% V205 b . . . . 9.8% V205 b . . . . a) From reference [12]. b) Silica-supported samples" extrapolated to the same reaction conditions reference [ 11 ].

0.97 1.83 1.95

No/mHz 2.5 6.1 5.8 1.9 4.1 8.6 7.7 7.5

from data presented in

Vanadium-based turnover rates for the oxidation of methanol matched those obtained previously using these same samples but a different reactor [ 12]. The oxygen-based turnover rates agreed very well with those obtained by Oyama and Somorjai. 4. DISCUSSION To more closely follow the protocol described by Oyama and Somorjai, we should have reduced our samples in flowing hydrogen. Indeed, polyoxometalates can be reduced with H2 but only with the assistance of catalytic quantities of Pt or Pd acting as a source of spilledover hydrogen [16]. We avoided that complication here by using a gas (methanol vapor) which is known to be oxidized by the catalysts and therefore is guaranteed to reduce the heteropolyanions at reaction temperatures below the decomposition temperature of the clusters. We also attempted to reduce the catalysts with He-diluted ethanol vapor but found that it produced large quantities of coke on the more acidic samples (PV 1, PV3). Upon reduction with methanol at 625 K, the vanadium ions in these samples exhibited an apparent, average degree of reduction of 0.5--0.8 electrons per atom (Table 3). This value

341

-3

-3 -4

Formaldehyde

.,.,.

-4 _

Acetaldehyde

~-5 8

~-6 "E

-

"7 I

O

O

'~

-8 1.60

I I I I 1 . 6 5 1 . 7 0 1 . 7 5 1.80 kK/T

I 1.85

-8 1.60

1.90

I 1.65

I 1.70

I I 1 . 7 5 1 . 8 0 1 . 8 5 1.90 kK/T

Figure 3. Arrhenius plots for the formation of formaldehyde or acetaldehyde from methanol or ethanol, normalized by the number of vanadiums (open symbols) and by the amount of oxygen uptake measured at 625 K (filled symbols). Lines on the right panel are calculated from the data reported by Oyama and Somorjai [ 11].

-3

-3 -4

Formaldehyde

m

9

~-5

E i..

I~ii II

9 $

Acetaldehyde

-4

:

R A

-5 8

O

--

~-6

~-6

O

-

o

O

r-

-7--8 1.60

-7---

A I I I 1 . 6 5 1 . 7 0 1.75 kK/T

I 1.80

I 1 . 8 5 1.90

-8 1.60

I 1.65

I 1.70

i I I 1 . 7 5 1 . 8 0 1 . 8 5 1.90 kK/T

Figure 4. Arrhenius plots for the formation of formaldehyde or acetaldehyde from methanol or ethanol, normalized by the number of vanadiums (open symbols) and by the amount of oxygen uptake measured at the temperature of reaction (filled symbols).

agrees well with that extrapolated to this temperature from the experiments by Oyama and co-workers. The larger ensembles had the lower O/V stoichiometry. We infer that a higher degree of adjacency of the vanadium ions in the clusters correlates with a small decrease in the amount of oxygen taken up per vanadium or, equivalently, with a small decrease in the extent of reduction for the same treatment.

342 Since the areal loading of the supported heteropolyanions is constant, the surface concentration of the vanadium in the samples varies by a factor of 14 from PV 1 to PV 14. Table 1 shows that the surface concentrations of higher ensembles (dimers, trimers) varies even more strongly among the samples. Yet, all of the samples exhibit an identical pattern of re-oxidation that is well represented by first order kinetics (Figure 2). For comparison, that figure also presents a curve corresponding to second-order kinetics. Evidently, the fit of such a curve to the data would be considerably worse than that shown if it accounted for the variation in the initial concentrations of the V ensembles. We propose, therefore, that the adjacency of the vanadium ions does not affect the rate at which they are each re-oxidized. To be consistent, we must then also hypothesize the existence of partially reduced oxygen species which are free to roam, gathering (or exchanging) electrons before finally resting in a vacant anion site. There is clear spectroscopic evidence for the reverse of this process during reduction of heteropolyanions [16-18], and, in general, for the existence of the partially reduced oxygen species on the surfaces of redox catalysts (e.g., [19]). It is known that the re-oxidation step is not so simple at lower temperatures. For example, Neumann and Levin [20] have found that re-oxidation is the rate limiting step when heteropolyanions act as homogeneous catalysts and that the reaction rates depend on the square of the catalyst concentration. It is reasonable to infer that temperatures comparable to those employed here are necessary for the rate of the re-oxidation step to be independent of the adjacency of the transition metal cations. For both methanol and ethanol, normalization of the rates of oxidative dehydrogenation by the number of vanadium ions present in the catalysts serves to group the Arrhenius lines together (Figure 3, open symbols). There is scarcely a factor of 2 variation among the Nv's at a given temperature, confirming the very weak dependence of the rate of either oxidative reaction on the adjacency of redox sites. Normalization by the amount of oxygen uptake also accounts for much of the variation in specific activity among the samples for both reactions. Thus, our results corroborate those of Oyama and Somorjai for ethanol oxidation and extend them to methanol oxidation. However, as can be seen from the superposition of data sets in Figure 3, there is a small residual variation in the rates normalized by oxygen uptake which closely resembles that found by Oyama and Somorjai [11] for the oxidation of ethanol. The variation in rates is about a factor of 3 between samples containing isolated vanadium ions (PV1, low loading samples) and samples containing higher oligomers (PV3, PVI4, high loading samples). Given the range of O/V ratios exhibited by our samples (Table 3), we suggest that this mild "structure sensitivity" arises from variations in the titration reaction with catalyst structure: the isolated V ions are, on average, slightly more oxidized (or slightly more reduced for a given oxidation stoichiometry) than are the larger ensembles. We speculate that this variation in degree of reduction persists under the conditions of steady state reaction. If so, then it may bear upon the results of Louis et al. [21 ] and of those from Iwasawa' s laboratory [22], all of whom have .attributed higher turnover rates to clusters of redox sites.

5. Summary Neither the oxidation of methanol to make formaldehyde nor that of ethanol to make acetaldehyde is sensitive to the adjacency of redox sites in the catalysts: the number of exposed vanadium serves well to normalize the reaction rate in both cases. This conclusion does not contradict a statement by other authors [23-25] that the selectivity of alcohol

343 oxidation is "structure sensitive." The selectivity of a catalyst is a ratio of rates of reactions. Selectivity alone, therefore, provides ambiguous information since it is not clear whether it is the numerator or the denominator (or both) that varies with the molecular and electronic structure of the surface of the catalyst. In the case of methanol oxidation catalyzed by supported polyoxometalates, we have found previously that changes in the observed product selectivity (aldehyde/ether) can be attributed just to changes in the turnover rate of production of the ether that correlate with the anionic charge of the supported clusters [ 12]. We expect that a similar pattern will hold for the oxidation of other alcohols provided that the extent of reduction of the catalyst can be controlled. The rates of oxidative dehydrogenation of methanol and ethanol are also normalized reasonably well by the amount of oxygen taken up by the catalysts. Thus, we urge the use of oxygen uptake as a convenient, easily implemented means of characterizing the surfaces of oxidation catalysts, particularly in those cases where it is difficult to count the number of accessible redox sites by other means. At the least, oxygen uptake can serve as a benchmark measurement to identify instances of unusual behavior. It will be especially interesting to learn whether the small variation (factor of 3) found by us and by Oyama and co-workers between NO measured on isolated and that measured for agglomerated sites applies to other catalysts and other reactions. Further resolution of the details of oxidative dehydrogenation requires the measurement of a catalyst's degree of reduction carried out during steady state reaction. We note that UVvisible spectroscopy offers a way to perform this measurement since many of the transition metal oxides which are active as oxidation catalysts exhibit striking color changes between their oxidized and reduced states. ACKNOWLEDGEMENTS Portions of this work were supported by the Exxon Educational Foundation, the State of Connecticut, the US National Aeronautic and Space Administration (HED), the US Department of Energy (DW) and the Ministry of Education, Science and Culture of Japan

(ST). REFERENCES 1. S. Brunauer, and P.H. Emmett, J. Am. Chem. Soc., 57, (1935) 1754. 2. S. Brunauer, and P.H. Emmett, J. Am. Chem. Soc., 62, (1940) 1732. 3. R. Grabowski, B. Grzybowska, J. Haber, and J. Sloczynski, React. Kinet. Catal. Lett., 2, (1975) 81. 4. B.S. Parekh, and S.W. Weller, J. Catal., 47, (1977) I00. 5. J. Valyon, and W.K. Hall, J. Catal., 84, (1983) 216. 6. N.K. Nag, K.V.R. Chary, B.M. Reddy, B. Rama Rao, and V.S. Subrahmanyam, Appl. Catal., 9, (1984) 225. 7. M. Niwa, Y. Matsuoka, and Y. Murakami, J. Phys. Chem., 91, (I 987) 4519. 8. M. Niwa, Y. Matsuoka, and Y. Murakami, J. Phys. Chem., 93, (1989) 3660. 9. M. Niwa, H. Yamada, and Y. Murakami, J. CataI., 134, (1992) 331.

344

10. S.T. Oyama, G.T. Went, K.B. Lewis, A.T. Bell, and G.A. Somorjai, J. Phys. Chem., 93, (1989) 6786. 11. S.T. Oyama, and G.A. Somorjai, J. Phys. Chem., 94, (1990) 5022. 12. C.M. Sorensen, and R.S. Weber, J. Catal., 142, (1993) 1. 13. M. Bolmer, Ph.D. Dissertation, University of Delaware (1981 ). 14. S. Kasztelan, E. Payen, and J.B. Moffat, J. Catal., 125, (1990) 45. 15. M.J. Bartoli, L. Monceaux, E. Bordes, G. Hecquet, and P. Courtine, in (P. Ruiz, B. Delmon, Eds.), New Developments in Selective Oxidation by Heterogeneous Catalysis, Elsevier, New York, 1992 p. 81.. 16. N. Mizuno, K. Katamura, Y. Yoneda, and M. Misono, J. Catal., 83, (1983) 384. 17. K. Eguchi, Y. Toyozawa, N. Yamazoe, and T. Seiyama, J. Catal., 83, (1983) 32. 18. I. Kawafune, Chem. Lett., (1989) 185. 19. M. Che, and T.A.J., Adv. Catal., 31, (1982) 77. 20. R. Neumann, and M. Levin, J. Am. Chem. Soc., 114, (1992) 7278. 21. C. Louis, J.M. Tatibou/~t, and M. Che, J. Catal., 109, (1988) 354. 22. Y. Iwasawa, Chemical Design Surfaces for Active Solid Catalysts, (D.D. Eley, H. Pines, P.B. Weisz, Eds.), vol. 35, p. 187. Academic Press, New York, 1987. 23. J.M. Tatibou~t, and J.E. Germain, J. Catal., 72, (1981) 375. 24. J.M. Tatibou6t, J.E. Germain, and J.C. Volta, J. Catal., 82, (1983) 240. 25. J.M. Tatibou~t, and J.E. Germain, C. R. Acad. Sci. Paris, 296, (1983) 613.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11 th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

345

Some evidences of a bifunctionnal mechanism for the reduction of NO on Pd based catalysts.

H. Praliaud a, A. Lemaire a, J. Massardier b, M. Prigent c and G. Mabilon c a Laboratoire d'Application de la Chimie h l'Environnement, Unite Mixte C.N.R.S. Universit6 Claude Bernard Lyon I, 43 boulevard du 11 novembre 1918, 69622 Villeurbanne Cedex, France b Institut de Recherches sur la Catalyse, C.N.R.S., conventionne a l~niversit6 Claude Bernard Lyon I, 2 avenue Albert Einstein, 69626 Villeurbanne Cedex, France c Institut Franr

du Petrole, 1-4 Avenue du Bois Preau, 92506 Rueil - Malmaison, France

ABSTRACT Catalysts with Pd deposited on various supports such as ZrO 2 or mixed supports such as AI203-ZrO2-BaO have been used for the elimination of NOx. Their activity in the presence of complex mixtures like CO-NO-O2-C3H6-CO2-H20 is better than that of Pd/Al20 3. This increase in activity is due not to a modification of the electronic properties of Pd but to a direct participation of the support in the process. From the analysis of the experimental results, a bifunctional mechanism is proposed and discussed.

I. I N T R O D U C T I O N For the elimination of NOx in the exhaust gases of gasoline motors, more and more effort has been devoted to the elaboration of new catalysts, potential substitutes of Pt-Rh based solids. Indeed, with such materials, the Rh/Pt ratio is higher than in the Pt-mine. It is economically important, therefore, to elaborate new catalysts either without Rh or with a lower amount of Rh than the catalysts of the present day. In the present work various Pd-based catalysts are prepared using supports such as ZrO 2 or mixed supports such as Al203-ZrO 2 and Al2Oa-ZrOE-BaO instead of Al203. The catalytic behaviour of these new solids is compared to that of an unmodified Pd/Al203 and of a Pt-Rh/Al203 solid which is considered as a pefformant three-way catalyst. Physic,o-chemical measurements are undertaken in order to determine the Pd electronic state and the redox and acidic properties of the supports. It is shown that these new supports induce an enhancement of the NO reduction which cannot be explained by a modification of the Pd electronic properties. A bifunctional mechanism is proposed and discussed in addition to the intrinsic roles of the metal and of the

346 support. For instance the support plays also specific roles in the stabilisation toward thermal sintering of Pd and in the change in the adsorption of the reagents due to the acidic properties. 2. EXPERIMENTAL 2.1. Materials Two simple supports have been used, a transition y (8) alumina (around 120 m2 g-I) and a tetragonal ZrO 2 (70 m2 g-l) with a low amount of the monoclinic form. The mixed support AI203-ZrO 2 (AIZr) was prepared by wet impregnation of AI203 with Zr(CsH702) 4 dissolved in toluene. After evaporation of the solvent and drying at 373 K, the acetylacetonate was decomposed under 02 up to 773 K (0.5 K/min). The AI203-ZrO2-BaO (A1ZrBa) support was obtained by dry impregnation of AI203ZrO 2 by barium nitrate, drying at 423 K and calcination under a flow of 0 2 at 723 K (2 K/min). It contains the y(5) AI20 3 support, Ba compounds (oxide, nitrate, carbonate), the tetragonal form of ZrO 2 and very probably a BaZrO 3 phase characterized by its X-ray line at 2.95 angstroms. The Pd-based catalysts were prepared by wet impregnation of the support with Pd(CsH702) 2 dissolved in toluene. After drying at 373 K, the acetylacetonate was decomposed under 02 up to 773 K (0.5 K/min). The catalytic behavior of these solids was compared to that of a reference Pt-Rh/AI20 3 solid. Some characteristics of the solids are reported in Table 1. Table 1 Main physico-chemical characteristics of the solids: weight percents of metals and additives, specific surface area, metallic dispersion. m

% Pd (weight) % additives SBET(m2 ~ - 1 ) Pd/Al203 b 06 120 Pd/ZrO2 0.88 70 Pd/AIZr 1.01 9.8 (Zr) Pd/AIZrBa 1.0 8.7(Zr),6.6(Ba) 92 Pt-Rh/AI20~ c l%Pt, 0.2% Rh 120 a) determined by H 2 chemisorption b) catalyst prepared by the Institut Fran~ais du Petrole (IFP) c) prepared by coimpregnation with H2PtCI6 and RhC13 .

.

.

.

Dispersion a 0.4 0.31 0.4 0.19 0.64

2.2. Physicochemical and catalytic measurements The redox properties of ZrO 2 and Pd/ZrO 2 were studied by programmed thermooxidations and therrnoreductions. The reduced and degassed samples were heated under a He + 1% 0 2 mixture from room temperature up to 773 K (10 K/rain), cooled down to 298 K under the He + 1% 02 mixture and purged with He. Then, the thermogrammed reduction was carried out with an Ar + 1% H 2 mixture. After degassing the adsorbed hydrogen under Ar the oxidation/reduction sequence was repeated. The electronic properties of Pd were studied by both physical and chemical techniques: X-ray photoelectron spectroscopy (XPS) and infrared spectroscopy using CO and NO as probe molecules

347 The XPS experiments were performed with the ESCALAB 200 R (Fisons Instruments) using the AIKo~ line (1486.6eV). The binding energies have been corrected of the charge effects (insulator supports) by using the values of the elements in the oxides: Zr3d level of ZrO 2 at 182.2 eV and Al 2p level of Al20 3 at 73.8 eV. Before XPS analysis, the conditioned samples, previously reduced or treated under the reactant mixtures, were transferred into the spectrometer without exposure to air. The infrared absorbance spectra were recorded at room temperature on a Fourier transform spectrophotometer (Brucker I.F.S. 110) with a resolution of 4 cm -1. To compare the integrated adsorbances of the various samples the weight of the pellet and the Pd content were considered. The samples were placed in a heatable cell where the catalysts were treated "in situ". Different kinds of experiments were carried out: i) adsorption of CO, NO and pyridine at 298 K, ii) heating of the solids under CO, NO or CO+NO mixtures, iii) CO+NO reaction performed in the infi'ared cell and followed by the adsorption of CO at 298 K after evacuation of the gaseous phase. The catalytic experiments were performed at the stationnary state and at atmospheric pressure, in a gas flow microreactor. The gas composition (NO, CO, O 2, C3H 6, CO 2 and H20 diluted with He) is representative of the composition of exhaust gases. The analysis, performed by gas chromatography (TCD detector for CO 2, N20, 02, N2, CO and flame ionisation detector for C3H6) and by on line IR spectrometry (NO and NO2) has been previously described (1). A small amount of the sample (10 mg diluted with 40 mg of inactive ct AI203 ) was used in order to prevent mass and heat transfer limitations, at least at low conversion. The hourly space velocity varied between 120 000 and 220 000 h-1. The reaction was studied at increasing and decreasing temperatures (2 K/min) between 423 and 773 K. The redox character of the feedstream is defined by the number "s" equal to {2 JOEl+[NO ] }/ {[CO]+9[C3H6] }.

3. RESULTS 3.1. Influence of the water addition to the CO-NO-O2-C3H6 mixture When H20 (10 vol %) is added to the CO-NO-O2-C3H 6 mixture, variations in activities are observed, as shown in Figure 1 The influence of H20 depends on the support, leading to a clear enhancement of the conversion (reduction of NO into N2) activity with respect to the light-off temperatures (T50) or to the total conversion [Pd/ZrO 2 : AT50 (NO) = - 45 K, AT50 (CO) = - 50 K, 100% NO conversion at 663 K instead of 89% NO conversion at 698 K] or to pratically no effect (Pt-Rh/AI203) A small enhancement is also observed, above 573 K, with Pd/A1203 From these results, a sequence for the NO reduction activity can be given: Pd/Al203 < Pd/ZrO2, PtRh/Al203. The NO reduction is higher with the ZrO 2 support than with the Ai20 3 one.

In the goal of optimizing the catalytic activities two more complex solids (Pd/A1203ZrO 2 and Pd/Al203-ZrOE-BaO ) have been tested. The mixtures were CO-NO-O2-C3H6 and CO-NO-OE-C3Hr-H20 (s - 1.03).

348

1 0 0

:

,

,~,~--

(a)

;._~ 0

.

~

se

(b)

0

0

* $

o

-~I ~ 473

573

0 673 Temperature

-

04)4) 4) 4) ,

8

o IP oo,oo* <

,--~

473

573

9 673

(K)

Figure 1. Influence of addition of H20 to the CO-NO-O2-C3H6 mixture (s = 1.03) on the CO (O, o) and NO (~,, O) conversions for the Pd/AI203 (a) and Pd/ZrO 2 Co) catalysts. (o, O) in the absence of H20, (O, ~ ) in the presence of H20 (10 vol. %).

In the absence of H20 the light-off temperature, T50, for NO is strongly improved on the two optimized catalysts (Table 2). The light-off temperature becomes close to that of the reference Pt-Rh/AI203 solid. Near T50 the conversions obey the sequence: PtRh, Pd/AIZrBa > Pd/AIZr > Pd/AI, Pd/Zr. At higher temperatures the NO conversion is enhanced for the Pd/AI-Zr-Ba and Pd/AI-Zr solids but the Pt-Rh/AI203 solid remains the best one. At the maximum of conversion, the activities obey the order: Pt-Rh > PdAIZrBa, PdAIZr > PdZr > PdAI Table 2 Light-off temperature, T50 (K), for NO in the presence of the CO-NO-O2-C3H6 and CO-NOO2-C3H6-H20 mixtures (s = 1.03) (flow rate 22 l/h) (4568 vpm CO, 623 vpm NO, 5611 vpm 0 2, 799 vpm C3H6, 10 vol. % H20 ). T50(NO)(K ) Pd/AI20~ Pd/ZrO2 Pd/AI20~- ZrO 2 Pd/AI203- ZrO2-BaO pt-Rh/Al203 ,,,

CO-NO-O2,C3H6 650 660 610 585 575

CO-NO-O2-C~FI6-H20 690 620 665 .. 590 5~/5

In the presence of water, with the exception of Pt-Rh/AI203, the best Pd-based catalysts contain ZrO 2 (Table 2, Figure 2), either alone (Pd/ZrO2) or associated with Ba (Pd/Al203-ZrOE-BaO). Above 673 K, Pd/ZrO 2 is the best catalyst for the NO reduction but the Pd/Al2Oa-ZrO2-BaO solid has the highest performances in the overall range of temperature. Near T50 the NO conversions obey the sequence: PtRh >Pd/AIZrBa> Pd/Zr > Pd/AIZr > Pd/Al.

349 At the maximum the sequence becomes: PtRh, Pd/Zr > Pd/AIZrBa > Pd/AIZr > Pd/AI It is noteworthy that the Pd/ZrO 2 and Pd/A1203-ZrO2-BaO activities are close to that of PtRh/AI203

.

100

.

.

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.

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.

, c~-."~ . ~ c , ~

A

qr~JP ~

9

~A |Ik,,

Q g=

o o

0 Z

50-

9

.

~ e

9

e" %* . , , . . - . - - * ~ ~. ........ 473,"" " 573 6 } 3 "" Temperature (K)

7;'73

Figure 2. NO conversion with the CO-NO-O2-C3H6-H20 (10 vol. %) (s = 1.03) mixture on: Pd/AI203, A Pd/ZrO2, m Pd/AI203-ZrO2 , ~ PtRh/AI203 and 9 Pd/AI203-ZrO2-BaO. Let us also notice that: i) the changes in the activity sequences with the temperature show the necessity to compare the activities at various temperatures, not only near the light-off temperature ii) since the activity sequences depend also on the mixture, for example on the presence or on the absence of H20, it is necessary to study complex realistic mixtures. Furthermore whatever the solid the N20 quantity remains moderate. 3.2. Influence of the support on the Pd electronic state

In order to control the effect of ZrO 2 on the electronic properties of Pd, an infrared study (using the CO probe molecule) and an XPS determination of the Pd binding energies have been performed. The infrared spectrum of CO irreversibly adsorbed at room temperature on the Pd/AI203 solid previously reduced at 673 K and evacuated at 623 K shows (Figure 3) the classic v CO bands assigned to lineary bonded CO (2065 cmq) and to multi-bonded CO (1960 and 1925 cm-l) on metallic Pd [2-7]. The low frequency bands are more precisely attributed to bridged species adsorbed on the (100) or (110) (1960 cmq) and (111) (1925cm q) faces of Pd crystaUites. The position of the bands is coverage-dependent with a downwards frequency shift as the CO coverage decreases upon evacuations at increasing temperatures. Upon evacuation, the linear and multibonded species are no longer seen at 423 K and 573 K, respectively. The spectra of CO irreversibly adsorbed at 298 K on Pd/ZrO 2 (Figure 3), Pd/AI203ZrO 2 and Pd/AI203-ZrO2-BaO reduced at 673 or 773 K are similar to the spectrum of CO on

350 Pd/A]20 3. Linear CO species (band at 2070 cm-1) and multi-bonded CO species (band between 1970 and 1945 cm -1 and shoulder between 1925 and 1920 cm -1) on Pd ~ are observed. The variations in wavenumbers are not strong enough to indicate electronic effects rather than changes in CO coverage.

1960 1925

. m , m ~

2065

t ~

1119'0

'

..2..2 0 0

'| ............2 "i 'l -O00

....... 1 8 0..... 0 i ' --'cm-1

Wavenumber

Figure 3. Infrared spectra of CO irreversibly adsorbed at 298 K on Pd/AI20 3 (a) and Pd/ZrO 2 (b) reduced at 673 K and evacuated at 623 K

/

Zr

31~/,

................. 349; .... ; Binding

Pd3d5/2

I

......341; ............... ; ............ 333'

; ..... eV

Energy

Figure 4. Original and decomposed XPS spectra of the Pd/AI203-ZrO 2BaO reduced "in situ" Pd 3d and Zr 3p photopeaks.

This has been corroborated by the complememary XPS study in which the spectra were recorded without air contacting the reduced samples. As already discussed, the binding energies are referred to an internal standard, ie., the AI 2p line of AI203 (73.8 eV) or the Zr 3d line of ZrO 2 (182.2 eV). The binding energies 0~d 3d5/2 and Pd 3d3/2 levels) are characteristic of non-modified metallic palladium ffigure 4) [8-10]. There is no detectable difference between the spectra of Pd/AI20 3 and the deconvoluted spectra Pd/Al2OyZrO2-BaO. In the case of Pd/ZrO 2, the low Pd content with respect to ZrO 2 and the closeness of the binding energies of Pd 3d and Zr 3p levels do not permit reliable deconvolution. We can conclude however that the electronic state of Pd ~ is not noticeably modified whatever the support.

351 3.3. Characterization of the supports From the previous results, it has been proven that the nature of the support, although it has no significant influence on the Pd electronic properties, modifies the catalytic properties of the solids. To permit a better understanding of these supports effects, the surface properties of the supports (in the presence of the metal) have been studied, in particular the acidic properties and the oxygen mobilities. The AI203 and ZrO 2 supports have been mainly onsidered.

As already described (1) the inhibition of the NO reduction by CO due to carbon deposits in mixtures containing hydrocarbons depends on the support, with the most acidic supports leading to higher amounts of carbon deposits. The nature (Bronsted or Lewis centers), the number, and the strength of the acidic sites of the Pd/AI203 and Pd/ZrO 2 solids have been checked using infrared spectroscopy of adsorbed pyridine and thermoprogrammed desorption of ammonia. The infrared spectra were recorded after equilibrating the reduced and evacuated solids with an excess of pyridine vapor and further evacuation at various temperatures. After evacuation at 423 K there is no more physically adsorbed pyridine. There is no characteristic band of pyridine adsorbed on Bronsted acid sites (no appearance of the 19b vibration at 154045 cm-1) [11,12]. The OH groups observed on the solids are thus non acidic. The existence of Lewis acid centers (coordinatively unsatured AP + or Zr 4+) is proven by the presence of the 19b vibration at 1440-50 cm-1 and of the 8a vibration at 1610-1620 cm-1. The absorbances of the 1440-50 cm -1 band show that the acidity difference between the Pd/AI203 and Pd/ZrO 2 solids is not significant. This is corroborated by a thermodesorption study of NH 3 adsorbed at 373 K, which shows that the quantity of adsorbed NH 3 is only slightly higher on Pd/AI203 (4.7 * 10-4 mol NH3/g solid) than on Pd/ZrO 2 (3.3 * 10-4 mol NH3/g solid). The different behavior between Pd/AI203 and Pc[/ZrO2 cannot therefore be explained only by difference in the acidic properties. Nevertheless after reduction of Pd/ZrO 2 and Pd/Al20-ZrO2-BaO , IR spectroscopy of adsorbed CO allows to detect, aside from the bands due to CO adsorbed on Pd ~ (2000-1850 cm-1), other absorption bands in the range 1200-1700 cm-1 characteristic of "carbonatecarboxylate-formate" structures [3, 13-16] located on the supports. These "carbonatecarboxylate" structures imply the formation of CO 2 arising either from the participation of surface oxygen species of the supports (since Pd has been totally reduced) according to the surface reaction COads + O s --> CO 2 + ~ (oxygen vacancy) or from the CO disproportionation. Such species (carbonates, carboxylates, formates ...) are not formed after CO chemisorption on Pd/AI20 3. With Pd/ZrO 2 previously reduced (no ionic Pd), the actions of C3H6 at 473 K or 673 K and of C3H8 at 673 K lead also to the formation of hydrogenocarbonates (1615, 1450, 1220 cm-1) due probably to the presence of"oxygen species" located on the support. To confirm this formation of reactive surface oxygen species and oxygen vacancies, thermoprogrammed oxidations and reductions have been performed on Pd/ZrO 2. On this Pd/ZrO 2, calcined under 0 2 at 723 K, reduced in flowing H 2 at 773 K and evacuated at 773 K, the oxidation is performed with a He-l% 02 mixture. At 298 K, the 0 2 consumption is weak (0.2 mmol.for 14.8 mmol. total Pd). This corresponds to the oxidation of

352 the superficial Pd according to the reaction: Pd s + 8902 --> PdsO. Upon heating from 298 K to 773 K (10 K/rnin), the 02 consumption reaches 10.3 retool, with the maximum of the peak at 658 K. The total 02 consumption (10.5 mmol. 02 for 14.8 mmol. Pd) exceeds the amount of O 2 needed for the Pd reoxidation into PdO. It can be assumed that this 02 excess (3.1 mmol.) is trapped by the support according to the reaction : ZrO2. x + x/2 02 --> ZrO 2. Then, the solid is cooled under the He-1% 02 mixture and an At-1% H 2 mixture is introduced at 298 K. The H 2 consumption which reaches 12.3 retool. (for 14.8 retool, total Pd) can arise from two main phenomena : the reduction of PdO and the adsorption of H 2. PdO + H 2 --> Pd ~ + H20 pd s + 1/~H2 ._> Pdsi_I (s: superficial) The H 2 absorption can be neglected since the H 2 pressure is low (17). During the heating from 298 K to 773 K (10 K/min) the H 2 consumption reaches 4.75 mmol. Thus, the total H 2 comsumption reaches 17 retool, for 14.8 retool, total Pd and it can be concluded that this excess o f H 2 corresponds to a partial ZrO 2 reduction: ZrO 2 + X H 2 --> ZrO2.x + H20. After a purge under Ar at 773 K a second temperature-programmed oxidation leads also to a consumption of 0 2 in excess with respect to the Pd amount. In the absence of palladium, the ZrO 2 support (after treatment under H 2 at 773 K in order to remove possible impurities) does not absorb 02 or H 2 during successive temperatureprogrammed oxidations and reductions until 773 K. Moreover, with the Pd/Al203 solid the 02 or H 2 consumption never exceed the amounts corresponding to Pd. In the presence of Pd, the ZrO 2 surface presents therefore some redox properties and, assuming an amount of 10 oxygen atoms per (angstrom) E, 2 or 3% of the superficial oxygen atoms would be reduced, leading to the formation of oxygen vacancies on ZrO 2. 4. DISCUSSION AND CONCLUSION Supports effects do not drastically modify Pd, this is shown by XPS and by IR spectroscopy of adsorbed CO. Nevertheless, the catalytic performances of the materials have been significantly improved with the supports containing zirconia. For instance, the activity of Pd/AI203-BaO-ZrO 2 nearly reaches the activity of the Pt-Rh/AI20 3 reference.

The role of the metal in the NO reduction by CO is clearly shown by CO, NO and CONO (1:1 mixture) adsorptions at room temperature and at reaction temperatures. i) At 298 K, the adsorption of NO and CO takes place on the same Pd centers and a competitive adsorption is observed. ii) Upon heating, NO is dissociatively adsorbed on Pd ~ according to the reaction : NOads --> Nads + Oads and Pd ~ is oxidized into Pd n+ as shown by the IR bands at 1815-1810 crn-l assigned to NO on Pd n+ and observed after heating under NO at 473 K and evacuation at 298 K. iii) The number of such Pd n+ sites is drastically decreased when the samples are put into contact, at 473 K, with CO or with a CO-NO (l:l) mixture. Thus CO 2 is detected in the gaseous phase. The Oads species created by the NO dissociation react with adsorbed CO according to the reaction COads + Oads --> CO 2 . This reaction cleanses the metal surface of excess adsorbed oxygen, permitting a new dissociative NO adsorption. These two reaction steps are generally invoked in the CO+NO mechanism [ 18-21 ] and, as can be concluded from the above results, occur mainly on Pd.

353

For the role of the support, two effects have to be considered i) An intrinsic effect: for instance the acidity or the stabilizing effect toward thermal sintering of Pd. The influence of the support acidity has been clearly illustrated by the hydrocarbon poisoning which is strongly decreased in the presence of less acidic supports [ 1]. ii) The effect of an oxygen-ions mobility which is connected to the redox properties of the support but needs the presence of the metal. In the present work, thermoprogrammed oxidations and reductions have shown the presence of "mobile and active oxygen atoms" on ZrO 2. With such solids, CO adsorbed on the metal can migrate to and react with these reactive oxygens giving CO 2 and an oxygen vacancy (in close proximity to the metal) according to the scheme : COads(Pd) +O- Zr-O --> CO 2 + [3- Zr-O. These oxygen vacancies can be directly involved in the water gas shitt reaction (CO + 1-120 --> CO 2 + 1-12)or in the NO dissociation or in a trapping of 0 2. In the first case the oxygen vacancy would be filled by 1-120, leading to the formation o f H 2 [22-25] according to the reaction: 1-120 + O- Zr- [3 --> H 2 +O- Zr-O In fact the water gas shitt reaction is much more promoted by the ZrO 2 support than by the AI203 one. While the CO conversion reaches 100 % at 623 K on Pd/ZrO 2 , it does not exceed 50 % at 773 K on Pd/AI20 3. Such a H 2 formation via the water gas shift reaction (or the steam reforming reaction which is also greatly enhanced on Pd/ZrO 2 ) implies the participation of both the metal and the support. H 2 is generally considered as a better reducing agent for NO than CO [18] and therefore, this H 2 formation via the water gas shitt reaction could explain the enhancement of activity observed in the presence of 1-120 for ZrO 2containing solids which possess "mobile oxygen atoms". Therefore, a bifunctional mechanistic scheme, including the participation of both the metal (via the adsorption of CO) and the support (via the formation of "oxygen vacancies" which are active sites for the H20 dissociative adsorption) seems quite relevant to explain the specific behavior, for the NO reduction in the presence of water, of samples containing ZrO2. Such active sites would be located at the metal-support interface and are linked to the redox properties of the support. In conclusion, a "specific role" of these dual "Metal-Support" sites at the metal-support interface has to be considered in addition of the "intrinsic roles" of the metal and the support and a bifunctional mechanism can be reasonably proposed. ACKNOWLEDGEMENTS This work has been carried out with the financial support of the "Groupement de Recherches catalyseurs d'epuration des gaz d'echappement automobile" funded by the "Centre National de la Recherche Scientifique, the "Institut Fran~ais du Petrole" and the PIRSEM (Programme Interdisciplinaire de Recherches Scientifiques pour rEnergie et les Matieres Premieres).

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5. .

7. 8. .

10. 11. 12. 13.

14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25.

A. Lemake, J. Massardier, H. Praliaud, G. Mabilon and M. Prigent, Catalysis and Automotive Pollution Control Ill, Elsevier, Amsterdam, (1995), in press. J.L. Duplan and H. Praliaud, Appl. Catal., 67 (1991) 325 A. A. Davydov, Infrared Spectroscopy of Adsorbed Species on the Surface of Transition Metal Oxides, (J. Wiley ed.), Rochester, (1984) A. M. Bradshaw and F. Hoffmann, Surf. Sci., 72 (1978) 513 D. Tessier, A. Rakai and F. Bozon-Verduraz, J. Chem. Soc. Farad. Trans., 88 (1992) 741 P. Gelin, A. R. Siedle and J.T. Yates, J. Phys. Chem., 88 (1984) 2978 A. Palazov, G. Kadinov, Ch; Bonev and D. Shopov, Surf. Sci., 188 (1987) 505 J. Z. Shyu, K. Otto, W.L.H. Watkins, G. W. Graham, R.K. Blitz and H.S. Gandhi, J. Catal., 124 (1988) 2. V. Pitchon, M. Guenin and H. Praliaud, Appl. Catal., 63 (1990) 333 Tran Thanh Phuong, J. Massardier and P. Gallezot, J. Catal. 102 (1986) 456 G.M. Alikina, A.A Davydov, I.S. Sazonova and V.V. Popovskii, Kinetik. i Katal. 28 (1987) 418 J. W. Ward, J. Catal., 10 (1968) 34 and 11 (1968) 271 A. Laachir, V. Perrichon, A. Badri, J. Lamotte, E. Catherine, J.C. Lavalley, J.E1Fallah, L. Hilake, F. Le Normand, E. Quemere, G.N. Sauvion and O. Touret, J. Chem. Soc. Farad. Trans.,87 (1991) 1601 Can Li, F. Domen, K.J. Maruya and T. Onishi, J. Catal., 141 (1993) 540 C. Binet, A. Badri, M. Boutonnet- Kizling and J.C. Lavalley, J. Chem. Soc. Farad. Trans., 90 (1994) 1023 E. Guglielminotti and F. Boccuzzi, J. Catal., 141 (1993) 486 E. Wicke and H. Brodowsky, Topics in Applied Physics, 29 (1978) 73 W. F. Egelhoff, The Chemical Physics of solid surfaces and Heterogeneous Catalysis,(D. A. King and D. P. Woodruff, eds.) Vol. 4 (1984) 397 D' Arc3, Lorimer and A.T. Bell, J. Catal. 59 (1979) 223 W. C. Hecker and A.T. Bell, J. Catal., 84 (1983) 200 S. H. Oh, G. B. Fisher, J. E. Carpenter and D. W. Goodman, J. Catal. 100 (1986) 360 B. Harrison, A. F. Diwell and C. Hallet, Platinum Metals Rev., 32 (1988) 73 J. C. Schlatter and P. J. Mitchell, Ind. Eng. Chem. Prod. Res. Dev., 19 (1980) 288 J. Barbier, Jr. and D. Duprez, Applied. Catal. B., 4 (1994) 105 M. Weibel, F. Garin, P. Bernhardt, G. Make, and M. Prigent, Catalysis and Automotive Pollution Control II, Elsevier, Amsterdam, 71 (1991) 195

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

355

Characterization of Pd-based Automotive Catalysts R. W. McCabe and R. K. Usmen Ford Research Laboratory, MD 3179, SRL, Dearborn, MI 48121-2053 1. ABSTRACT Characterization studies were undertaken to determine the cause of large differences in activity between various commercial automotive catalysts after aging for 75 or 120 h on an accelerated engine-dynamometer cycle. In all, a set of nine catalysts was examined, comprised of both fresh and aged Pal-only, Pd/Rh, and Pt/Rh catalysts. Catalyst activity, as measured by CO/NOx crossover efficiencies in dynamometer airfuel sweep tests, showed no correlation with either noble metal dispersion or noble metal surface area. The amount of stored oxygen required to obtain 100% CO conversion in the A/F-modulated dynamometer sweep experiments was estimated at 15 p-mol O/g-cat.. Bench reactor experiments involving both titration of pre-oxidized catalysts with CO and cyclic CO oxidation confirmed that the 15 p-mol O/g-cat. storage requirement represents a threshold level separating high- and low-activity catalysts. Formation of bulk PdO is the main oxygen storage mode in the 120 h dynamometeraged Pd-based catalysts. Pd loading is important: the higher the Pd loading, the greater the capacity for oxygen storage via PdO. Dispersion of Pd in the dynamometer-aged catalysts (2-6%) was too low to account for significant oxygen storage via chemisorbed oxygen. Furthermore, temperature-programmed reduction experiments and comparisons of oxygen uptakes on catalysts with and without rare earth oxides both indicated that the rare earth oxides play little role in oxygen storage after 120 h dynamometer aging. Not only does the formation of bulk PdO account for the quantities of oxygen stored in the aged catalysts, but the observed rates of oxygen uptake are consistent with bulk Pd oxidation kinetics reported by Remillard et al [J. Appl. Phys. 71 (9), 1992, pp. 4515-4522]. 2. INTRODUCTION Three-way automotive catalysts based on palladium, rather than the more expensive metals platinum and rhodium, have long been desired. However, Pd is more sensitive than Pt to poisoning by lead (Pb) compounds [1-4]. Consequently, widespread commercial use of Pd-based automotive three-way catalysts (TWC) was delayed in the U.S. until the early 1990s, by which time residual Pb concentrations in unleaded gasoline had decreased to negligible levels. The past five years have witnessed

356

considerable research and development of various types of Pd-based TWCs including Pd-only [5-11 ], Pd/Rh [12-15], and Pt/Pd/Rh '~rimetal" catalysts [16,17]. Aside from its historically low price compared to Pt and Rh, Pd has distinct catalytic properties which make it a desirable component of today's three-way catalysts. Chief among these is thermal durability, particularly Pd's ability to maintain activity under high-temperature lean (i.e. excess O2) conditions. Pd also has excellent light-off characteristics, especially when deployed at higher concentrations than traditional Pt/Rh catalysts. Highly loaded Pd-based catalysts are thus an obvious choice for socalled close-coupled or starter catalysts mounted close to the exhaust manifold [5,18,19]. Such catalysts reach operating temperatures much faster than underbody catalysts but also experience higher warmed-up operating temperatures and greater risk of thermal deactivation. The present study was initiated to understand the causes of large differences in performance of various catalyst formulations after accelerated thermal aging on an engine dynamometer. In particular, we wished to determine whether performance characteristics were related to noble metal dispersion (i.e. noble metal surface area), as previous studies have suggested that the thermal durability of alumina-supported Pd catalysts is due to high-temperature spreading or re-dispersion of Pd particles [20-

25].

Catalyst performance (as evaluated in dynamometer sweep evaluations) did not correlate with noble metal particle dispersion. Instead, bench reactor experiments involving titration of preadsorbed oxygen with CO showed that total Pd load in@ rather than Pd surface area is the key factor affecting performance. Pd serves as its own oxygen storage agent through formation of bulk PdO, and the amount of PdO formed depends primarily on the amount of Pd available, not the surface area of the Pd. 3. EXPERIMENTAL 3.1. Catalysts Table 1 lists characteristics of the catalysts. Those labeled "TWC" are commercial formulations from Ford's catalyst suppliers, each letter designating a different formulation. The commercial catalysts all contained various rare earth and alkaline earth oxide promoters and stabilizers in addition to the noble metals. Two simple Pd-on-alumina reference catalysts (A and B) were prepared in our laboratory by impregnating alumina-coated ceramic monoliths with aqueous solutions of Pd nitrate. The laboratory-prepared catalysts were dried and calcined in air at 550 ~ for 5 h prior to evaluation. Noble metal concentrations were determined by x-ray fluorescence. Some of the commercial formulations were aged on an engine dynamometer for 75 or 120 h according to a standard 4-mode aging procedure with an inlet exhaust gas temperature of 760 ~ (peak catalyst temperature ca. 900 ~ [6]. The dynamometer aging cycle simulates vehicle aging of catalysts - 75 h for 50,000 miles and 120 h for 100,000 miles.

357

Table 1 Catalyst Description Catalyst

Source

Aging

A) Pd/AI203 B) Pd/AI203 C1) Pd-only TWC C2) Pd-only TWC C3) Pd-only TWC D) Pd-only TWC E) Pd/Rh TWC

Lab Lab Supplier Supplier Supplier Supplier Supplier

Fresh Fresh Fresh 75 h 120 h Fresh 120h

F) Pd/Rh TWC

Supplier

120h

G) Pt/Rh TWC

Supplier

120h

Pd concentration (%) 0.30 0.79 0.66 0.66 0.66 0.60 0.30 (0.026 Rh) 0.33 (0.036 Rh) (0.21 Pt) (0.042 Rh)

Oxygen storage component (~) No No Yes Yes Yes No Yes Yes Yes

(1) rare earth oxide component

32. ~ meUxx~ Conversion efficiencies of the dynamometer-aged catalysts were measured in a standard A/F sweep test on an engine dynamometer [6]. The sweep experiments were carried out at 450 ~ and 85,000 h"1 space velocity (volumetric basis; standard conditions). The sweep ranged from 0.5 A/F lean of stoichiometry to 0.5 A/F rich of stoichiometry with imposed A/F perturbations of +_0.5A/F at 1 Hz. After sweep evaluation, small samples of catalyst were removed from the front region of the brick for chemisorption and flow reactor experiments. Chemisorption measurements employed the CO-methanation technique of Komai et al [26] as modified in our laboratory [27] for application to automotive catalysts. In particular, modifications were made to the pretreatment to avoid the formation of Pd hydrides [27]. We have found the method well-suited to automotive catalysts, both because of its high sensitivity and apparent freedom from complications due to adsorption of CO on sites other than noble metal sites. In previous studies involving a series of Pd/Rh and Pt/Rh TWCs aged on vehicles, we obtained good correlation between apparent dispersions (and noble metal surface areas) measured by the CO methanation technique and those determined from both x-ray diffraction and transmission electron microscopy [28,29]. As with other chemisorption methods, the CO-methanation technique does not distinguish between different noble metals. Thus, Rh was treated equivalently to Pt or Pd, and a standard adsorption stoichiometry of 1 CO molecule per surface metal atom was assumed for all three noble metals. In general, contributions from Rh are expected to be small due to the low concentrations employed (one-tenth the Pd loading in the case of the Pd/Rh catalysts). Thus, the Rh concentration was simply added to the Pd or Pt loading and treated as Pd or Pt.

358

T~ration of pre-dosed oxygen by CO was carded out at 500 ~ in a 1" o.d. quartz reactor tube housing a 3/4" diameter by 1/2" long catalyst button. The reactor contained two solenoid-controlled three-way valves th= were used to inject alternating pulses of secondary feed streams (each 0.1 I.Jmin) into a main feed of N2 carder gas (2.9 L/min). Two types of experiments were conducted. One involved pre-dosing the catalyst for 120 s with 0.87% 02 in the main feed. The O2 flow was then stopped, the system was purged for 120 s, and the adsorbed oxygen was titrated from the catalyst by injecting alternating 15 s pulses of CO and N2 into the N2 carder stream (the CO concentration was 0.3% after dilution with the carder N2 stream). The second type of experiment exposed the catalyst to alternating pulses of CO (0.3%) and 02 (0.185%) for durations of 15 s for 02 and variable times between 1 and 7 s for CO. Conversion of CO and formation of CO2 were monitored by non-dispersive infrared analyzers. 4. RESULTS

DyrBmonteter ev uaSons

Figure 1 shows sweep data for two of the dynamometer aged catalysts: the 120 h aged Pd-only ((33) catalyst (Fig. 1A), and the 120 h aged Pd/Rh (E) catalyst (Fig. 1B). Despite equivalent aging, the Pd-only catalyst gave much higher conversions, especially around the stoichiometric point. CO/NOx cross-over efficiencies of the other dynamometer-aged catalysts are reported in Table 2. The Pd-only catalyst stands out, showing crossover efficiencies in excess of 95% after both 75 and 120 h aging. In contrast, the formulations containing Pd/Rh or PURh have crossover efficiencies between 50 and 54%.

Figure 1. Engine dynamometer sweep plots of 120-hr aged (A) Pd-only (C3) and (B) Pd/Rh (E) catalysts at 450~ and modulations of +0.5 A/F at 1Hz.

359

4.2. Noble metal dispersions and surface areas Table 2 lists the apparent dispersions obtained from the CO methanation technique. No correlation is observed between dispersion and catalyst performance as measured by the CO/NOx crossover efficiencies. The C2 and C3 Pd-only TWCs, despite their extremely high CO/NOx crossover efficiencies, gave apparent dispersions of 3.5 and 3.0% after 75 and 120 h aging versus higher values of 5.9% for the Pd/Rh catalyst (E) and 4.3% for the Pt/Rh catalyst (G), both of which displayed low CO/NOx crossover efficiencies. Even between the two Pd/Rh catalysts, catalyst E has an apparent dispersion more than four times that of catalyst F, yet the two are nearly identical in their CO/NOx crossover efficiencies. Table 2 Catalyst Properties Catalyst

CO/NOx crossover efficiency

(%)

A) Pd/AI203 B) Pd/AI203 C1) Pd-only TWC C2) Pd-only TWC C3) Pd-only TWC D) Pd-only TWC E) Pd/Rh TWC F) Pd/Rh TWC G) Pt/Rh TWC

NA NA NA 96 98 NA 50 52 54

Dispersion (%)

10.1 9.5 10.8 3.5 3.0 6.6 5.9 1.6 4.3

NM Surface Area (m2/g-cat.)

0.17 0.42 0.40 0.13 0.11 0.22 0.11 0.03 0.04

Table 2 also lists the noble metal surface areas normalized to the total mass of the catalyst. The surface areas were calculated directly from the dispersion data taking into account the different mass of noble metal in each catalyst and assuming a constant site density of 1x1019/m2. As with dispersion, no clear correlation exists between mass-specific noble metal surface areas and CO/NOx cross-over efficiencies. 4.3. Oxygen tJtr"~n expedrnents The engine dynamometer sweep evaluations were carried out under modulated air-fuel conditions of + 0.5 A/F ratio at 1 Hz. To achieve high conversions under these conditions, the catalyst must store oxygen during lean excursions in order to convert CO and HC under rich excursions. Likewise, some of the oxygen related inhibition of NOx reduction on the lean side is mitigated by replenishment of oxygen to the storage agent. The lack of a correlation between the CO/NOx crossover frequencies of these catalysts and either noble metal dispersion or mass-specific surface area suggests, in turn, that oxygen storage in aged catalysts is not strongly dependent on either noble metal dispersion or noble metal surface area.

360

120 h aged Pt/Rh (G) Feed CO A M U

~2 tn q0 c o r m 4r b., t) ,,4

-

Fresh (Cl) Pd-Only

-

0

CJ

Feed

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1

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6

CO Pulse Number

Figure 2. Trtration of pre-adsorbed oxygen by pulses of CO (0.3% CO in N2) at 500 ~

Experiments involving titration of preadsorbed oxygen with pulses of CO in nitrogen were carded out to assess the oxygen storage capacity of both the fresh and dynamometer aged catalysts. Figure 2 shows the pulse profiles for the two catalysts with the highest (CI) and lowest (G) oxygen uptakes. The areas under the CO pulses were integrated to determine the amount of oxygen pre-adsorbed during the 120 s O2 pretreatment at 500 ~ W'rth the exception of the fresh CI Pd-only catalyst, all of the other catalysts produced CO breakthrough in excess of 90% after the third pulse. The cumulative O-atom uptakes, expressed as /~-mol O/g-cat., are summarized in Table 3. Both the first-pulse uptakes and the cumula'dve uptakes corresponding to the first three pulses are listed. Table 3 also contains the theoretical oxygen uptakes associated with both chemisorbed oxygen (i.e. assuming one chemisorbed O-atom per surface noble metal atom) and formation of bulk noble metal oxides (i.e. PdO or PtO).

Table 3 Oxy.qe.n........capac...ities(#-mol O/Q-~.)... Catalyst

- Theoretical O uptakesSurface O ~ Bulk O 2

- O titrated by CO 1st pulse 1st 3 pulses

A) Pd/AI20 3 B) Pd/AI20 s C1) Pd-only TWC C2) Pd-only TWC C3) Pd-only TWC D) Pd-only TWC E) Pd/Rh "IWC F) Pd/Rh TWC G) PURh TWC

1.0 1.7 6.8 2.7 1.8 3.4 1.9 0.4 0.6

11.5 24.8 35.5 27.2 21.6 20.9 9.2 12.7 5.2

28.2 74.2 62.0 62.0 62.0 56.4 30.6 34.4 14.0

1~ ...assumes i'" O-atom per surface""or""buJk noble metal atom. '..........

15.6 32.0 66.8 39.0 32.0 31.0 14.3 23.0 7.2

361

Two key observations can be drawn from the data in Table 3: 1) the oxygen consumed in either the first pulse or the first three pulses is much greater than that which can be attributed to oxygen adsorbed on the surface of the noble metal particles, and 2) with the exception of ~ fresh (CI) Pd-only IWC, none of the other catalysts have either first-pulse or cumulat~e CO uptakes that exceed the theoretical oxygen uptake associated with formation of bulk noble me'~ oxide. Even though all of the aged commercial TWCs contain oxygen storage agents, the quantities of oxygen taken up after 120 h dynamometer aging are not sufficient to require storage via those agents. 40

40

z~Pd w OS A "" Pd w/o OS w 30 • Pt/Rh a. o * Surface O ~ 20

& ,e=e

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.

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7O 80 20 30 40 50 60 Bulk oxygen capacity (as PdO) Figure 3. Plot of O-atoms consumed in the first CO pulse (/~-mol O-atoms/g-cat.) vs. the bulk oxygen capacity of the catalyst (PdO basis). 0

10

4.4. Role of bulk PdO The importance of bulk PdO as an oxygen storage component is illustrated in Figure 3 which plots the amount of oxygen consumed in the first CO pulse versus the theoretical bulk oxygen capacity of each Pd-containing catalyst (expressed as/~-mol PdO/g-cat.). The solid curve is fit to the data for the three catalysts which do not contain rare earth oxygen storage agents plus an additional point at the origin reflecting the experimental observation that negligible oxygen is consumed on a blank alumina catalyst. Given the absence of rare earth oxides, and recognizing that the amount of stored oxygen is far greater than that available from chemisorption (as shown by the data points denoted by diamonds in Fig. 3), the only other source of oxygen is from reduction of bulk PdOo Thus the solid curve can be taken as representative of bulk oxidation of Pd during the 120 s exposure to O2 at 500 ~ Note that all of the 12=3h dynamometer aged catalysts have oxygen uptakes on or below the curve defined by the catalysts which do not contain rare earth oxygen storage agents. This suggests that rare earth oxides do not contribute significantly to oxygen uptakes after 120 h dynamometer aging. The only catalysts showing greater oxygen

362

storage are the fresh and 75 h aged Pd-only TWCs (C1 and C2), and it is likely that rare earth oxides do contribute to oxygen uptakes in those catalysts. Interestingly, the C1 and C2 catalysts are the only pair which show a correlation between oxygen uptake and noble metal dispersion (i.e. the oxygen titrated by the first CO pulse drops from 35.5 to 27.2/~-mol O/g-cat. as the dispersion drops from 10.8% (C1) to 3.5%

(C2)).

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Figure 4. 1-12TPR traces of fresh (CI) and b) 120 h aged (C3) Pd-only TWCs (30 ~ heating rate in 9% i"La/Ar). Inset: TPR traces of a 0.32% Pd/15% CeO2/AI=O3 catalyst fresh and after aging for 1 h at 900 ~ in a laboratory flow reactor. 4.5. TPR The absence of a strong contribution from rare earl~ oxygen storage agents is confirmed by hydrogen temperature-programmed reduction (TPR) experiments. Figure 4 shows hydrogen TPR traces for the fresh and 120 h aged Pd-only catalysts (C1 and C3). The fresh catalyst contains a broad peak centered around 100 ~ which has an area of 127/~-mol 1-12/g-cat. (corresponding to twice the theoretical I-Iz uptake of 62 /~-mol I-Iz/g-cat. required to reduce PdO to Pd metal). The "extra" area is attributed to the reduction of surface oxygen associated with ceria in intimate contact with Pd as shown by the close similarity to the TPR trace of a fresh 0.32=,(, Pd/15% CeO2/AI203

363

catalyst prepared in our laboratory (Fig. 4 inset). After 120 h dynamometer aging, the low-temperature feature of the C3 catalyst peaks near 40 ~ and decreases in area to 23/~-mol H2/g-cat. (Fig. 4b). Again, the changes upon aging are similar to those of the reference 0.32% Pd/15% CeO2/AI20~ catalyst after aging for 1 h at 900 ~ in a laboratory reactor (Fig. 4 inset). Changes in TPR spectra of the type shown by both the Pal-only catalyst and the model Pd/CeO2/AI20 s catalyst are characteristic of large reductions in contact area between Pd and rare earth oxides effected by thermal aging [30]. Pd in intimate contact with ceria [30,31] or lanthana [32] gives a TPR feature between 100 and 200 ~ whereas Pd on alumina has its chief TPR feature at temperatures between 0 and 60 ~ [23,32,33]. In fact, partial reduction of PdO may be occurring in our experiments prior to the start of the temperature ramp from 0 ~ The TPR data thus support the oxygen uptake experiments in suggesting that rare earth oxides contribute little to the uptakes observed for the 120 h dynamometer-aged catalysts. 5. DISCUSSION Automotive catalysts have traditionally been designed with the objective of maximizing dispersion of the noble metals. Some exceptions exist, such as the well-documented observation that oxidation rates of saturated hydrocarbons increase with increasing noble metal particle size [34,35]. However, one of the main objectives in designing automotive catalysts is to stabilize the dispersion of the noble metals against thermal sintering. Pd undergoes complex changes in oxidation state and structure in response to variations in temperature and gas environment [20-25,33,36-40]. Spreading and/or redispersion of Pd (as a cationic form of Pd) has been reported in oxidizing environments at temperatures between 700 and 900 oC [20-25]. At first glance, such a mechanism would appear to offer a plausible explanation for the thermal durability of Pd-basecl catalysts, particularly in aging cycles of the type reported here involving lean (i.e., oxidizing) modes. The data of this study, however, indicate that dispersions of Pd-based catalysts after thermal aging are quite low, and no greater than comparably aged PURh catalysts. Catalytic activ'rty, as measured in modulated NF sweep experiments, neither correlates with noble metal dispersion nor with noble metal surface area. As shown in Fig. 3, oxygen uptakes (as reflected in amounts of CO oxidized) show an increasing trend with the bulk oxygen capacities of the catalysts (expressed as the theoretical amount of bulk PdO which can be formed). Comparing the oxygen uptake data of Fig. 3 to the CO/NOx crossover efficiencies of Table 2, a threshold level of oxygen uptake is suggested, between 13 and 20/~-mol O/g-cat., above which the catalyst can store enough oxygen to ensure high efficiency during the A/F perturbations encountered in the dynamometer sweep test. This range is consistent with a stored oxygen demand of 15.3/~-mol O/g-cat. which we have estimated as required to ensure stoichiometric oxidation of all reducing species during the 0.5 s, -0.5 delta NF half cycle of the dynamometer sweep (centered at the stoichiometric A/F ratio) [41]. To a first approximation, the high activity of catalyst C3 results simply from its having an oxygen uptake capacity above the threshold requirement, whereas the other

364 dynamometer-aged catalysts have oxygen uptakes below the threshold requirement. The oxygen uptakes shown in Fig. 3 reflect much longer oxygen exposures (120 s) and larger CO titers (48 #-mol CO/g-cat.) than charactedstic of the dynamometer sweep cycle. Therefore, additional bench reactor CO oxidation experiments were carried out at shorter time intervals cycling between 15 s pulses of O2 (0.185 mol%) and 2 to 7 s pulses of CO (0.3 mol%). Results are summarized in Table 4, with the catalysts listed in order of highest to lowest CO conversions. The most active catalyst (C1 Pd-only TWC) stored sufficient oxygen to convert a 7 s CO pulse with 98% efficiency. At the other extreme, the PURh catalyst (G) gave only 48% conversion of a 7 s CO pulse and reached only 9'2% conversion with a 2 s pulse. The most pertinent condition for comparing the bench reactor and dynamometer results is at 5 s, since the bench reactor CO dose at 5 s (16/~-mol CO/g-cat.) is close to the dynamometer stored oxygen demand (15.3/~-mol O/g-cat.) Significantly, we find close quantitative agreement between the CO conversions at 5 s CO pulse length and the CO/NOx crossover efficiencies reported in Table 2. In both cases, catalysts E, F, and G give conversions in the 50-55% range whereas the C2 and C3 catalysts give conversions in excess of 85%. The good quantitative agreement between the engine dynamometer data and the laboratory pulsed CO oxidation experiments supports our interpretation that high CO/NOx crossover efficiencies in dynamometer sweep evaluations reflect oxygen storage above a threshold level. Table 4 Per cent conversion of CO pulses of various len.qths Catalyst

C1) Pd-only TWC C2) Pd-only TWC D) Pd-only TWC C3) Pd-only TWC B) Pd-AI20 3 A) Pd-AI20 s E) Pd/Rh TWC F) Pd/Rh TWC G) Pt/Rh TWC

7 98 NM NM NM 62 NM 42 NM 48

length of CO pulse (s) ............. 5 4 3 >99 95 89 86 78 67 51 51 NM

>99 >99 95 96 89 NM 67 NM 59

>99 >99 >99 > 99 >99 92 82 NM 71

>99 >99 >99 >99 >99 >99 >99 93 92

Given the connection between oxygen uptakes and catalyst sweep performance, the lack of a correlation between noble metal dispersion (and surface area) and catalyst performance implies that oxygen uptake does not depend on noble metal dispersion or surface area. The lack of a correlation is not surprising given that, 1) the quantities of oxygen involved, 2) the similarities in oxygen uptakes between catalysts with and without oxygen storage agents, and 3) the absence of TPR features characteristic of noble metal-rare earth oxide interactions all point to bulk oxidation of PdO as the

365

dominant oxygen storage mechanism in the 120 h dynamometer-aged Pd-based catalysts. For bulk oxidation of Pd to be the dominant oxygen storage mechanism, Pd oxidation kinetics must be rapid enough to account for the quantities of oxygen stored. Data reported by Remillard et al [38] on the air oxidation of thin Pd films indicate that the rates are indeed fast enough to account for the oxygen uptakes observed for the aged catalysts in these experiments. The 500 ~ 120 s pre-oxidation employed in the CO titration experiments would produce an oxide film thickness of 89 Angstroms using the growth expression reported in Ref. 38 (assuming equivalent kinetics in air and in the 0.87% O2/N2 mixture of our experiments). Taking the C3 catalyst, for example, with a dispersion of 3% (corresponding to a mean particle diameter of 373 Angstroms [42]), the formation of an outer PdO band of 89 Angstrom thickness would result in shrinkage of the metallic core to 274 Angstroms and growth of the overall diameter to 452 Angstroms. The increase in particle diameter owes to both the lower density of PdO compared to Pd and the presence of about 21% void volume in the PdO band [38]. Approximately 61% of the Pd in the particle is oxidized. Experimentally, 120 s oxidation of the C3 catalyst yielded a 3-pulse CO consumption corresponding to oxidation of about 52% of the Pd in the catalyst. Thus, the bulk oxidation kinetics are fast enough to account for the experimental observations. Even at the shorter 15 s oxygen exposure of the cyclic CO oxidation experiments, the kinetics of Remillard et al predict a 35 Angstrom thick oxide film. This corresponds to 25% oxidation of the Pd (or 15.2/~-mol O/g-cat.), just slightly less than the uptake of 16/~-mol O/g-cat. required for complete conversion of the 5 s CO pulse in the cyclic CO oxidation experiments (Table 4). Note that the oxidation kinetics of Remillard et al were obtained on 1-/~m-thick sputtered films of Pd on quartz, indicating that Pd oxidation kinetics are rapid enough, even for large-grain Pd particles, to account for oxygen uptakes required in both the dynamometer and bench reactor activity evaluations. The conclusions reached in the present study, namely that the catalytic activity of aged Pd-based catalysts depends primarily on Pd loading, begs the question '~hat is the function of the various additives such as ceria?". One obvious answer is that much effort has gone into stabilizing support components against loss of surface area and associated occlusion of noble metal particles. Both rare earth and alkaline earth oxides are important in this regard. In addition, we have addressed catalytic activity under very limited conditions -- modulated A/F sweep experiments carried out with lowsulfur fuel on an engine dynamometer. Factors such as noble metal dispersion and effects of promoters/stabilizers may be more important under other conditions, e.g. during light-off, at different temperatures and space velocities, with high-sulfur fuel, or in different A/F regimes during aging and evaluation. Also, we have focussed exclusively on CO oxidation, whereas different effects of dispersion and promoter/stabilizers may obtain for HC oxidation and for NOx reduction. Our study does suggest, however, that at least part of the Pd should be deployed in the form of large particles to obtain good CO and NOx conversions under modulated A/F conditions. The par~cles will consist of a metallic core with a roughened outer band capable of facile interconversion between Pd oxide and Pd metal.

366

6. SUMMARY Chemisorption measurements, combined with oxygen uptake, TPR, and pulsed CO02 experiments were employed to determine the source of large differences in dynamometer sweep performance of a series of Pt/Rh, Pd/Rh and Pd-only TWCs after dynamometer aging. The following observations have been made: 1) Apparent noble metal dispersions of 75 and 120 h dynamometer-aged TWCs range from about 2 to 6%. 2) CO/NOx cross-over efficiencies of aged catalysts in dynamometer sweep experiments do not correlate with either noble metal dispersion or noble metal surface area. 3) Oxygen uptakes in both dynamometer and bench reactor experiments at 500 ~ are much too great to attribute to oxygen chemisorption. 4) After 120 s oxygen exposure (500 *C), all of the dynamometer-aged Palbased catalysts gave oxygen uptakes that could be accounted for by the formation of bulk PdO. 5) Dynamometer-aged (120 h) catalysts showed no evidence for oxygen storage via rare earth oxides. 6) Formation of bulk PdO is the primary oxygen storage mechanism in the dynamometer-aged Pd-based catalysts. 7) A threshold level of oxygen storage (via bulk PdO) is required to reach high CO/NOx conversion levels in dynamometer sweep tests; Pd loading, rather than dispersion or surface area, is the most important factor affecting oxygen uptakes. 8) Rates of oxygen uptake in the dynamometer-aged catalysts are consistent with published oxidation kinetics of 1-/~m-thick Pd films. 7. ACKNOWI.EDGMENTS We thank K. S. Patel and D. M. DiCicco for providing the dynamometer-aged catalysts and sweep evaluation data. E. Gulari and C. Sze (U. of Michigan) assisted with the design of the pulsed reactor system. REFERENCES 1. M. Shelef, K. Otto, and N.C. Otto, Adv. in. Catal. 27 (1978) 311-365. 2. H.S. Gandhi, W.B. Williamson, E.M. Logothetis, J. Tabcock, C. Peters, M.D. Hurley, and M. Shelef, Surf. & Interface Anal. 6(4) (1984) 149. 3. W.B. Williamson, D. Lewis, J. Perry, and H.S. Gandhi, Ind. Eng. Chem. Prod. Res. Dev., 23 (1984) 531. 4. R.L Klimisch, J.C. Summers, and J.C. Schlatter, Amer. Chem. Soc. Adv. Chem. Ser. 143 (1975) 103. 5. J.C. Summers, J.F. Skowron, and M.J. Miller, Soc. of Automotive Eng., Paper

367

930386 (1993). 6. J.S. Hepburn, K.S. Patel, M.G. Meneghel, H.S. Gandhi, and Engelhard and Johnson Matthey Three Way Catalyst Development Teams, Soc. of Automotive Eng., Paper 941058 (1994). 7. M. Harkonen, M. Kivioja, P. Lappi, P. Mannila, T. Maunula, and T. Slotte, Soc. of Automotive Eng., Paper 940935 (1994). 8. J.C. Summers, J.J. White, and W.B. Williamson, Soc. of Automotive Eng., Paper 890794 (1989). 9. H. Muraki, Soc. of Automotive Eng., Paper 910842 (1991). 10. H. Tanaka, H. Fujikawa, and I. Takahashi, Soc. of Automotive Eng., Paper 930251 (1993). 11. T. Yamada, K. Kayano, and M. Funabiki, Soc. of Automotive Eng., Paper 930253 (1993). 12. Y.-K. Lui and J.C. Dettling, Soc. of Automotive Eng., Paper 930249 (1993). 13. J.K. Hochmuth and J.J. Mooney, Soc. of Automotive Eng., Paper 930219 (1993). 14. J.C. Summers, W.B. Williamson, and J.A. Scaparo, Soc. of Automotive Eng., Paper 900495 (1990). 15. H. Muraki, H. Sobukawa, M. Kimura, and A. Isogai, Soc. of Automotive Eng., Paper 900610 (1990). 16. B.H. Engler, E.S. Lox, I~ Ostgathe, T. Ohata, I~ Tsuchitani, S. Ichihara, H. Onoda, G.T. Garr, and D. Psaras, Soc. of Automotive Eng., Paper 940928 (1994). 17. A. Punke, U. Dahle, S.J. Tauster, and H.N. Rabinowitz, Soc. of Automotive Eng., Paper 950255 (1995). 18. D. Ball, Soc. of Automotive Eng., Paper 922338 (1992). 19. Z. Hu and R.M. Heck, Soc. of Automotive Eng., Paper 950254 (1995). 20. E. Ruckenstein and J.J. Chen, J. Catal. 70 (1981) 233. 21. J.J. Chen and E. Ruckenstein, J. Phys. Chem. 85 (1981) 1606. 22. E. Ruckenstein and J.J. Chen, J. Colloid Interface Sci. 86 (1982) 1. 23. H. Lieske and J. Volter, J. Phys. Chem. 89 (1985) 1841. 24. J.W.M. Jacobs and D. Schryvers, J. Catal. 103 (1987) 436. 25. J.G. McCarty and Y-F. Chang, Scripta Metal. et Mater. 31 (1994) 1115. 26. S. Komai, T. Hattori, and Y. Murakami, J. Catal. 120 (1989) 370. 27. R.K. Usmen, R.W. McCabe, and M. Shelef, "Proceedings of the Third Congress on Automotive Pollution Control," Elsevier, Brussels, Belgium, in press. 28. M.H. Yao, D.R. Liu, R.J. Baird, R.K. Usmen, and R.W. McCabe, ext. abstr., "Proc. 52rid Ann. Mtg. Microscopy Soc. of Amer.," C.W. Bailey and A~J. Garratt-Reed, eds., San Francisco Press, Inc., p.776, 1994. 29. M.H. Yao, D.R. Liu, R.J. Baird, R.K. Usmen, and R.W. McCabe, submitted to J. Catal., Nov., 1995. 30. R.K. Usmen, R.W. McCabe, G.W. Graham, W.H. Weber, C.R. Peters, and H.S. Gandhi, Soc. of Automotive Eng., Paper 922336 (1992). 31. B. Harrison, A.F. Diwell, and C. Hallett, Plat. Met. Rev. 32 (1988) 73. 32. S. Subramanian, R.J. Kudla, C.R. Peters, and M.S. Chattha, Catal. Letts. 16 (1992) 323. 33. T.R. Baldwin and R. Burch, Appl. Catal. 66 (1990) 359. 34. T. Tokoro, I~ Hori, T. Nagira, T. Uchyima, and Y. Yoneda, Nippon Kagaku Kaishi

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12 (1979) 1646. H. Yao, Y. Yao, and K. Otto, J. Catal. 56 (1979) 21. R.F. Hicks, Q. Haihua, M.L. Young, and R.G. Lee, J. Catal. 122 (1990) 295. P. Briot and M. Primet, Appl. Catal. 68 (1991) 301. J.T. RemiUard, W.H. Weber, J.R. McBride, and R.E. Soltis, J. Appl. Phys. 71 (1992) 4515. 39. D. Konig, W.H. Weber, B.D. Poindexter, J.R. McBride, G.W. Graham, and I~ Otto, Catal. Letts. 29 (1994) 329. 40. R. Burch and F.J. Urbano, Appl. Catal. A: Gen. 124 (1995) 121. 41. The stored oxygen demand was estimated from the difference between the actual conversions measured at the stoichiometric point with +0.5 A/F perturbations at 1 Hz and steady state conversions measured at the +0.5 and -0.5 A/F extremes. 42. F.H. Ribeiro, M. Chow, and R.A. Dalla Betta, J. Catal. 146 (1994) 537. 35. 36. 37. 38.

j.w. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

369

P r o c e s s d e v e l o p m e n t for the selective h y d r o g e n o l y s i s o f CC12F2 ( C F C - 1 2 ) into CH2F2 ( H F C - 3 2 ) A. Wiersma l, E.J.A.X. van de Sandt 2, M. Makkee 1, H. van Bekkum 2 and J.A. Moulijn I Department of Chemical Process Technology, Section Industrial Catalysis 2 Department of Organic Chemistry and Catalysis Delft University of Technology, Julianalaan 136, 2628 BL, Delft, The Netherlands

A palladium on activated carbon catalyst is a suitable catalyst for the selective hydrogenolysis of CC12F2 into CH2F2. Its stability is good: no deactivation occurred after 800 hours of operation. Even after addition of possible recycle components methane and CHC1F2, only minor deactivation was observed during 1600 hours of total operation. The catalyst performance and stability strongly depend on the hydrogen to CC12F2 feed ratio. A ratio of 6 leads to higher conversion, higher selectivity to CH2F2 and higher stability. At low H2 to CC12F2 ratios the mechanism of deactivation is presumably coke deposition. However, the ratio cannot be raised too high because then sintering of the palladium particles causes deactivation. Addition of methane to the feed leads to additional deactivation, and, therefore, methane is only allowed in a recycle stream when high hydrogen to CC12F2 feed ratios are applied. Addition of CHC1F2 to the feed does not lead to a higher CH2F2 yield. The reaction pathway to CHCIF2 and CH2F2 is different from the route to methane. The selectivity to methane depends on the adsorption mode of CC12F2 and is independent of both the CC12F2 and the hydrogen concentration. The ratio between CHC1F2 and CH2F2 is mainly determined by the CC12F2 concentration. A CFC-destruction process, which can destruct about 90% of the CFCs in use, based on this catalyst is both technically and economically feasible. The main features of this process are a liquid cooled, multi-tube fixed bed reactor, excess hydrogen, 100% CC12F2 conversion and a hydrogen recycle, in which methane is allowed.

1. INTRODUCTION The term CFCs is a general abbreviation for ChloroFluoroCarbons. They have been extensively used since their discovery in the thirties, mainly as refrigerant, foam blowing agent, or solvent because of their unique properties (non toxic, non flammable, cheap). However, after the first warning of Rowland and Molina [ 1] in 1974 that CFCs could destroy the protective ozone layer, the world has moved rapidly towards a phase-out of CFCs. Because the destruction of stratospheric ozone would lead to an increase of harmful UV-B radiation reaching the earth's surface, the production and use of CFCs is prohibited (since January 1, 1995 in the European Union and since January 1, 1996 worldwide). However, the depletion of stratospheric ozone will continue in spite of this prohibition. This is caused by slow diffusion of CFCs present in the troposphere to the stratosphere and

370 the eventual emission of the CFCs still in use [2]. It is, therefore, of utmost importance to prevent the banked CFCs (45% CC12F2 as refrigerant, 45% CCI3F as foam blowing agent) from being emitted into the atmosphere. Commercially available CFC destruction processes are incineration and pyrolysis, but these processes only have a limited capacity. The long term replacements for CFCs in their application as refrigerant are most probably HFCs. Especially HFC-134a (C2H2F4) is often mentioned, but the total number of possible replacements still increase. CH2F2 (HFC-32) or a mixture thereof are good replacements in heavy duty cooling applications, ' because CH2F2 has excellent cooling properties and in addition a lower global warming potential than HFC-134a. Combining the fact that there is not enough destruction capacity for the banked CFCs and that the market for CFCreplacements such as CH2F2 is growing, a challenging task is to convert the waste CFCs into valuable HFCs. At Delft University of Technology a catalytic process is under development in which the harmful CC12F2 is converted into the valuable, ozone friendly CH2F2. With this process both the waste materials CC13F, which can be converted into CC12F2 by use of HF, and CC12F2 can be converted. Because of the limited time available to develop this process, the catalyst development is strongly related to a future process design. Palladium on a purified activated carbon support has been selected as a very suitable catalyst for the reaction. We have reported that the performance of this catalyst looks very promising and that a CFC hydrogenolysis plant based on this catalyst is both technically and economically feasible [3-5]. This paper deals with the stability of the selected catalyst, the long term influence of the hydrogen to CC12F2 feed ratio on the catalyst performance and the influence of the possible recycle components methane and CHC1F2 on the performance of the catalyst.

2. E X P E R I M E N T A L 2.1. Preparation of catalyst Catalysts were prepared by incipient wetness impregnation of palladium chloride dissolved in HCI (CI/pd2+=10) on an activated carbon support. 50 Gram of the gas activated, peat based carbon extrudates (Norit RB1, d=l mm, 1=3-5 mm, BET=1060 m2/g) were consecutively washed with 0.5 M caustic hydroxide, water, 0.5 M hydrochloric acid, and water in a flow set-up prior to the impregnation. The catalysts were dried overnight at 373 K and subsequently activated at 623 K for 1 hour in nitrogen. Reduction of catalyst was done insitu with hydrogen at room temperature and subsequently the temperature was raised to the desired level under flowing hydrogen. The CC12F2 was introduced and after 10 hours of catalyst activation, the temperature was set at the experimental temperature.

2.2. Testing of catalyst The activity tests of the catalyst were carried out in a microflow reactor set-up in which all the high temperature parts are constructed of hastelloy-C and monel. The reactor effluent was analyzed by an on-line gas chromatograph with an Ultimetal Q column (75 m x 0.53 mm), a flame ionization detector, and a thermal conductivity detector. The composition of the feed to the reactor can be varied, besides the temperature, pressure, and space velocity. The influence of the recycle components CHC1F2 and methane was tested by adding these components to the feed. In total five stability experiments of over 1600 hours were performed. In each

371

experiment temperature, pressure, and WHSV were the same. The H2 to CC12F2 feed ratio ranged from 1.5 (hydrogen limiting) to 10 (excess hydrogen). The experimental set-up was the same for each stability test: After 830 hours on stream methane was added to the feed as to give a CC12F2 to methane feed ratio of about 1. The methane was removed from the feed after 1000 hours. After 1150 hours CHC1F2 was added to give a ratio of CHCIF2 to CC12F2 of 4 and was removed after 1300 hours. After 1650 hours on stream the experiments were ended. In the experiment with the hydrogen to CC12F2 feed ratio of 10, this ratio was doubled after 680 hours of operation.

3. R E S U L T S The main products of the CC12F2 hydrogenolysis are CH2F2, CHC1F2, and methane. The initial performance of the catalyst, after activation, is strongly dependent on the HE to CCI2F2 feed ratio. The influence of the H2 to CC12F2 feed ratio on the initial catalyst performance is depicted in Figure 1. Both the conversion and the selectivity to the desired product CHEF2 increase with increasing H2 to CC12F2 feed ratio, while the selectivity to CHC1F2 increases. The selectivity to methane is not influenced by the HE to CC12F2 feed ratio. These differences become even more pronounced as a function of time on stream. This can be seen in Figure 2 in which the influence of the HE to CC12F2 feed ratio on the catalyst performance after 700 hours of operation is depicted. 100

25

conv. CC12F2

I00

=

O

~176176

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12

o

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Figure 2. Catalyst performance after 700 hours.

Thus the catalyst performances change as a function of time on stream. The extend of these changes depends on the H2 to CC12F2 feed ratio. In Figures 3a and 3b the catalyst performance during the first 800 hours of operation is depicted for two different H2 to CC12F2 feed ratios. The results show that not only the catalyst performance is better at higher H2 to CC12F2 feed ratios, but also the stability of the catalyst is better. At a ratio of 1.5 both the catalyst activity and selectivity keep declining during the 800 hour experiment. At the ratio of 6, however, no selectivity change is observed after 300 hours of operation and after 450 hours also the activity of the catalyst remains unchanged.

372

100 e

~

~

~' S ~'~ 80 "; ~ 60 ~ "~ 40 ~

g

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~

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5

0

0 0

200 400 600 Time on stream [hi

800

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0 0

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(a) H2 / CC12F2 = 1.5. (b) H2 / CC12F2 = 6. Figure 3a, b. Performance of catalysts as a function of time on stream. The effect of further increase of the high H2 to CC12F2 ratio on the catalyst performance is depicted in Figure 4. In this experiment the ratio was doubled to 20 after 680 hours. The steady-state catalyst performance is not only more enhanced at the ratio of 10, but it is also reached far more quickly in comparison to the results in Figure 3b. However, an increase of the ratio to 20 does not lead to a further increase in catalyst performance and at this ratio the catalyst also tends to deactivate. 100

9

~" 80

conv. CCI2F2

ratio doubled

o 60 o 40 e,,,

o 20

25 20 15-, ._>. 10oo

(D

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100

200

300

400 500 600 Time on stream [h]

700

800

0 900

Figure 4. Performance of catalyst as a function of time on stream (after 680 hours H2 to CC12F2 feed ratio doubled from lO to 20). Table 1 summarizes the performances of the catalysts for different H2 to CC12F2 feed ratios after about 700 hours of operation. A low HE to CCl2F2 feed ratio is poor choice for three reasons: low conversion, low selectivity, and low stability, while a high ratio leads to a stable catalyst, with a superior performance. A stable catalyst performance has only been reached with both the HE to CC12F2 feed ratios of 6 and 10. When the ratio of l0 is doubled, the catalyst also tends to deactivate.

373

Table 1 Catal~,st performance after 700 hours ratio H2/CC12F2 1.5 2.2 3 6 10 20 ,,

,,, ,,,

conv. CC12F2 50 60 67 79 88 84

sel. CH2F2 53 67 75 85 88 88

sel. CHC1F2 37 24 16 7 4 3

sel. CH4 7 7 7 6 6 7

%

deact. 17 14 11 5 3 -

The effect of addition of methane to the feed is depicted in Figures 5a and b. During methane addition the amount of methane produced cannot be exactly determined and, therefore, the selectivity to methane is assumed to be the same as prior to the addition of methane. The influence of methane is larger at lower H2 to CC12F2 feed ratios. The addition causes dilution and will result in both a lower CC12F2 and hydrogen concentration. Thus, methane addition leads to a lower CC12F2 conversion. An unexpected effect is the higher selectivity to CH2F2, which coincides with a lower selectivity to CHC1F2. This selectivity change was observed in all experiments, but with higher H2 to CC12F2 feed ratio, the change in selectivity was less pronounced. After removal of methane from the feed, the conversion increases again, but does not return to its original value. As can be seen from Figure 5a, during the addition of methane additional deactivation of the catalyst was observed, especially at lower hydrogen to CC12F2 feed ratios.

~9 '~ "~ .,-

= 0

~

t._.

o

I00 90 80 7o 60 50 40 30 20 10 0 600

methane -.9 sel. CH2F2

se

2~"-

sel. methane 700

50 45 40 35 30 25 20 15 ~0 5 0

~" >, "~ ~ ~ ,___,

800 900 1000 1100 Time on stream [hi

(a) H2 [ CCI2F2 = 2.2. Figures 5a, b. Influence of addition of methane.

methane ' 100 25 I..91 sel CH2F2' ~' 90 Ll_ "~ 80 20 -~ 70 ~ 60 15 ~ 50 "~40 sel CHCIF2 10 "~ 30 ~ 20 sel methane I~5 ~ 10 ~ 0 600 700 800 900 1000 1100 Time on stream [hi I

l

I

>

'1

(b) H2 / CC12F2 = 6.

No influence was observed in the selectivities of addition of CHC1F2 to the feed even at low H2 to CC12F2 feed ratios. The influence of addition of CHC1F2 on the conversion is comparable to the influence of methane. Figure 6 gives a comparison of the CC12F2 conversions in all experiments as a function of time on stream. Both in the presence of methane and of CHCIF2 additional catalyst deactivation was observed, especially at the lower hydrogen to CC12F2 feed ratios. None of the catalysts restores its initial activity completely after the methane is switched off. After the CHC1F2 is switched off at the ratio of 6 the

374

catalyst stops to deactivate but does not restore its initial activity. Therefore, the catalyst deactivation is irreversible under the conditions applied. methane

I0080

..............

CHC1F2

' .............. i

..........

H2

,.__..

g 60 40

r

;'2

20 [

ti

0

0

250

500

750

I000

1250

1500

1750

Time on stream [h]

Figure 6. CC12F2 conversion as a function of time for several hydrogen to CCI2F2 feed ratios. The main by-products of the reaction were ethane, CH3F, CH3CI, and propane. The selectivities as a function of time are depicted in Figure 7a and b. At low hydrogen to CC12F2 feed ratios the selectivities to ethane and propane are higher. Addition of methane or CHC1F2 to the feed also leads to an increase in the selectivities to ethane and propane. methane CHCIF2

3

methane

.................................

CHC1F2

0

I sel CH3F

.

,

,

i ~

0 0

250

500

250

750 1000 1250 1500 1750

H2 / CC12F2 = 1.5.

,

~

,~

set, CH3C1

....-

750 1000 1250 1500 1750

Time on stream [hi

Time on stream [hi

(a)

500

.,

(b)

H2 / CC12F2 = 6.

Figure 7 a, b. Selectivities to by-products as a function of time on stream.

4. D I S C U S S I O N 4.1. Catalyst stability The catalyst activity, selectivity, and stability are simultaneously influenced by the H2 to CCI2F2 feed ratio. Also the time needed to reach a steady-state catalyst performance depends

375

on this ratio. Apparently it takes some time before the adsorption/desorption equilibrium between all the reactants has been established. The catalyst deactivation at lower H2 to CC12F2 ratio can be tentatively explained by coke deposition on the catalyst surface. As the concentration of hydrogen is increased, coke precursors are hydrogenated from the catalyst surface. The amount of ethane and propane produced, which can be an indication for coke formation, is proportional to the deactivation and the H2 to CC12F2 feed ratio. In separate studies with a model catalyst evidence for formation of a palladium carbide phase was found also indicating coke formation on the catalyst surface[6]. This suggests that high hydrogen concentrations should be used to avoid the coke deposition. It can, however, be expected that the hydrogen concentration cannot be increased infinitely because palladium is expected to sinter easy in a 100% hydrogen atmosphere at reaction temperatures [ 11 ]. On the other hand CFCs can be expected to regenerate a deactivated palladium on activated carbon catalyst, as has been claimed by Kellner [12]. Therefore, catalyst stability can be obtained when the hydrogen concentration is high enough to prevent coke formation and the CC12F2 concentration is high enough for in-situ stabilization of the palladium dispersion. This explains why only two stable operations, at hydrogen to CC12F2 ratio of 6 and 10, were observed. When the H2 to CC12F2 feed ratio of 10 is doubled to 20 the catalyst starts to deactivate. This is probably caused by sintering of the metallic palladium, but also the presence of an optimum in activity and selectivity as a function of disperion could explain the observed phenomena. Thus the determination of the dispersion under reaction conditions is important. The addition of methane leads to a lower CC12F2 concentration and a lower hydrogen concentration. Thus the conversion decreases by addition of methane. The addition of methane also leads to a higher selectivity to ethane and thus coke formation. This effect can best seen at the ratio of 6, where the stable catalyst starts to deactivate after addition of methane. At the ratio of 20 no influence of methane on the selectivities and no additional deactivation were observed. Therefore, the mechanism of deactivation is better explained by sintering at this H2 to CCI2F2 ratio. It is not clear whether the added methane interferes with the catalyst surface and has an influence on the desorption of methane from the surface or that the observed additional ethane is caused by a lower hydrogen concentration. There are, however, indications that the observed additional deactivation is caused only by a lower hydrogen concentration. This can be seen in the influence of addition of CHC1F2 to the feed. Although there is little influence of addition of CHC1F2 on the selectivities, a comparable influence on the catalyst stability as addition of methane is observed. The catalyst deactivation increases by the addition of CHC1F2, and this deactivation coincides with a higher selectivity to ethane. Thus, the most probable explanation for the additional deactivation is the lower hydrogen concentration, and, therefore, higher coke formation during addition of both CHC1F2 and methane. This explains why the effects of extra deactivation are less at higher ratios and not observed at the high ratio of 20.

4.2. Catalyst performance The catalyst performance depends on the H2 to CC12F2 feed ratio. The selectivities to CHEF2 and CHC1F2 are influenced by the HE to CCIEF2 feed ratio, while the selectivity to methane is independent of this ratio. We have previously proposed a reaction mechanism with serial reactions on the catalyst surface and minor readsorption of the intermediate products, which is depicted in figure 8 [4,5]. Thus the kinetics of the reaction follows mainly parallel reaction pathways, in which the selectivities are not influenced by the conversion, and a

376 remarkable low change in selectivities as a function of temperature is observed. Such a behaviour has also been observed by other authors [7-9]. A mechanism with a serial reaction on the catalyst surface can explain the observed changes in selectivity of CHCIF2 and CH2F2 as a function of H2 to CC12F2 feed ratio. A higher hydrogen concentration would lead to more hydrogen adsorbed on the catalyst surface thus favouring the adsorption steps in which more hydrogen are involved. However this mechanism does not fully explain the constant selectivity to methane. An explanation for this phenomenon would be that the methane is formed via a different route on the catalyst surface. This could be caused by either different sites on the catalyst surface or a difference in the adsorption of CC12F2. It is not likely that the methane is formed on different sites because the selectivity to methane remains unchanged in spite of deactivation of the catalyst. Therefore, the conclusion would be that the 'methane' and the 'non-methane' sites deactivate at the same rate. A difference in the adsorption of CC12F2 would be a more simple explanation. Figure 9 shows part of a reaction scheme for the CC12F2 hydrogenolysis reaction. In this scheme all the hydrogenolysis reactions starting from CC12F2 are depicted.

HO,F

7,F

[CCI2F'2 ~

C ,F OH,

CHCI2F "

12

/ [CHCIF2} 22

= CH2CIF 31

\/ l

Figure 8. Reaction mechanism with apparent parallel reaction kinetics.

= CH2CI2

-

32

=

= CH3CI

\/

40

',,

CHeF - - 41

CH, 50

Figure 9. Reaction scheme of CFCs starting from CC12F2.

If during the adsorption of CC12F2 a carbon-chlorine bond is broken the products CHC1F2 and CH2F2 are preferentially formed. The sequential reaction to methane via CH3F will be of minor importance. This can be expected because the carbon fluorine bond is stronger than the carbon chlorine bond. If on the other hand a carbon-fluorine bond is broken during adsorption, the reaction follows a different path. In this case intermediates, which can be expected to be more reactive, are formed, which results in a high selectivity to methane. The ratio between the selectivity to methane and the sum of the selectivities to CHCIF2 and CH2F2 depends on the adsorption mode of CCI2F2, which is independent of the hydrogen concentration. This explanation is consistent with the observation that the sum of the selectivities to CHC1F2 and CHEF2 is the same during each experiment. The lower selectivity to CHCIF2 during addition of methane is remarkable. Because of the lower hydrogen concentration during addition a higher selectivity to CHC1F2 can be expected. A tentative explanation would be that the selectivity to CHCIF2 is proportional to the CCI2F2 concentration. The higher selectivity to CHCIF2 is then caused by coupled desorption of CHC1F2 and adsorption of CC12F2. CHC1F2 is removed from the surface by adsorption of CCIEF2, and, therefore, the selectivity to CHCIF2 is proportional to the CCIEF2 concentration. This also explains the fact that CHC1F2 has no influence on the catalyst performance in spite of the fact that the reactivity of CHCIF2 is 5% of the CCl2F2 reactivity [4,5]. CHC1F2 does not adsorb on the catalyst surface in the presence of CC12F2.

377 Thus it is clear from both the remarkable constant selectivity to methane and the influence of addition of methane on the selectivities to CHCIF2 and CH2F2, that the serial reaction mechanism has to be modified in order to explain the observed phenomena. Therefore, more measurements will be performed in order to determine the adsorption of CFCs on the metal surface and the influence of other products, such as CH2Fz and HC1 on the catalyst performance [ 11 ].

4.3. Process design The research is strongly related to a future plant operation. From these stability tests a preliminary layout of a future process can be determined. Obviously a high hydrogen to CC12F2 feed ratio is needed. This leads to a higher conversion, higher selectivity, and higher stability. The excess hydrogen cannot be purged and has to be recycled. It can be expected that some light products such as methane and ethane will accumulate in the recycle. From the methane addition experiments it can be concluded that some build-up of methane in the hydrogen recycle can be tolerated without a negative influence on the catalyst performance. The additional methane has to be compensated for with a higher hydrogen to CC12F2 feed ratio to avoid coke formation. However, the hydrogen concentration cannot be raised too high and therefore, the amount of methane in the recycle is limited. The methane to CC12F2 feed ratio is estimated to have a maximum of 1. The amount of methane in the recycle can be controlled with a purge. A scheme of a plant layout is depicted in Figure 10. recycle

D.. purge

Hydrogen .._

product H2F~

oil water

Feed pretreatment

waste

acids

Reactor

Acid removal

Lights separation

Distillation

Figure 10. Process lay out. In the feed pretreatment section oil and water are removed from the recovered or converted CC12F2. The reactor type will be a multi-tubular fixed bed reactor because of the exothermic reaction (standard heat of reaction -150 kJ/mol). After the reactor the acids are selectively removed and collected as products of the reaction. In the light removal section the CFCs are condensed and the excess hydrogen is separated and recycled. The product CH2F2 is separated from the waste such as other CFCs produced and unconverted CC12F2. The waste will be catalytically converted or incinerated. A preliminary process design has shown that such a CFC-destruction process would be both technically and economically feasible.

378 5. CONCLUSIONS Palladium on activated carbon is a very suitable and stable catalyst for the selective conversion of CC12F2 into CH2F2. The performance and stability of the catalyst strongly depend on the H2 to CC12F2 feed ratio. At low H2 to CC12F2 feed ratio, the mechanism of deactivation is presumably coke deposition, while at high ratios sintering of palladium causes catalyst deactivation. An optimum in both catalyst performance and stability as a function of the hydrogen to CC12F2 feed between 6 and 20 is found. The reaction follows mainly parallel pathways, rather than the expected serial reaction. Methane is formed via a reaction pathway different from the formation of CHCIF2 and CH2F2. The ratio between the selectivity to CHC1F2 and CH2F2 is proportional to the concentration of CC12F2. A process for the conversion of CC12F2 into CH2F2 would include a multi-tube fixed bed reactor with a hydrogen recycle in which a limited amount of methane is allowed. This process would be both technically and economically feasible. REFERENCES 1. 2.

M.J. Molina and F.S. Rowland, Nature, 249 (1974) 810. United Nations Environmental Programme, "Report of the Ad-Hoc Technical Advisory Committee on ODS Destruction Technologies", (May 1992). 3. A. Wiersma, E.J.A.X. van de Sandt, M. Makkee, H. van Bekkum, and J.A. Moulijn, Dutch Patent Application, NL 94.01574, (1994). 4. A. Wiersma, E.J.A.X. van de Sandt, M. Makkee, H. van Bekkum and J.A. Moulijn, in Proceedings of the 1st World Congress Environmental Catalysis, Pisa, Italy, G. Centi, C Cristiani, P. Forzatti and S. Perathoner (Eds.),(May 1-5 1995) 171. 5. A. Wiersma, E.J.A.X. van de Sandt, M. Makkee, C.P. Luteijn, H.van Bekkum and J.A. Moulijn, Catal. Today, 27 (1996) 257. 6. E.J.A.X. van de Sandt, A. Wiersma, M. Makkee, H. van Bekkum and J.A. Moulijn, submitted for publication in catalysis today. 7. B. Coq, J.M. Cognion, F. Figueras and D. Tournigant, J. Catal. 141 (1993) 21. 8. G. Moore and J. O'Kelly, European Patent Application 508660, Imperial Chemical Industries (1992). 9. R. Ohnishi, W.-L. Wang and H. Ichikawa, Stud. Surf. Sc. Catal., 90 (1994), 258 10. H. Jin, S.Y. Jeong, B.S. Kim, J.M. Lee and S.K. Ryu, in Proceeding of the 6th carbon conference, Granda, Spain, (July 3-8 1994), 352. 11. E.J.A.X. van de Sandt, A. Wiersma, M. Makkee, H. van Bekkum, and J.A. Moulijn, to be published. 12. C.S. Kellner, J.J. Lerou, V. Rao and K.G. Wuttke, WO patent 91/04097, Du Pont de Nemours, (1991).

Acknowledgement: R. Artan, R.A.M. Avontuur, B.J. Bezemer, J.H.A. van den Hoogen, N.B.G. Nijhuis, Y. Verbeek, and C.P. Luteijn are greatfully acknowledged for their work in the preliminary process design. Sponsors:

AKZO-Nobel, Dutch Ministry of Housing and Environmental Affairs, AlliedSignal Fluorocarbon Europe B.V., Johnson Matthey Plc., KTI b.v., and the European Union.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis- 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

379

Catalytic fluorination o v e r c h r o m i u m oxides. Preparation of hydrofluorocarbons

S. Brunet, B. Boussand and J. Barrault

Laboratoire de Catalyse - URA CNRS 3 50 - Eeole Sup6rieure d'Ingenieurs de Poitiers 40 Avenue du Recteur Pineau - 86022 POITIERS Cedex - FRANCE

Abstract

The fluorination of CF3CH2CI into CF3CH2F over chromium oxides is accompanied by a dehydrofluorination reaction (formation mainly of CF2=CHCI). This dehydrofluorination is responsible for the deactivation of the catalyst. A study of the dehydrofluorination reaction of CF3CH2CI proves that the reaction is favoured when the degree of fluorination of chromium oxide increases. Consequently it would be favoured on strong acid sites. Adding nickel to chromium oxide decreases the formation of alkenes and increases the selectivity for fluorination while the total activity decreases. Two kinds of active sites would be present at the catalyst surface. The one would be active for both the reactions of dehydrofluorination and of fluorination, the other only for the fluorination reaction.

1. INTRODUCTION The F / CI exchange in chloroalkanes is a route to HFCs. For example, different routes can be possible for the synthesis of CF3CH2F [ 1,2 ]. Our focus is on its preparation from CF3CH2CI and HF with chromium (III) oxide as a catalyst. This fluorination is accompanied by a dehydrofluorination which produces chloroalkenes (mainly CF2=CHCI ) resulting in a deactivation of the catalyst. Indeed this haloalkene could polymerise and thus lead to coke formation. The reactions involved are : CF3CH2C1 + HF

.~

~ CF3CH2F + HCI

CF3CH2CI

-.-

~

CF2 = CHCI + HF

The catalytic activity for F/CI exchange depends on the amount of reversibly oxidized sites with a linear relationship between the activity and the number of such sites [3,4].

380 The aim of our work is to study, under adequate operating conditions, the dehydrofluorination reaction of CF3CH2CI so as to determine the nature of the sites involved in the fluorination and the dehydrofluorination of CF3CH2CI. Thus a selective poisoning of dehydrofluorination sites would allow to increase the selectivity for the fluorination reactions. With this in view, we studied the development of the dehydrofluorination reaction of CF3CH2CI as function of the degree of fluorination of chromium oxide. Moreover, nickel and chromium oxide catalysts were prepared and tested for the dehydrofluorination reaction. Nickel oxide, a basic compound [5], could poison selectively the sites involved du~ng the dehydrofluorination reaction.

2. EXPERIMENTAL

2.1. Catalyst preparation The chromium oxides were prepared according to the following procedure [6-8]. Chromium oxide resulted from the dehydration of chromium hydroxide obtained by the addition of an ammonia solution (5M) to a solution of chromium nitrate (0.5M). The final pH was equal to 7.5 and the hydroxide formed was kept constantly stired and heated at 80~ for 1 h so as to obtain complete precipitation. This solid was filtered and washed three times with hot distilled water and dried for 16 h in an oven at 90~ It was then submitted to a dynamic thermal treatment under nitrogen at 380~ for 8 h. The chromium oxide formed was cooled down under the same vector gas. The mixed chromium and nickel (5% and 10% Ni atomic) catalysts were prepared by dehydration of mixed chromium and nickel hydroxides prepared by adding an ammonia solution to a solution of chromium and nickel nitrate in order to maintain a pH = 7+_1. The final pH was equal to 7.5. The hydroxyde treatment was the same as the one already described for chromium oxide.

2.2. Catalytic fluorination and dehydrofluorination The fluorination of CF3CH2CI was carried out at 380~ under atmospheric pressure in a fixed bed dynamic reactor. In a first operation, the catalyst was fluorinated "in tim" for 2 or 70 hours by HF mixed with nitrogen (ratio N 2 : HF: = 5 : 4) at 380~ CF3CH2CI was then injected into the reactor in the presence of HF. The operating conditions of this fluorination of CF3CH2CI were: temperature = 380~ catalyst weight = 50 mgo contact time = O.Ols, HF : CF3CH2CI : N 2 = 4 : 1 : 5. The same operation was repeated with a nonfluorinated catalyst. The catalytic activity for the CF3CH2F formation was measured both over the prefluorinated and the non-fluorinated catalyst after a 2 hour reaction with CF3CH2CI. The products resulting from the reaction were injected with an automatic sampling valve into a GIRA GC 181 gas phase chromatograph and analyzed with a flame ionization detector. The separation was carried out in a capillary column DB5 (J and W Scientific).

381 The catalytic activity for the dehydrofluorination of CF3CH2CI was measured at 320~ under atmospheric pressure, in a pulse flow reactor [9]. Pulses of pure CF3CH2CI were injected into a helium stream every ten minutes. The amount of ehloroeompound was adjusted (40.9 pmol) in order to obtain a conversion ea 10 %. The products resulting from the reaction were injected into a Varian 3400 gas phase chromatograph and analyzed with a flame ionization detector. The separation was made in a capillary column BP5 (SGE). The catalytic activity of the catalyst was measured at~er a 5 hour reaction with CF3CH2CI (by the amount of ehloroalkene formed : CF2=CHCI , CFCI=CHCI (Z and E)). 2.3. Fluorine titration

The fluorine titration of chromium oxide was carried out at the Elf-Atoehem Research Center, Pierre-B6nite.The catalyst mineralization was carried out in a Parr bomb by reaction with sodium peroxyde. Fluorine ions were then titrated by a potentiometric method with a specific fluoride electrode. 3. RESULTS 3.1. Transformation of CF3CH2C! over a non fluorinated chromium oxide

The transformation of CF3CH2CI was studied at 320~ in a pulse flow reactor. Indeed, in a dynamic reactor, the significant alkene formation leads to a rapid deactivation of the catalyst. The reaction is carried out in absence of I-IF in order to favour the dehydrofluorination reaction. Products distribution is shown in Fig. 1. 1,50

E

],oo

00000000

0

E

.~0,50
1

2

Time (h)

3

4

Figure 1. Products of CF3CH2CI transformation without HF over chromium oxide versus the time (T = 320~ flow reactor, 1 pulse = 40.9~tmol CF3CH2CI ) (OCF3CH2F (2 CF2=CHCI

A CFCI=CHCI Z) o CFCI=CHCI E)

382 The main product of the reaction was CF2=CHCI. Moreover, two secondary alkenes CFCI=CHCI (Z and E) were formed from CF 2 = CHCI by successive chlorine-fluorine exchanges [ 10]. After a fast initial deactivation, the catalyst stabilized. The catalytic activity for the dehydrofluorination reaction was estimated after a 5 hours reaction from the sum of the amount of alkenes. The formation of CF3CH2F could also be observed although not very significant because limited by the production of I-IF liberated during the dehydrofluoration reaction. Moreover at the end of the experiment there was less CF3CH2F formed than CF2=CHCI, which means that the formation of CF3CH2F was limited both by the formation of CF2=CHCI and by the fluorination of the catalyst. Consequently the catalytic activity for the CF3CH2CI fluorination reaction was measured in the presence of an excess of HF in a fixed bed dynamic reactor (fig. 2). The catalyst deactivation was then low, CF3CH2F being preponderantly formed. CF2=CHCI is the only alkene formed.In those conditions the formation of this alkene is at the equilibrium state. The catalytic activity for the fluorination reaction was calculated after a 2 hour reaction. 0,50 t~ 0,40"7

o

0,30 "

E E

0,20 ._~
,,

,

0

I

I

1

2

Time (h)

l

I

3

4

5

Figure 2. Products of CF3CH2CI transformation with HF present over chromium oxides versus the time (T = 380~ ; HF " CFaCH2CI 9N 2 = 4 91 " 5) (O CF3CH2F

Q CF2=CHCI )

3.2. Influence of the degree of fluorination of chromium oxide

In table 1, we show the influence of the degree of fluorination of the chromium oxide (measured by F/Cr atomic ratio.) on the transformation of CF3CH2CI. We remarked that the longer the fluorination of the catalyst lasts, the greater the degree of fluorination Moreover this fluorination decreases slightly the specific surface of the catalyst.. Fluorination and dehydrofluorination activities per unit of surface show the same evolution whatever the degree of fluorination However the dehydrofluorination reaction compared to the fluorination reaction is favoured. The ratio between the catalytic activity for the fluorination reaction and the

383 dehydrofluorination reaction decreases when the degree of fluorination increases. The fluorination of chromium oxide leads therefore to a decrease in the selectivity for the fluorination reaction.

Table 1 Catalytic activity for the fluorination of CF3CH2CI and for the alkene formation after 5 hours of CF3CH2CI pulses over chromium oxide with different degree of fluorination Fluorination time (hour)

0

2

70

Specific area (m2.g - 1)

178

114

111

Degree of fluorination

0

0.6

1.60

F/CI exchange activity (mmol.h-l.m -2)

0.31

0.48

0.23

Dehydrofluorination activity (mmol.h- 1.m-2)

1.0

2.0

1.4

F / CI exchange activity dehydrofluorination activity

0.31

0.24

0.17

(F/CO

3.3. Effect of nickel addition on chromium oxide

Two Ni-Cr catalysts containing one 5 % nickel atomic and the other 10 % were tested for fluorination and dehydrofluorination reactions of CF3CH2CI. The results were compared to those obtained with chromium oxide alone. These reactions are performed only with the non fluorinated catalysts (table 2). The presence of nickel did not affect the specific surface, nevertheless the activity for the fluorination reaction decreased particularly for the catalyst with 10 % nickel. Moreover a significant decrease in the formation of alkenes could be observed for the Ni-Cr catalysts. Consequently the presence of nickel disfavoured the dehydrofluorination reaction while increasing the selectivity for the fluorination reaction.

384 Table 2 Catalytic activity for the fluorination of CF3CH2CI and for the alkene formation after 5 hours of CF3CH2CI pulses over mixed Ni-Cr oxydes (Ni/(Ni + Cr) = 5 % or 10 %) or Cr203 Catalyst

0

5

10

Specific area (m2.g - 1)

178

198

206

F/CI exchange activity (mmol.h-l.m-2)

0.31

0.24

O.18

Dehydrofluorination activity (mmol.h-l.m-2)

1.0

0.38

0.42

F / Cl exchange activity dehydrogenation activity

0.31

0.63

0.43

Ni/(Ni + Cr) (molar %)

4. DISCUSSION

We observed the dehydrofluorination reaction was favoured when the degree of fluorination of the catalyst increased. Indeed, we showed that the ratio between fluorination activity and dehydrofluorination activity decreased when the degree of fluorination of the catalyst increased. Thus when the degree of fluorination increased, the selectivity for the fluorination reaction decreased. The dehydrofluorination reaction required the rupture of the C-F bond of the CF3CH2CI molecule while the fluorination reaction involved the rupture of the C-CI bond. The C-F bond being harder to split than the C-CI bond [ 11], the dehydrofluorination reaction require stronger adsorption sites than the fluorination reaction. The fluorination of chromium oxide caused an increase of surface site Lewis acidity. Kemnitz and al.[12] as well as Peri [13], showed that on fluorinated alumina the progressive substitution of F for O and OH led, thanks to inductive attracting effect of fluorine, to an increase of the Lewis acidity of r sites. Hence the dehydrofluorination reaction was favoured on strong acide sites. The nickel addition in chromium oxide decreased the formation of alkenes which was smaller than the one observed in the presence of just chromium oxide. It is to be remarked that the decrease of alkene formation was independent of the quantity of nickel in the catalyst. However, the catalytic activity for the fluorination reaction decreased when the nickel content increased. Thus the addition of nickel in small quantities allowed to increase the selectivity for the fluorination reaction.We could suggest that nickel substitute

385 chromium atom on the surface and modify the catalyst strong acidity which is principaly involved in the dehydrofluorination reaction. The presence at the catalyst's surface of active sites which made possible the dehydrofluorination reaction (C-F rupture) and the chlorine-fluorine substitution (C-CI rupture) was then supposed. There would also exist active sites which would only allow the second reaction (C-CI rupture). The addition of nickel, by suppressing the sites which allow the two reactions (fluorination and dehydrofluorination) decreases the total activity. Thus, it would be possible to increase the selectivity for the fluorination reaction, nethertheless, this would induce a decrease of the catalytic activity.

Acknowledgements We are grateful to the Elf Atochem company for their financial support.

REFERENCES 1. L.E. Manzer, Catal Today 13, (1992) 13.1 2. D.R. Coulson, P. W. J. G. Wijnea, J. J. Lerou, and L. E. Manzer, J. Catal 140, (1993) 103 3. J. Barrault, S. Brunet, B. Requieme and M. Blanchard, J. Chem, Soc, Chem., Commun, 374 (1993). 4. S. Brunet, B. Requieme, E. Matouba, J. Barrault and M. Blanchard, J. Catal., 152, (1995) 70-74. 5. 6. 7. 8. 9. 10.

J.M. Trillo, G. Munuera, and J. M. Griado., Cat. Rev., 7, (1973) 51-83. R.C. Burwell Jr., G. L. Hailer, K. C. Taylor and J. F. Read. Adv. Catal., 20 (1969) 4. W. Gamier, L. M. Lettoffe and P. Llaudy, Bull. Soc. Chim. Ft., 91 (1984) 3-4. R.C. Burwell Jr., G. L. Hailer, K. C. Taylor and J. F. Read, Adv. Catal., 20 (1969) 29. D.Duprez, J. Chim.. Phys. 80(1983)487-505. S. Brunet, B. Requieme, E. Colnay, J. Barrault, M. Blanchard, Appl. Catal., B, 5, (1995) 305-317. 11. L. Kolditz, G. Kauschka, W. Schmidt, 2. Anorg. Chem., 41-54, (1977). 434. 12. J. B. Peri, J. Phys. Chem. 72,8, (1968) 2917-2925. 13. A. Hess and E. Kemnitz, J. Catal., 149 (1994).449-457.

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J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996Elsevier Science B.V. All rights reserved.

387

O x i d a t i v e C o u p l i n g of M e t h a n e to E t h y l e n e w i t h 85% Yield in a G a s R e c y c l e E l e c t r o c a t a l y t i c or C a t a l y t i c R e a c t o r - S e p a r a t o r M. Makria, Y. Jiang b, I. V. Yentekakis a and C. G. Vayenas a aDepartment of Chemical Engineering, University of Patras, GR 26500 Patras, Greece bDalian Institute of Chemical Physics Chinese Academy of Sciences, Dalian 116012, China Methane was oxidatively coupled to ethylene with very high yield in a novel gas recycle reactor-separator operated in a batch or continuous flow mode. The recycled gas passes continuously through a molecular sieve trap in the recycle loop which adsorbs and thus protects from further oxidation a controllable percentage of ethylene and ethane. The products are obtained by subsequent heating of the molecular sieve trap. Ethylene yield up to 85%, i.e. 88% selectivity to ethylene at 97% CH 4 conversion, has been obtained in the batch mode of operation using a Ag-Sm203 or Ag electrocatalyst and electrochemical supply of oxygen through a ZrO2-Y20 3 solid electrolyte. Using the continuous flow mode of operation with gaseous 0 2 supply in the recycle loop and a Sr(lwt%)/La20 3 catalyst we have obtained ethylene yields up to 50%, i.e. 65% C2H4 selectivity at 76% CH4 conversion. The synergy of the catalytic and molecular sieve materials is discussed and modelled in view of the predominantly consecutive nature of the oxidative coupling of methane (OCM) network. 1. INTRODUCTION Keller and Bhasin were first to report in 1982 [1] on the catalytic one-step oxidative dimerization or "coupling" of methane (OCM) to C2 hydrocarbons, ethane and ethylene. Numerous investigations have followed this seminal work and a large number of catalysts have been found which give total selectivity to C 2 hydrocarbons higher than 90% at low (<2%) methane conversion [2-6]. However, it was generally found that the total C2 hydrocarbon selectivity decreases drastically with increasing conversion of methane, so that Yc 2 (the total C2 hydrocarbon yield) was always found, until very recently, to be less than 30% [1-9]. Achieving C2 hydrocarbon yield in excess of 50% is a necessary requirement for the development of an economically viable industrial process. The reason for the low C2 selectivity values at high methane conversion and thus the reason for the low measured Yc2 and YC2H4 yield values of earlier

388 studies is that the desired products, ethylene and ethane, are far more reactive with oxygen than methane and therefore are easily oxidized to CO/CO 2 when their concentrations become comparable to t h a t of methane, i.e. for high m e t h a n e conversion. Aris and coworkers [10] recently showed t h a t Yc 2 and YC2H4 can be increased up to 50% and 17%, respectively, by using a Sm20 3 catalyst in a simulated countercurrent moving bed chromatographic reactor (SCMBCR) to carry out the OCM reaction. The observed significant improvement in C 2 y~ld (-50%) in the case of the SCMBCR is due to the partial separation and removal of C 2 hydrocarbons from u n r e a c t e d m e t h a n e and oxygen. Despite the reactor complexity, which may not be easy to adapt to industrial practice [11], this pioneering work created a great deal of interest and has underlined the importance of rapid C2 removal from the CH 4 catalytic oxidation zone. We have recently presented preliminary results obtained with a novel catalytic or electrocatalytic OCM reactor-separator which almost entirely eliminates the problem of the high reactivity of the C 2 hydrocarbons during the OCM reaction [12]. In this concept a gas recycle catalytic or electrocatalytic reactor is combined with an appropriate molecular sieve trap in the recycle loop (Linde molecular sieve 5A maintained at <70~ which traps and thus protects an easily controllable percentage (up to 100%) of ethylene and of ethane produced during each gas cycle [12]. An important feature of this molecular sieve material is that it traps ethylene much more effectively than ethane and thus leads to very high ethylene yields, and ethylene to ethane ratios up to 30 for batch operation [12]. The process is simple and appears promising for industrial application [13]. In this work we present results obtained both with batch and continuous flow operation of the gas-recycle reactor-separator utilizing Ag and Ag-Sm20 3 electrocatalysts and Sr(lwt%) La20 3 catalysts, in conjunction with Linde molecular sieve 5A as the trapping material, and discuss the synergy between the catalytic and adsorption units in view of the OCM reaction network. 2. EXPERIMENTAL The recycle reactor is shown schematically in Figure 1. It consists of a catalytic or electrocatalytic reactor unit with a bypass loop, a recycle pump and a molecular sieve trap unit. The l a t t e r comprises one or two packed bed columns in parallel each containing 2-10 g of Linde 5A molecular sieve pellets. On line gas chromatography (Shimadzu 14A) was used for the analysis of CH 4, 0 2, CO, CO 2, C2H 4 and C2H 6 in the reactants and products. Two types of reactors were used: One was a CSTR type consisting of an Y203 (8mol%)-stabilized ZrO 2 (YSZ) tube (length 15 cm, diameter 2 cm) closed flat at one end with an appropriately machined water-cooled stainless steel reactor cap attached to the other end, thus allowing for continuous gas feed and

389 removal [9,14]. In this reactor the catalyst was a porous Ag film (mass 150 rag; superficial surface area 10cm 2) or a porous cermet of the same mass and surface area consisting of 20 wt% Sm20 3 (doped with 1% CaO) and 80% Ag. The catalyst was coated on the inside walls of the O2--conducting YSZ tube (Fig. 1). Catalyst film preparation and characterization details are given elsewhere [9,14]. The second reactor was a plug flow quartz reactor (length 5 cm, diameter lcm) with a fixed catalyst bed of Sr(lwt%)/La20 3 pellets (0.5-1mm diameter, total mass 0.5g). The YSZ reactor could be operated catalytically or electrocatalytically, depending on the mode of oxygen addition. Oxygen could be supplied either electrochemically by means of the solid electrolyte wall of the reactor (electrocatalytic operation) or in the gas phase (catalytic operation) (Fig. 1).

Figure 1. Schematic of the gas recycle electrocatalytic or catalytic reactor separator; WE, CE and RE are the working, counter and reference electrodes respectively; PCV: product collection vessel.

390 In the case of electrocatalytic operation, a g a l v a n o s t a t was used to apply constant currents I between the catalyst and a counter electrode deposited at the outer walls of the YSZ tube. In this way, oxygen is supplied to the Ag-based catalyst at a rate I/2F mol O/s, where F is Faraday's constant. In this case the catalyst acts as an electrocatalyst [9,12,14]. Appropriate setting of two on-off valves (Fig. 1) allows the system to be operated either as a batch recycle reactor or as a continuous-flow steady-state recycle reactor. In this work we present results obtained with the YSZ reactor operated in the batch mode with electrochemical oxygen addition, and with the quartz plug flow reactor operated in the continuous-flow steady-state mode. In the case of continuous flow operation, the molecular sieve trap comprised two packed bed units in parallel in a swing-bed a r r a n g e m e n t (Fig. 1), t h a t is, one unit was m a i n t a i n e d at low t e m p e r a t u r e (<70~ to c o n t i n u o u s l y t r a p the reactor products while the other was heated for --30 rain to 300~ to release the products in a slow stream of He. On-line GC analysis (Shimadzu GC 14A) was used to m e a s u r e product selectivity and m e t h a n e conversion. Details on the analysis procedure used for batch and continuous-flow operation are given elsewhere [12]. The molecular sieve trap was found to trap practically all ethylene, CO 2 and H20 produced; a significant, and controllable via the adsorbent mass, percentage of ethane; and practically no methane, oxygen or CO, for t e m p e r a t u r e s 50-70~ The trap was heated to -300~ in order to release all trapped products into the recirculating gas phase (in the case of batch operation), or in a slow He stream (in the case of continuous flow operation).

3. R E S U L T S AND DISCUSSION 3.1. Ag-based electrocatalysts, batch o p e r a t i o n Silver films and Ag-CaO-Sm20 3 c e r m e t s were chosen as the anodic e l e c t r o c a t a l y s t s because of their h i g h electrical conductivity, which is necessary for electrocatalytic operation, and also because of their high (>95%) selectivity to C 2 hydrocarbons at very low (<2%) CH 4 conversions [9]. Figs. 2a and 2b show typical results obtained with the electrocatalytic YSZ recycle reactor operated in the batch mode. The initial CH 4 partial p r e s s u r e was 20 kPa, i.e. 20% CH 4 in He at atmospheric pressure. The recirculation flowrate was 220 cm3/min. Oxygen was supplied electrochemically, at a rate I/2F, t h r o u g h the solid electrolyte walls of the YSZ tube. The applied current was always maintained at low levels, so t h a t the catalyst potential Vwc with respect to the Ag counter electrode was always m a i n t a i n e d negative. Higher c u r r e n t s l e a d i n g to V w c > 0 were found to lead to poor C 2 s e l e c t i v i t y . Consequently all results reported here can be considered as obtained in the chemical cogeneration mode [14], i.e., with s p o n t a n e o u s cell operation and

391 some electrical energy production. Figure 2 depicts the effect of CH 4 conversion and current on the total C 2 hydrocarbon selectivity and yield (Fig. 2a) and on the C2H 4 selectivity and yield (Fig. 2b). For any fixed m e t h a n e conversion, decreasing current, i.e. oxygen supply, causes a pronounced increase in C 2 selectivity and yield (Fig. 2a) and also:~n C2H 4 selectivity and yield (Fig. 2b). (b)

(a) I00

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.

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..... ~,::. ,.'. 4-0 ........... o . . . . .

" 00

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835~

.o

=

o

=

ox .

,~-o"" ."

I00

Cc~,,%

20 40 60 CH4 C o n v e r s i o n , Cc~,%

I 1

Figure 2. Effect of m e t h a n e conversion and applied c u r r e n t on the C 2 hydrocarbon (a) and on the ethylene (b) selectivity (filled symbols) and yield (open symbols). (Reprinted with permission from the AAAS, ref. 12). As shown in Fig. 2a for I=5mA and 97% CH 4 conversion, the C 2 selectivity is 91%, corresponding to a C 2 yield of 88%, which is the highest C 2 yield obtained so far for the OCM reaction. More importantly under these conditions 97% of the C 2 hydrocarbon products is ethylene (Fig. 2b), i.e. the ethylene selectivity is 88%, the ethylene yield is 85% and the ethylene to ethane ratio is 30 (Fig. 3). The effect of CH 4 conversion on the total C2, C2H4, C2H6 hydrocarbon selectivity and yield is shown in detail on Figure 4 for the case of I=5mA. Interestingly, the ethylene selectivity can increase with increasing methane conversion. This is because of the predominantly consecutive nature of the OCM reaction network: 02.

02-

02.

2CH4 --e 2CH3" --e C2H6 --+ C2H4 --+ 2CO2 1

2

3

(1)

4

The m o l e c u l a r sieve a d s o r b e n t t r a p s ethylene q u a n t i t a t i v e l y , thus practically freezing step 4. Ethane trapping is only partial, thus the desired step 3 is not decelerated significantly. Steps 1,3 and 4 are predominantly catalytic or electrocatalytic, depending on the mode of oxygen addition,

392

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PRODUC T5:

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:

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PRODUC TS

Figure 3. Gas chromatogram of initial and final gas composition for the maximum C~. yield point of Fig. 2a. (Reprinted with permission from the AAAS, ref. 12).

Figure 4. Effect of methane conversion for I=5 mA on ethylene, ethane and total C 2 hydrocarbon selectivity and yield. Lines from kinetic model discussed below. Solid lines" C2H 4 and C2H6; Dashed lines: C 2

whereas step 2 is predominantly homogeneous, as established elegantly by Lunsford and co-workers [3,5]. Thus, the observed dependence of Sc 2, 8C2H4, SC2H6, YC2, YC2H4 and YC2H6 on methane conversion shown in Fig. 4 can be rationalized easily on the basis of the above consecutive reaction network. In a separate set of experiments, oxygen was supplied from to the gas phase at a rate comparable to I/2F, instead of electrochemically. It was found that both modes of oxygen addition give practically the same results. This is due to the complete oxygen conversion, since in a recent study we have found that, at moderate oxygen conversion levels in a single pass reactor, electrochemically supplied oxygen can be more active and selective than gas phase oxygen [9]. Substituting Ag with the Sm203-CaO-Ag cermet or substituting YSZ with a-A120 3 as the cermet support gave practically the same results [15]. The ethylene selectivity (Fig. 5) and thus the ethylene yield depend strongly on the adsorbent mass (Fig. 5). For fixed catalyst mass, oxygen supply I/2F and methane conversion there is an optimal amount of adsorbent for maximizing ethylene selectivity and yield (Fig. 5). Excessive amounts of adsorbent cause quantitative trapping of ethane and thus a decrease in ethylene yield according to the above reaction network. This shows the important synergy between the catalytic and adsorbent units which significantly affects the product distribution and yield.

393 1oo

80

..,

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.--'*"

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.-"'" 8C2H6 T

0

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z 3 Adsorbent

Figure 5. Effect of adsorbent m a s s in the molecular sieve t r a p on the ethylene, ethane a n d total C 2 selectivity at a fixed m e t h a n e conversion of 15%. Recirculation flowrate 220 cm 3 STP/min

= 835"C

4 5 6 weight, g

7

3.2. LaeO3-based c a t a l y s t , continuous f l o w o p e r a t i o n Figure 6 shows typical results obtained with the plug-flow q u a r t z reactor containing 0.5 g of S r ( l w t % ) / L a 2 0 3 c a t a l y s t operated in the continuous flow recycle mode. The inlet CH 4 partial pressure was 20 k P a (20% CH 4 in He) at inlet flowrates of 7.1 and 14.3 cm 3 STP/min. A 20% O2 in He m i x t u r e was supplied directly, at a flowrate F02 , in the recycle loop via a needle valve placed after the reactor (Fig. 1). The m e t h a n e conversion was controlled by adjusting Fo 2, which was kept at appropriately low levels so t h a t the oxygen conversion was always higher t h a n 95%. In this way the oxygen partial p r e s s u r e in the recycle loop is very low and no explosive mixtures with CH 4 can form. 100~

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Figure 6. Continuous flow steady-state operation: (a) Effect of oxygen stream flowrate on C2 selectivity a n d yield; (b) corresponding effect of m e t h a n e conversion on the selectivity and yield of C2H4 and C2Hs; C a t a l y s t Sr/La203; T=750oC; recirculation flowrate 200 cm3/min.

394 Figure 6a shows the effect of Fo 2 on the C 2 selectivity and yield. The C 2 yield is up to 53%. Figure 6b refers to the same experiments and shows the corresponding effect of CH 4 conversion on the selectivity and yield of ethylene and ethane. The ethylene yield is up to 50% (65% ethylene selectivity at 76% methane conversion). To the best of our knowledge this is the maximum ethylene yield obtained for the OCM reaction under continuous-flow steadystate conditions. It is worth noting that for similar operating conditions, batch operation gives higher selectivity and yield values than continuous-flow steady-state operation. This is because in the latter case the methane concentration in the recycle loop is lower than the average methane concentration during each batch [12]. High oxygen to methane ratios favor CO 2 formation. The improvement in C2 selectivity and yield of the present work, for both batch and continuous flow steady-state operation, is not attributable to the properties of the catalyst used. This is shown clearly in Fig. 7 which compares some of the highest literature C2 selectivity and yield values on the basis of Figure 8 of the review article of Lee and Oyama [4], the pioneering work of Aris and coworkers [10], the present results using the novel gas recycle reactor in the continuous flow (catalytic) and batch (electrocatalytic) mode and the results obtained under the same operating conditions but without gas recycling and molecular sieve trapping (single pass conditions). It is clear that the very significant improvement in C2 selectivity and yield of the present work, for both batch and continuous flow steady-state operation, is a result of the reactor design. Thus, as shown in Fig. 7, the Sr-doped-La20 3 catalyst is comparable to some of the best reported OCM catalysts particularly at low CH4 conversions and gives single-pass C 2 yields up to 13% but the Ag or Ag-Sm20 3 catalysts or electrocatalysts give single pass C2 yields of less than 4%. It is also evident I00

80,

~ 60

k

0

L ~A ~

9

~

~

40

"" ref 4

I N

r~

20

20 40 6O 8O CH4 Conversion, Cc~,,~o

100

Figure 7. Effect of methane conversion o n C 2 selectivity for some of the best state-oft h e - a r t OCM catalysts (A, based on ref. 4), the simulated chromatographic reactor of Aris and coworkers (A, ref. 10) and the present work. (m): Ag e l e c t r o c a t a l y s t , single pass; (O): Ag electrocatalyst with recycle and trapping; (0) Sr/La203 catalyst, single pass; (O) Sr/La20 3 catalyst with recycle and trapping. Open symbols, b a t c h operation; filled symbols, continuousflow steady-state operation.

395 from Figure 7 t h a t batch or periodic batch (unsteady-state) recycle operation gives higher performance t h a n steady-state reactor operation, as discussed previously. 3.3. Elec~m~talytic a n d catalytic reactor modp.llng A simplified mathematical model was developed for the novel OCM reactor. One version of the model, presented here, describes batch operation. A second version addressing continuous flow operation will appear elsewhere [16]. The model assumes a well-mixed gas phase composition in the recycle loop, a well justified assumption in view of the very high (10-200) recycle ratio values used in the present work. For the batch electrocatalytic version we also neglect volume changes and assume linear kinetics for steps 1,3 and 4 of the consecutive OCM network (1), i.e.: rl=r2=klAYcH4; r3=k2AYc2H6 ; r4=k3AYc2H4

(2)

where A is the catalyst-electrode surface area. We also assume complete oxygen conversion and linear kinetics for the rates, rt,C2H6 and rt,C2H4 , of molecular sieve trapping of C2H 4 and C2H 6, i.e.: rt, C2H6=kt,2AtYc2H6 ; rt, C2H4=kt,3AtYc2H4

(3)

where A t is the molecular sieve surface area. We thus obtain the following simplified dimensionless mass balance equations: dCcH4/dx = -CcH4

(4a)

dCc2H6/dl; = (1/2) CCH4 - ~-2Cc2H6 - kt,2Cc2H6

(4b)

dCc2H4/d~; = ~2Cc2H6- ~-3Cc2H4- ~t,3Cc2H4

(4c)

where the symbols are defined in Table 1. The ethylene yield is computed from: YC2H4 = 2 CC2H4 + Io~t,3CC2H4d~;'

(5)

The set of equations 4a, 4b and 4c was solved analytically with initial o

conditions CCH4=I, (i.e. YC2H4/YCH4) CC2H4= CC2H6--0 and the resulting CCH4(t), CC2H4(t) and CC2H6(t), which contain the parameters k2, ~3' kt,2 and ~t,3' were used to fit the data of Fig. 4. As shown in this Figure the electrocatalytic batch model provides a quantitative fit to the experimental data.

396 Table 1 Model parameters o

/o

o

CCH4 = YCH4/YCH4 ; CC2H6= YC2H6YCH4 ; CC2H4 = YC2H4/YCH4 [CH4]~ Initial CH 4 concentration, mourn 3 ; t" Batch operating time, s V: Reactor-separator loop volume, m 3

Greek symbols ~L2=k2/kl; ~3=k3/kl; ~-t.2=kt~2At/klA; ~t.3=k~3At/klA;

't=klA~[CH4] ~

4. CONCLUSIONS Methane can be oxidatively coupled to ethylene with very high yield using the novel gas recycle electrocatalytic or catalytic reactor separator. The ethylene yield is up to 85% for batch operation and up to 50% for continuous flow operation. These promising results, which stem from the novel reactor design and from the adsorptive properties of the molecular sieve material, can be rationalized in terms of a simple macroscopic kinetic model. Such simplified models may be useful for scale up purposes. For practical applications it would be desirable to reduce the recycle ratio p to lower values (e.g. 5-8). This requires a single-pass C 2 yield of the order of 15-20%. The Sr-doped La20 3 catalyst of the present work is close to this target. Further reduction in p can be achieved via catalyst and molecular sieve adsorbent optimization. REFERENCI~ 1. 2. 3. 4. 5. 6. 8. 7. 9. 10. 11. 12. 13.

G.E. Keller and M.M. Bhasin, J. Catal. 73(1982) 9 W. Hinsen and M. Baerns, Chem. Ztg. 107 (1983) 223 T. Ito, J.H. Lunsford, Nature 314 (1985) 721 J.S. Lee and S.T. Oyama, Catal. Rev. -Sci. Eng. 30 (1988) 249 J.H. Lunsford, Catal. Today 6 (1990) 235 D. Eng and M. Stoukides, Catal. Rev. -Sci. Eng. 33 (1991) 375 J.M. Fox, Catal. Rev.-Sci. Eng., 35(2) (1993) 169 G. Renesme, J. Saint-Just Muller, Catal. Today, 13 (1992) 371 P. Tsiakaras and C.G. Vayenas, J. Catal. 144 (1993) 333 A.L. Tonkovich, R.W. Carr, R. Aris, Science 262 (1993) 221 J. Haggin, C&E News, pp. 4-5, October 11, (1993) Y. Jiang, I.V. Yentekakis and C.G. Vayenas, Science 264(1994)1563 Science 264, 1513, June 10, 1994; Chemistry and Industry 12, June 20, 1994; C&En News p. 41, June 13, 1994 14. C.G. Vayenas, S. Bebelis, I.Yentekakis, H.-G. Lintz, Cat. Today 11(1992)303 15. Y. Jiang, PhD Thesis, University of Patras (1994) 16. M. Makri, Y. Jiang, I.V. Yentekakis, C.G. Vayenas, in preparation (1995)

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All fights reserved.

397

T h e A c t i v e O x y g e n f o r D i r e c t O x i d a t i o n o f M e t h a n e to M e t h a n o l in the P r e s e n c e o f H y d r o g e n Ye Wang and Kiyoshi Otsuka Department of Chemical Engineering, Tokyo Institute of Technology, Ookayama, Meguro-ku, Tokyo 152, Japan Abstract

Methane was selectively converted to methanol by a mixture of oxygen and hydrogen at >573 K and atmospheric pressure over Fe05AIo.sPO4 catalyst, while, in the absence of H 2, the conversion of methane was notably low and only a trace of methanol was formed. Characterization of the catalyst by XPS before and after the reaction suggested the redox of Fe(IlI)/Fe(lI) on the catalyst surface during the oxidation of methane by a H~_-O2 gas mixture. The adsorbed oxygen species generated on the catalyst in a H~-O 2 gas mixture and its reactivity with methane were studied by in situ FT-IR spectroscopy. The absorption band at 895 cm 1 was observed in a H2-O 2 gas mixture when the temperature was raised above 573 K. The isotope substitution of 1602 with 1802 shifted the absorption band at 895 cm 1 to 849 cm 1. Three absorption bands at 895, 870 and 849 cm 1 were observed when the mixture of 1602, 16OXsOand 1802 was used with H 2 at >573 K. These observations strongly suggest that the band at 895 cm -1 is ascribed to a peroxide species adsorbed on the iron site of the catalyst. The intensity of the band due to this oxygen species decreased greatly in the presence of methane at :-473 K as a result of the reaction with methane. The new bands at 2936, 2870 and 1050 cm 1, which were assigned to the stretching vibrations of CH 3 groups and of C-O of methoxide species, appeared simultaneously. The absorption band at 3668 cm 1, the stretching vibration of the adsorbed OH group, increased at the same time. The absorption bands at 3668, 2936 and 2870 cm 1 were shifted to 2668, 2208 and 2124 cm 1, respectively, when CD 4 was used instead of CH 4. These results suggest that the adsorbed peroxide reacts with methane at >473 K, producing methoxide and OH groups as intermediates for the formation of methanol. 1. Introduction

Selective oxidation of methane, the most abundant and the least reactive alkane, to useful oxygenates such as methanol and formaldehyde would be of great value compared to the current process of methane to methanol through synthesis gas [1]. However, the current state is far from industrial realization. Although it has been reported that formaldehyde could be produced from the oxidation of methane with oxygen over solid catalysts, it seems very difficult to obtain methanol at low pressures (-:101 kPa) [2-6]. On the other hand, it is well known that methane monooxygenase (MMO)catalyzes the oxidation of methane to methanol by oxygen under mild conditions [7]. The accumulated knowledge for the MMO catalytic cycle suggests that oxygen is reductively activated on the iron centers by electrons and protons supplied from a reductant such

398 as NADH or NADPH, generating an active oxygen species, probably a high oxidation state iron oxo species, which directly converts methane to methanol [8-11]. In the light of such studies, we expect that if a reductant is cored with oxygen in the reaction the heterogeneous oxidation of methane to methanol using solid catalyst may also be realized. In this case, a cheap and easy handling gaseous reductant is most desirable. Recently, using hydrogen as an activator of oxygen, we have succeeded in converting methane selectively to methanol at atmospheric pressure over iron phosphate catalysts [12, 13]. The previous studies have suggested that a new oxygen species generated on FePO 4 surface in H2-O2 gas mixture is responsible for this oxidation. In this study, the active oxygen species generated on a similar catalyst to FePO 4 and the role of the active oxygen species in the selective oxidation of methane to methanol are studied in detail using in situ FTIR spectroscopy.

2. Experimental section 2.1 Catalyst The iron aluminum phosphate catalyst used in this study was prepared by a sol-gel method. FeCI 3, AICI3 and NH4H2PO 4 were used as the raw materials and were dissolved into water (30 ml). The mole ratio of Fe:AI:P was adjusted to be 0.50:0.50:1.00 in the solution. Propylene oxide (-40 ml) was added to the solution at 273 K. After 2h, the solution gelled and the gel was aged for 20 h at room temperature. The gel was dried at 333 and 393 K for 20 h after being thoroughly washed with ethanol. The resultant was calcined at 573 K for 20 h and at 823 K for 10 h in a flow of oxygen. The catalyst thus prepared is denoted as Fe05Alo 5PO 4, hereafter. The BET surface area of this catalyst was 275 m 2 g-X. The high specific surface area was necessary for gaining an adequate absorbance of IR of the adsorbed species.

2.1. Characterizations of the catalyst and evaluation of the catalytic activities The catalyst prepared above was characterized by X-ray diffraction, X-ray photoelectron and Mrssbauer spectroscopic studies. The catalytic activities were evaluated under atmospheric pressure using a conventional gas-flow system with a fixed-bed quartz reactor. The details of the reaction procedure were described elsewhere [13]. The reaction products were analyzed by an on-line gas chromatography. The mass balances for oxygen and carbon between the reactants and the products were checked and both were better than 95 %. 2 . 3 . In situ F T I R measurements The catalyst for the in situ FTIR-transmission measurements was pressed into a selfsupporting wafer (diameter 3 cm, weight 10 mg). The wafer was placed at the center of the quartz-made IR cell which was equipped with two NaCI windows. The NaCI windows were cooled with water flow, thus the catalyst could be heated to 1000 K in the cell. A thermocouple was set close to the sample wafer to detect the temperature of the catalyst. The cell was connected to a closed-gas-circulation system which was linked to a vacuum line. The gases used for adsorption and reaction experiments were O~ (99.95%), ~80~ (isotope purity, 97.5%), H 2 (99.999%), CH 4 (99.99%) and CD4 (isotope purity, 99.9%). For the reaction, the gases were circulated by a circulation pump and the products were removed by using an appropriate cold trap (e.g. dry-ice ethanol trap). The IR measurements were carried out with a JASCO FT/IR-7000 spectrometer. Most of the spectra were recorded with 4 cm x resolution and 50 scans.

399

3. Results and Discussion 3.1. Catalytic activity of the Feo.sAi0.sPO 4 catalyst for the oxidation of

CH 4

Figs. 1 (A) and (B) show the conversion of CH 4 and the product selectivities obtained in the presence and absence of H a over the Fe0.sAlo 5PO 4 catalyst. It is apparent from Fig. 1 (A) that the CH 4 conversion increases markedly (e.g. 6.7 times at 723 K) when H 2 is cofed with O2. In the absence of H 2, the conversion of CH4 does not occur at <673 K, while the reaction proceeds at 573 K in the presence of H 2. The apparent activation energies for CH 4 conversion calculated from Fig. I(A) are 147+_.10 and 220___10 kJ mol a in the presence and absence of H z, respectively. These results strongly suggest that the cofeeding of H,_ with 02 generates a new active center for the conversion of CH 4 over Fe0.sAlo.sPO4 catalyst. The other important effect of the cofeeding of H 2 is that CH 3OH is selectively produced. As shown in Fig. 1 (B), only a trace amount of C H 3OH is obtained in the absence of H2, while CH 3OH becomes a main product in the presence of H a. These results further suggest that a new active species selective for the oxidation of CH4 to CH 3OH is generated by the cofeeding of H~_ and 02 over the catalyst.

3

100

(B)

(A) .c

2

ra~

"~ 50

> r

- ~

,/~

.

r 0 550

0 600

650 T/K

700

750

550

600

650 T/K

700

750

Fig. 1 Oxidation of CH4 in the presence and absence of H2 over Fe0.sAl0.sPO4 catalyst. (A) CH4 conversion; (O) in the absence of H2, (Q) in the presence of 50.7 kPa Hz; (B) Product selectivities; black symbols: in the presence of 50.7 kPa He; white symbols: in the absence of H2. ( S ) and (C)) CH3OH, (O) and ( ~ ) HCHO, (A) and (/X) CO, ( I ) and (7-1) CO2. Conditions: P(CH4) = 33.8 kPa, P(O2)= 8.4 kPa, W/F= 0.028 g h dm-3

The AIPO 4 prepared by a similar method to that used for the preparation of the Fe0.sAlo sPO 4 catalyst was also tested for the oxidation of CH 4 by 02 both in the presence and absence of H 2. No reaction occurred at 573-723K, suggesting that the iron was an indispensable active component for the conversion of C H 4. AS compared with the FePO 4 catalyst reported previously [13], the space time yield for the sum of CH3OH and HCHO (per unit gram)at 723 K increased for about 10 times for the Fe05Alo 5PO4 catalyst, i.e., from 1.36 mmol h 1 g-I for FePO 4 to 14.1 mmol h a g.1. This must be ascribed to the high specific surface area of the catalyst used in this study.

400 3 . 2 . Characterization of the Feo.sAio.sPO4 catalyst The results of the XRD measurement showed that the Fe0.sA10.sPO4 catalyst was "~most in amorphous state. Only a very broad peak at 20 of ca. 23 degree was observed. The M6ssbauer spectroscopic study on this catalyst showed one doublet of iron with the isomeric shift of 0.31 mm s a (ot-Fe was used as the reference) and the quadrupole splitting of 0.62 man s 1. These parameters are very close to those observed for FePO 4 [13, 14], suggesting that the iron cation in the catalyst is tetrahedrally coordinated with oxygen and isolated by four PO4 tetrahedral units. Such coordination circumstance was suggested to be a key factor for the iron site effective for the oxidation of CH4 to CH3OH by H2-O 2 gas mixture [15]. After the reaction for 5 h in a reactant stream of CH,, O z and H 2 (P(CH4)= 33.7, P(O2)= 8.4 and P(H2)= 50.7 kPa), the catalyst was analyzed 200 by XRD, M6ssbauer and XPS studies. As regarding the XRD and M6ssbauer spectroscopic measurements, obvious changes were not observed before and after the reaction. On the other hand, a marked change was observed in the XPS spectrum of the catalyst after the reaction. As r shown in Fig. 2, besides the peak at 57.7 eV, N which was the only peak of Fe3p obtained for the sample before the reaction and was ascribed to Fe(III), a clear shoulder at 56.1 eV was observed after the reaction. This can be ascribed to the ! I I l Fe(lI) on the catalyst surface. The same 64 60 56 52 phenomenon has been reported for FePO4 catalyst [13]. Such observations suggest the occurrence of Binding Energy/eV the redox of iron between Fe(llI) and Fe(ll) during Fig. 2 XPS spectra for Fe 3p. the reaction. We believe that this redox plays a (a) before the reaction, (b) after the key role in the formation of a new active center and reaction in a reactant stream of CH4, thus is important in the selective oxidation of CH4 H2 and 02. to CH3OH as will be discussed later.

I

3 . 3 . A d s o r b e d s p e c i e s on the catalyst in H2-O 2 gas mixture In situ FTIR spectroscopy was used to study the adsorbed species generated on the catalyst surface in the presence of H 2 and 02. Before the experiment, the catalyst wafer was pretreated by 02 (5.3 kPa) at 723 K for 1 h followed by evacuation at the same temperature in vacuum (ca. 6x10 3 Pa) for 2 h. After the pretreatment, the temperature was decreased to a desired one in vacuum and IR spectrum was recorded at that temperature. The spectra of the catalyst wafer recorded at different temperatures were used as the background ones for the adsorption studies described below. Fig. 3 shows the IR spectra of the adsorbed species generated on the surface in the presence of both H 2 and 02 at different temperatures. No obvious absorption band due to the adsorbed species was observed at 298 K. When the temperature was increased to 473 K, two weak bands at 3740 and 3670 cm 1 assigned to the stretching vibrations of a non-acidic and acidic OH groups, respectively, were observed. These two bands were also observed in H 2

401 alone at >473 K. Moreover, the isotope substitution of 1602 with ISoz in the H2-O 2 mixture did not affect the position of these OH stretching bands. Therefore, the OH groups must be formed due to the interaction of H z with the lattice oxygen of the catalyst surface. Particularly, the acidic OH at 3670 cm 1 may be generated as a result of the reduction of the surface Fe(IlI) according to the following reaction, -Fe(IlI)- O-P-O- + 1/2H z ---, -Fe(II) + HO-P-O(1) Besides the absorption bands attributed to OH groups, another weak band at 1075 cm 1 was observed at 473 K. This band can tentatively be assigned to an 02- species. The formation of 02- species with an absorption band at 1050-1200 cm a was reported for many oxides [1618]. Here, the O.,- species must be formed on the reduced iron site through one electron transfer from Fe(II) to the adsorbed Oz. However, this 02- species was very unstable because it disappeared when the temperature was increased above 473 K, probably due to the desorption or to the transformation to u% another type of oxygen species. As shown in Fig. 3, when the 0.01 O.Ol temperature was increased to 573 K, an absorption band at 895 crn 1 appeared and the bands of OH groups became stronger. The r162 intensity of the absorption band at 895 cm ~ 673K increased with a rise in temperature to 623 and 673 K. The stability of the species giving the absorption band at 895 cm 1 was examined under vacuum at different 623K temperatures. Only a slight decrease in the intensity was observed when the system was degassed at <473 K. However, a decrease 573K was observed when the temperature was raised above 473 K. The band completely H3K disappeared at >623 K. These results indicate that the absorption at 895 cm 1 was ascribed to a chemisorbed species peculiarly 298K generated and stable in the presence of H 2 and O2. It should be noted that the band at 895 cm 1 has never been observed in the | ! presence of O., alone. 1400 800 ;00 3 4 0 0 3 0 0 0 2 6 0 0 In order to identify the new adsorbed Wavenumber/cmq species of the band at 895 cm a, the effect of the isotope substitution of 160z with 1sOz Fig. 3 IR spectra of adsorbed species on was examined. The spectra shown in Fig. 4 the catalyst in H2-O2 at different were recorded for the catalyst in the presence temperatures. of H 2 and aSO2 in the wavenumber region of P(H2)= 13.3 kPa, P(O2)- 1.33 kPa. 700-950 cm -1 with 100 scans. The results show that the absorption band at 849 cm 1

I

402 appears when the temperature is higher than 573 K. This means that the absorption band at 895 cm -1 is shifted to 849 cm 1 when 1602 is replaced by 180,_. This result strongly suggests that the absorption at 895 cm 1 is ascribed to an adsorbed oxygen species on the catalyst surface. However, the results obtained above do not allow us to distinguish whether this adsorbed oxygen species is a diatomic type oxygen such as peroxide species (022) or a monoatomic type oxygen such as iron oxo species (Fe=O). Generally, both the adsorbed peroxide and M=O type oxygen species (where M is a metal) give IR absorption band at 800-950 cm 1 [ 17-21]. We can calculate the isotope shift in IR absorption due to the substitution of 1602 with 1802 by assuming the adsorbed species to be an O, 2 or a Fe=O species. If the absorption is ascribed to the O-O stretching mode of a peroxide species, the ratio of the wavenumbe~ for the 180-180 and 160-160 stretching modes should roughly be given by the ratio of the square root of the reduced masses, viz. [(1/18 + 1/18)/(1/16 + 1/16)] 1/2 or 0.943. Thus, the band at 895 cm -1 should be shifted to 844 cm 1 by the substitution of 1602 with 1802 in this case. If the adsorbed species is of Fe=O type, a similar calculation suggests that the substitution of 1602 with 1802 will shift the peak at 895 crn ~ to 855 cm -1. As described above, the experimental results showed that the band at 895 cm 1 was shifted to 849 cm 1 due to the replacement of 1602 with 1802 in the H2-O 2 gas

[0.005

473K

i 900

i 800

700

Wavenumber/cm-I

Fig. 4 IR spectra of the adsorbed species on the catalyst in Hz-18(b_ at different temperatures. P(H2)= 9.3 kPa, P(1802)= 0.93 kPa.

1 900

I 800

700

Wavenumber/cm-1

Fig. 5 IR spectra of the adsorbed species on the catalyst in the presence of H2 and a mixture of 1602, 16Oa80 and 1802_ at different temperatures.

403

mixture. Therefore, we still cannot identify the oxygen species, viz. peroxide or iron-oxo species? Thus, in this context, further experiments using an isotopic mixture of 1602, ~60180 and 180z were carried out. IR spectra were recorded under the mixture of H a, 1602, 160~80 and 180, at different temperatures. The initial pressures of H e and O2 at 298 K were 8.7 and 0.93 kPa, respectively. The analytical result of the composition of the isotopic oxygen by a quadrupole mass spectrometer was 2.0:0.5:1.0 for the 1602"160180:1802. The IR spectra obtained in the region of 950-700 crn 1 were shown in Fig. 5. Three absorption bands at 895, 870 and 849 crn a were observed at 573 K and their intensities were increased further with the increase in temperature to 623 K. These observations can only be explained by assuming that these absorption bands are ascribed to the O-O stretching mode of the adsorbed peroxide. The calculation shows that if the band at 895 cm 1 is ascribed to the stretching of 160-160, the wavenumbers for the stretching vibrations of 160-180 and ~80-180 should shift to 870 and 844 cm 1. Therefore, the observed bands at 895, 870 and 849 cm 1 can be assigned to the O-O stretching vibrations of 160-160, 160-180 and 180-180, respectively. In conclusion, an adsorbed peroxide species is formed on the catalyst surface in the presence of H,_ and O,. at e573 K.

3 . 4 . The reaction of the adsorbed oxygen species with methane As described above, the temperature needed for the formation of the adsorbed peroxide was ca. 573 K, which was almost the same as that for the initiation of the oxidation of CH 4 to CH3OH in the presence of H, and O2. This peroxide may directly be related to the selective oxidation of CH 4 to CH3OH. Therefore, the reactivity of this adsorbed peroxide with CH4 was investigated by in situ FTIR spectroscopy as follows. The adsorbed peroxide species for the reaction was prepared in a gas mixture of H a and O~ (P(H2)= 13.3 kPa, P(O2)= 1.33 kPa)at 673 K. Then, the temperature was rapidly lowered to 373 K and the gas phase was evacuated at this temperature. The intensity of the band at 895 cm ~ attributed to the peroxide was not affected by evacuation at this temperature. After the evacuation, CH4 gas (5.33 kPa) was introduced to the cell. IR spectra were recorded at different temperatures in the atmosphere of CH 4. Fig. 6 shows the IR spectra where the absorption due to the gaseous CH 4 at the corresponding temperatures has been subtracted. No obvious change in the intensity of the band at 895 cm x occurred after CH4 was introduced at 373 K. When the temperature was increased to 473 K, the intensity of the band at 895 cm 1 decreased markedly and three new bands at 2936, 2870 and 1050 cm -1 appeared simultaneously. The former two bands can be assigned to the anti-symmetric and symmetric stretching vibrations of C H 3 groups and the latter to the C-O stretching vibration for the methoxide species [22]. The reverse band at ca. 3010 cm ~ arose due to the decrease in the amount of the gaseous CH4 because the gaseous CH,, shows a very strong absorption band at 3010 cm ~. Accompanying the formation of the methoxide, the intensity of the band at 3668 cm I, assigned to the acidic OH group on the catalyst, increased obviously. In order to confirm that the absorption bands appeared at 3668, 2936 and 2870 cm -~ are derived from the reaction of CH4 with the adsorbed peroxide, CD 4 was used instead of CH 4. The IR spectra recorded in the regions of 4000--2000 cm 1 and of 1500-700 cm ~ are shown in Fig. 7. The experimental procedures for the results of Fig. 7 were the ~ m e as those for Fig. 6. When the temperature was increased to 473 K in the presence of CD4, the bands at 895 cm 1, ascribed to the adsorbed peroxide, decreased remarkably. The new bands at 2692, 2208 and 2124 cm 1 were observed simultaneously. Obviously, these are the isotopic bands of those

404 of 3668, 2936 and 2870 cm 1 observed in the case of C H 4 a s the reactant in Fig. 6. The ratios of wavenumbers for the corresponding pairs of the bands were 1.375, 1.330 and 1.351, respectively. All these values are reasonable ones for O-H/O-D, C-H/C-D (symmetric) and CH/C-D (anti-symmetric)stretching vibration modes. In conclusion, the adsorbed iron peroxide generated in the gas mixture of H 2 and 02 reacts with C H 4 at >473 K, forming adsorbed OH and methoxide species on the catalyst surface. Both species must be the intermediates in the formation of CH3OH.

I e~

I

O.Ol

,~,~

I

O.O1

II

tt%

' a2

I

O.Ol

I

0.01

!

J

r~l r~l .--I ~ r {'xll

~I~

573K ~ 573K

o

< 173K

473K

;73K 173K !

4000 3800

3400

3000

2600

1400

800

I

I

3000

2000

1400

1

I

801

Wavenumber/cmq

Wavenumber/cm-I

Fig. 6 IR spectra of the adsorbed species on the catalyst after the reaction of CH4 with the oxygen species generated in a H2-O2 gas mixture at different temperatures.

Fig. 7 IR spectra of the adsorbed species on the catalyst after the reaction of CD4 with the oxygen species generated in a H2-O2 gas mixture at different temperatures.

The effects of the partial pressures of H 2 and 02 on the formation of the adsorbed peroxide species were examined. These results have been compared with the kinetic results for the conversion of CH, by using the flow system. As shown in Fig. 8 (A), the surface concentration of the peroxide increased roughly linearly with a rise in the partial pressure of H~.. On the other hand, it was saturated at a low partial pressure of O2 (Fig. 8 (B)). Very similar trends were observed for the kinetic measurements for the conversion rate of CH 4 as functions of the partial pressures of H2 and 02 as shown in Fig. 9. These observations further support that the peroxide species is responsible for the partial oxidation of CH4.

405

~/~

0.03 _ (A)

0.03

0.02

0.02

.= < 0.01

0.01

t~ t.... O

I 1 1 0 10 15 20 0 1 2 3 P(H2)/kPa P(O2)/kPa Fig. 8 The intensity of the band at 895 cm -1 vs. the pressures of H2 (A) and 02 (B). T= 673 K; (A) P(O2)= 1.3 kPa; (B) P(H2)= 13.3 kPa.

5

"7, .m

_~ 0.6

(B)

(A) 0.4-

o4

o= "7.~ 0.2 w

= rj

~

0_

0

O.2-

658 K

20

40 P(H2)/kPa

60

00

1

1

1

1

1

2

3

4

5

P(O2)/kPa

Fig. 9 The conversion rate of CH4 vs. the pressures of H2 (A) and 02 (B). (A) P(O2)= 8.4 kPa; (B) P(H2)= 50.7 kPa.

3.5.

Reaction

mechanism

As described earlier, it is suggested that H 2 reduces the iron sites on the catalyst surface, forming water and OH groups. Thus, the cofeeding of H 2 must guarantee a high concentration of the active iron site, viz. Fe(II), as indicated by the XPS studies. Oxygen must be adsorbed on this reduced iron site by accepting electrons, initially forming O~.- species (absorption band at 1075 cm 1) as demonstrated in Fig. 3. This oxygen species itself may be very unstable because the band at 1075 cm 1 disappeared at >473 K. It was probably further reduced into 02- species by electrons trapped in Fe(ll) sites. The in situ FTIR results have shown that an adsorbed peroxide is generated on the catalyst surface in a H2-O: gas mixture at a573 K. Although we cannot exclude the possibility that the true active oxygen species could be an atomic oxygen intermediate (O-) resulted from the splitting of the adsorbed peroxide, the results of Figs. 6-9 strongly suggest that the peroxide species itself is responsible for the activation of CH 4. The results obtained in this study suggest the reaction mechanism in Fig. 10. As can be seen in Fig. 10, besides the redoxing iron sites, the proton accepting and donating sites, viz. phosphate groups, are also important for the generation of active oxygen species as well as for the accomplishment of the catalytic cycle. We believe that the phosphate groups directly bonded to the iron site play these specific roles [15].

406

CH3OHy

O~ . 0 ~ i )...,.Fe(II
oI < H 2

CH3

o

H+

o.

o..

H §

o ..o o'r<(H)'o~ P~ o"F~~ o

. H 2 0 ~ ~ '-O2 c,,3 o.,,o.--

t

H F~(n~ o \ p / O~ oi

\

o J<~

O

O22-

~ 0 o

O,1

o P~ . . oo/Fe(ll~,,,O .

(67b)2_

CH 4

" "~'O ~P\, 0., oY~(In ~ Fe(lll) 0o

Fig. 10 Proposed reaction mechanism.

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

J.M. Fox, Catal. Rev. -Sci. Eng., 35 (1993) 169. R. Pitchai and K. Klier, Catal. Rev. -Sci. Eng., 28 (1986) 13. M.J. Brown and N.D. Parkyns, Catal. Today, 8 (1991) 305. N.D. Parkyns, C.I. Warburton and J.D. Wilson, Catal. Today, 18 (1993) 385. O.V. Krylov, Catal. Toady 18 (1993) 209. T.J. Hall, J.S.J. Hargreaves, G.J. Hutchings, R.W. Joyner and S.H. Taylor, Fuel. Proc. Tech., 42 (1995) 151. 7. J. Colby, D.I. Stirling and H. Dalton, Biochem. J., 165 (1977) 395. 8. H. Dalton and J. Green, J. Biol. Chem., 264 (1989) 17698. 9. B.G. Fox, J.G. Bomeman, L.P. and J.D. Lipscomb, Biochemistry, 29 (1990) 6419. 10. S.-K. Lee, J. C. Nesheim and J. D. Lipscomb, J. Biol. Chem. 268 (1993) 21569. 11. M.J. Rataj, J.E. Kauth and M.I. Donnelly, J. Biol. Chem., 166 (1991) 18684. 12. Y. Wang and K. Otsuka, J. Chem. Soc. Chem. Commun., (1994) 1893. 13. Y. Wang and K. Otsuka, J. Catal., 155 (1995) 256. 14. J. M. Millet, C. Virely, M. Forissier, P. Bussi6re and J.C. Vedrine, Hyperfine Interact., 46 (1989) 619. 15. Y. Wang and K. Otsuka, Submitted for publication. 16. C. Li, K. Domen, K. Maruya and T. Onishi, J. Am. Chem. Soc., 111 (1989) 7683. 17. K.H. Jacob, E.Kn6zinger and S. Benfer, J. Chem. Soc. Faraday Trans.,90 (1994) 2969. 18. F.AI-Mashta, V. Lorenzelli and G. Busca, J. Chem. Soc. Faraday Trans. I, 78 (1982) 979. 19. M. Nakamura, S. Fujita and N. Takezawa, Catal. Lett., 14 (1992) 315. 20. P.J.M. Carrott and N. Sheppard, J. Chem. Soc., Faraday Trans. I, 79 (1983) 2425. 21. D. Klissurski, K. Hadjiivanov and A. Davydov, J. Catal., 111 (1988) 421. 22. T.P. Beebe, J.E. Crowell, J.T. Yates, J. Phys. Chem., 92 (1988) 1296.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

407

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

Photocatalytic Production o f Methanol and Hydrogen fi'om Methane and Water Charles E. Taylor, Richard P. Noceti*,

Joseph R.

D'Este, and Donald V. Martello

U.S. Department of Energy, Pittsburgh Energy Technology Center, P.O. Box 10940, Pittsburgh, PA 15236-(N40 (USA) Phone: (412) 892-5955, FAX: (412) 892-4152 *To whom correspondence should be sent ABSTRACt Investigation of direct conversion of methane to a'ms~rtation fuels has been an ongoing effort at PETC for over 10 years. One of our cmrent areas of research is the conversion of methane to methanol, under mild conditions, using light, water, and a semiconductor p h o t ~ y s t . Research in our laboratory is directed toward adapting the chemistry developed for photolysis of water to that of methane conversion. The reaction sequence of interest uses visible light, a doped tungsten oxide photocatalyst and an electron transfer molecule to produce a hydroxyl radic~. Hydroxyl radical can then react with a methane molecule to produce a methyl radical. In the preferred reaction pathway, the methyl radical then reacts with an additional water molecule to produce methanol and hydrogen. 1. INIRODUCHON Methane is produced as a by-product of coal gasification, either in a stand-alone process or as part of the direct or indirect liquefaction of coal. Depending upon the gasifier design and operating conditions, up to 18% of the total gaseous product may be methane. In addition, there are vast proven reserves of geologic methane in the world. A global research effort is under way in academia, industry, and govemment to find methods to convert methane to useful, more readily tramlx~rtable and storable materials. Methanol, the initial product of methane oxidation, is a desirable product of conversion because it retains much of the original energy of the methane while satisfying tmnspor~ion and storage requirements. A liquid at room temperature, methanol could be transported to market using the existing petroleum pipeline and tanker network and distribution infrastrucaz~. Methanol may be used directly as a fuel or may be converted to other valuable products (i.e., other wamportation fuels, fuel additives, or chemicals). At present, the direct oxidation of methane to methanol suffers from low methane conversion and poor methanol selectivity. A process for the direct oxidation of methane to methanol, with high yield and with high selectivity, is desirable. A long-term goal of our research group is to explore and evaluate novel pathways for the direct conversion of methane to liquid fuels, chemicals, and intermediates. One of our current areas of research is the conversion of methane to methanol, under mild conditions, using light, water, and a semiconductor photocatalyst, qhe use of three relatively abundant

408 and i n e ~ i v e reactants - light, water, and methane - to produce methanol is an attractive process option. The products of the reaction of interest, methanol and hydrogen, are both commercially desirable as fuels or chemical intermediates. Research in our laboratory is directed toward applying the techniques developed for the catalytic photolysis of water [1,2] and the photochemical conversion of methane to methanol [3,4]. The main advantage of using a photocatalyst to promote the photoconversion of methane to methanol is that the presence of the catalyst, in conjunction with an electron transfer agent, allows reaction to occur with visible light instead of with ultraviolet. This greatly simplifies reactor design and will permit flexibility in the selection of the light source. 2. BACKGROUND

It has been reported [1,2] that methane may be photochemically converted to methanol by first sparging it through a heated (--90~ water bath in order to sattrate it with water vapor and then exposing it to ultraviolet light at a wavelength of 185 nm in a quartz photochemical reactor. The suggested reaction pathway, shown in Scheme 1, proposes the

hv

H20

I. __>185nm= hV

1"!20 CH4 +

CH:~

+

=

H2 + 8 9

(1)

X > 185nm .OH =

892 +,OH

(2)

CH3~ + H20

(3)

H20

CI"I3OH +

(4)

=

892

SCHEME 1. Photochemical Conversion of Methane. Reference 1.

L.a/V~

hV

L >410nm

=

e~+

+

hvB

(1)

e ~ + MV 2+

=- MV, +

(2)

+ + H20 hvB

=

(3)

MV'++ H +

=

2.0H-----~ t-1202 ~

H + + ,OH 892 + MV 2+

(4)

H20

(5)

890 2 +

SCHEME 2. Catalytic Photolysis of Water. Reference 5. e-c~n= electron in Jr. ~.,x.J conduction band, h VB = positive hole in valance band, MV - methylviologen.

initial production of hydroxyl radical through photolysis of water. This radical may then react with a methane molecule to produce methyl radical. In the preferred reaction, the methyl radical then reacts with another water molecule to produce methanol and hydrogen. Catalytic photolysis of water to hydrogen and oxygen occurs during hxa~ation of liquid water with visible light (at wavelengths longer than 410 nm) in the presence of a solid photocztalyst suspended in the solution (Scheme 2) [5]. The photolysis sequence of interest initially produces a hydroxyl radical through the reaction of water in the presence of a doped tungsten oxide photocatalyst and an electron transfer molecule, methyl viologen dichloride hydrate (1, l'dimethyl-4,4'-bipyridinium dichloride). The proposed mechanism invokes the coupling of two hydroxyl radicals to form hydrogen peroxide, which decomposes to water and oxygen. By combining these reactions, hydroxyl radicals, generated with the photocatalyst and the electron transfer reagent, should react with methane to produce methyl radicals. In our

409

proposed reaction pathway (Scheme 3), methyl radicals react with an additional water molecule to form methanol and hydrogen, identical to Scheme 1.

hV

LaNkO3

-

Z. > 4 1 0 n m -

h

4-

ec8+ hvB

(1)

el~B + MV2+

= MVo +

(2)

+ , H20 hvB

.~ H * + , O H

(3)

MV,++ H +

=

(4)

CH4 + "OH

"

CH~ + H20

= CH3OH +

892 + MV 2+ CH3~ * H20

(5) 892

(6)

SCHEME 3. Photocatalytic Conversion of Methane e-Z 2 B = electron in 9 9 .. conduction band, h-VB = posmve hole in valance band, MV = methylviologen.

FIGURE 1. Schematic of photocatalytic reactor.

Previous research by our group [6] has confirmed literature reports [1,2] that it is possible to photolyze methane, saturated with water vapor, to produce methanol and hydrogen. In a modification of the above experiment, we were also able to photolyze methane sparged through a photochemical reactor filled with water. Recently, we began investigating the photocatalytic conversion of methane in water. 3. EXPERIMENTAL The reactor, a commercially supplied quartz photochemical reaction vessel, was fitted to meet the needs of this research (Figure 1). This included use of a Teflon-coated magnetic stirring bar in the reactor, a fritted glass sparger, a nitrogen line used to cool the UV lamp, and an injection port. Deionized water was distilled prior to use. The semiconductor photocatalysts were synthesized following a modification of the proc~ure in the literate [4]. Four dopants, copper, lanthanum, platinum, and a mixture of copper and lanthanmrt, were selected for study on the ttmgsten oxide catalyst base. In a typical experiment, 1.0000 grams of the sintered catalyst is suspended, by mechanical stirring, in water (--750 mL) containing an electron-transfer reagent, methyl viologen dichloride. A mixture of methane (5 ml_Jmin) and helium (16 mL/min) is sparged through the photocatalytic reactor. The helium is an intemal standard for on-line analysis of the reactor effluent. The reaction t ~ t u r e is maintained at --94~ by circulation of heated (-q 20~ silicone oil in the outer jacket of the reactor. A high-pressure mercury-vapor quartz lamp is used as the light source. The spectral characteristics and energy output of the lamp, supplied by the lamp's manufacturer, show that --46% of the radiated energy of the UV lamp used in this study is in the visible region. The outer surface of the lamp is cooled by a stream of nitrogen gas, while the lamp's immersion well is cc~led by a flow of tap water. The gaseous products of reaction

410 are analyzed on line and in real time by a quadrupole mass spectrometer. Liquid products are condensed from the gas stream at 0~ and analyzed by HP 5890A capillary gas chromatography. The chromatograph was equipped with a 60 m x 0.25 mm i.d. fused silica column coated with a 0.25 ~rn film of a crossed-linked version of SP-1000~ (commercially known as Nukol| Helium was used as the carrier gas with an average linear velocity of 30 cm/sec at 50~ After injection, the column was held at 50~ for ten minutes, then temperature programmed from 50~ to 200~ at 4 ~ per minutes with a five minute hold at 200~ The injector t ~ e was 200~ and the FID temperature was 250~ The injection volume was --2~tL. 4. RI~ULTS

The first series of experiments was conducted with no photocatalyst in the reactor. ~ g these experiments, a temperature d ~ d a n c e of the reaction was observed; photoconversion of methane decreased sharply with decreased temperature and was not observed below -~70~ This observation implies that a non-photochemical process is part of the reaction sequence. Several experiments were performed where the t ~ of the reactor was allowed to cycle between 60~ and 95~ In all experiments, as the t e ~ e of the reactor decreased, conversions of methane and the production of methanol decreased and were not observed below ~70~ The effect was reversible; when the reactor temperature increased above -~70~ conversion of methane and the production of methanol resumed and increased with temperature. Figure 2 shows the result of an experiment without photocatalyst where the reactor t ~ e was maintained at 97~ &ring the run. Note that conversions of methane remain relatively constant at -4% and production of hydrogen, methanol, FIGURE 2. Typical results of noncztalytic oxygen, and c.zrtx>n monoxide remain methane photoconversion. Observed flow of constant during the experiment. The products in mlJmin of CH3OH ( )'a~ large oscillations in the conversion of (O----O), H2 (t~---O), arid CD (D C3~ methane and the production of methanol were not observed during percent conversion of CH4 ( ............. ). Reaction conditions: He = 16 mL/min, CH4 = 5.0 this experiment. The four doped tungsten oxide mUmin, reactor temperature = 97~ catalysts (noted above) were synthesized and used in this study. The catalysts were analyzed by scanning electron microscopy (SEM), energy dispersive x-ray spectroscopy (EDS), x-ray diffraction (XRD), and electron

411 spectroscopy for chemical analysis (ESCA). For all catalysts except the platinum-doped tungsten oxide, these techniques were not able to detect any differences between the tungsten oxide as received and the unsintered-doped oxide because the level of doping, ~ 4 atom percent, is below the detection limits of these instruments. The sintering process produced differences that were detectable by SEM and XRD. After sintering, XRD dam showed the doped tungsten oxides to be more crystalline than the tmsintered materials as evidenced by the separation of a broad diffraction peak into two separate peaks having 2-theta values of 28.8 ~ and 42.0 ~ (Figure 3). Analysis of the sintered, doped tungsten oxides by SEM revealed that the sintered materials contained larger crystaUites with smoother edges. SEM and EDS analysis of the platinum-doped tungsten oxide photocatalyst after sintering showed the presence of platinum particles on the surface of the tungsten oxide. Figure 4 is a back-scattered electron image of the sintered platinum-doped tungsten oxide photocatalyst (the bright spheres are platinum). Analysis of the sintered platinum-doped tungsten oxide by ESCA revealed that the platinum on the surface is Pt". The catalysts were tested for their ability to catalytically photolyze water prior to their me in the methane conversion experiments. We were able to reproduce photolysis results reported in the literature [4] using these catalysts under similar conditions. All of the following FIGURE 4. Backscatter SEM Ofg~h3hr~t experiments were conducted with the photocamlyst. Bright circles in micro are photocatalyst and electron transfer platinum as identified by EDS. agent in the reactor. Figure 5 shows the results of a typical photocatalytic methane conversion experiment. Methane conversions are --4% with hydrogen and methanol as the main products of reaction. Note that after the UV lamp is turned off, the detected flow of methanol

412 decreases slowly to zero (over --2 hours). It was hypothesized that this behavior was due to stripping of methanol from the water in the reactor by the reactant gases. 0 . 0 s ~ ~ loo To confima this, methanol was injected t /I I ='~~__~;"4"x] into the reactor, previously filled with 0.04} // ON ON 9": ~ _ J ~ u ~ x , q l /-O- -F F I G" 750 mL at the operating I/ I L/ 1% temperature,w and er the concentration of 0.03~// I /~ ~/ ] ~n~' methanol in the gas flow from the => t./t_.. I ~ I u~F~ I t 6~ reactor was measured. Behavior :: similar to the reaction experiment was

!/I

o.ol

t~

I~

~:~~_~c=~

Iff

observed.

~

~J / ~ ~ ~ , ~ ' _ _ 4 TIME (HOURS)

6

~_140 820

FIGURE 5. Typical results of photocatalytic methane conversion. Observed flow of products hamtSn~ of~aoH ( ~ ch ( o - - o ~ I-h ( o - - - e ) , and CO ( m ~ m ) , ar~ reactor water temperature ( .............). WO3kLa photocatalyst.

Gas chromatographic analysis (Figure 6) of the liquid product that had condensed at O~ revealed the presence of methanol and acetic acid. Fm'ther analysis to identify other components by C~-MS was not possible due to the low concentration of products in the trap. The products were diluted by water carried over from the reactor in the flow of helium that is used as an intemal standard. As noted previously, the proposed reaction sequence of interest initially produces a hydroxyl radical,

which then reacts with methane to produce methanol. To test the validity of this hypothesis, a 30% solution of hydrogen peroxide, a good source of hydroxyl radicals, was injected into the reactor &wing p h o t ~ y t i c methane conversion. Figure 7 (a different photocatalyst preparation than Figtm~ 5) shows the results typical of the 30 peroxide injection experiments. After peroxide injection, 25 conversion of methane increases fi'om --4% to--10% methanol production increases 17 fold, and caz~n dioxide increases 5 fold, along with modest CH3COOH / 6.990 increases in hydrogen and carbon CH3OH monoxide. Introduction of hydroxyl ___.k radicals to the reactor leads to a 0 5 10 15 20 25 30 35 40 45 greater fraction of product going to TIME (MINUTES) methanol as evidenced by methane conversion increasing 2.5 times, whereas methanol production increases FIGURE 6. GC of condensed liquid product. 17 times. The increase in carbon dioxide is from "deep" oxidation of

I

413

0.20 ,..,.

200 pL 30% H202/ ADDED il~"~l.. \,

9

1

~9

i

10 2"

d!

8~E

N o.o5 -.................... tu

r~

]

~

.... w ~ ~O .

a.6

3.5

.

.

.

.

.

.

.

.

.

.

.

.

3.7

3.8

3.9

0

4.0

TIME (HOURS)

FIGURE 7. Results of addition of 200 laL of 30% hydrogen peroxide solution to a steadystate photocatalytic ~ i o n . Observed flow of products in ml_/min of CH3OH ( .... ), 0 2 !o-----o~ I-h ( e - - - e ~ c ~ e - - e ~ and O3 ( [ 3 ~ D ) , and percent conversion of CH4 ( ............. ). WO 3kLa photocatalyst.

,,,., 0.010 ._= E -J E 0.008

CH4 FLOW STARTEDf

~ 0.006 U. a LU

0.004

w

i

"'

L

a 0.001

1

TIME (HOURS)

FIGURE 8. Methanol production from various doped WO 3 photocxaalysts. WO3 dopants of La ( ~ ) , Pt ( O ~ O ) , La\Cu 50/50 (e-----el r~ c a a ~ (e- - - e l and Cu (D- - ~ l

[

methane and/or fiarther oxidation of the carbon containing products. Note the drop in methane conversion to zero for approximately 12 minutes after injection of the hydrogen peroxide. Prior to injection of hydrogen peroxide, a steady-state condition existed between the methane dissolving in the water and methane being consumed. It is likely that the introduction of excess hydroxyl radicals depleted the dissolved methane. This resulted in little methane available for conversion until steady-state conditions could be reestablished. Four doped tungsten oxide photocatalysts were prepared, platinmn, lanthanmn, copper, and a 50/50 mixture of copper and lanthanum. These catalyst were tested for their catalytic activity against a blank experiment, where all reaction conditions were identic~ except, that in the blank, no caA_alystwas present. Figtres 8-11 display the results for these experiments. After steady state conditions for the reactions were established (the reactor was at operating temperature and was being irradiated by the UV lamp), methane flow was started. In Figure 8 the production of methanol as a fimction of time is displayed. As shown in this figure, the lanthanum doped catalyst exhibits an increase in methanol production over the n o n ~ y d c reaction. The platinum and lanthanum/~ doped catalysts exhibit approximately the same production of methanol as the noncatalytic reaction. The presence of copper on tungsten oxide inhibits the production of methanol. The effects on hydrogen production for the various catalysts is shown in Figure 9.

414 Prior to the inEoduction of methane to the reactor, the n o n ~ y t i c reaction produces the most hydrogen. The only catalyzed system that produces hydrogen in significant quantities is the c/~4 0.14 one doped with lanthanum/eo~. FLOW .=_ A~er methane flow begins, the E 0.12 _J lanthanmn catalyst exhibits a large g O.lO increase in hydrogen production. A -0.08 modest increase is also exhibited by tl a the non-catalyzed and ,,, 0.06 t-c~ lanth~um\copper doped reactions. t.u 0.04 ]'he results of the addition of u.! o 0.02 l 2001LtL of 30% hydrogen peroxide solt~ion to the reactor converting oJ ~ 15 ~2 0 25 0 0.5 methane under steady-state conditions TIME (HOURS) are shown in Figtres 10 and 11. All of the calalysts exhibited larger peak production of methanol than the noncatalytic reaction after injection of FIGURE 9. Hydrogen production from various hydrogen peroxide. The lanthanum doped WO 3 photocaIalysts. WO3 dopants of doped catalyst showed the largest La ( 3 , ~ (O O), La\Cu 50/50 increase of all the catalyst with the copped doped catalyst exhibiting the least. The effect on the production of hydrogen was not as dramatic as those 0.20 ...., ... observed for methanol production with --200pL[/%, .=_ hydrogen peroxide injection. All the E H202 catalysts exhibited a slight increase in ""1 0.15 g hydrogen production while the noncatalytic reaction exhibited a decrease .J u. 0.10 in hydrogen production until the O LU I-hydrogen peroxide was consumed. LLI Analysis of the copper doped ~LU 0.05 E~ photocatalysts was conducted using scanning electron microscopy in an attempt to relate catalyst morphology d~)~ L.J ~LO.~ 0.4"" ~ ~ 0 . 6 0.8 1 1.2 characteristics to the low conversion TIME (HOURS) and selectivity of this material. Some areas of the surface of the copper FIGURE 10. Effects of addition of 200 ~tL of doped catalyst ~ e d smoother than 30% hydrogen peroxide solution on methanol that of the undoped WO3 ffigtre 12). production. WO dopants of La ( L ), Pt The grain structure of the tungsten ( O ~ O ) , La\C~ 50/50 mixture ( O ~ O ) , no oxide was obscured by this apparent catalyst ( 1 1 ~ 1 1 ) , and Cu ( ~ ~ [ 1 3 ) . coating. The smooth layer was examined using energy dispersive x-

,•o•1

415

FIGURE 12. photocatalyst.

SEM of Sintered WO3'\Cu

ray spectroscopy (EDS) in an attempt to detect differences in composition that might explain the change in morphology fi~om the extx~ed. . WO3 grain structm'e. However, if mcreas~ levels of copper were present in the smooth areas, they were below the limits of detection of EDS. This result is not surprising, since the coating was estimated to be around 0.1~tm thick, while the analysis depth was arotmd l~trn. Fm~er examination of the catalyst using X-ray photoelectron SlXCtroscopy (XPS) did indicate the presence of coptxr, but the signal was too small to accmmely determine its chemical state. The XPS copper peak was not distinct since the coming was not continuous over all catalyst particles, and the X-ray probe for XPS was of millimeter size. This resulted in contributions to the coptxr s ~ from both, coptxr in solid solution in the tungsten oxide, and from copper in the coating. It is possible that under the processing conditions, a thin layer of copper tungstate formed at the surface of the catalyst that effectively formed a barrier to the catalytic WO3 surface. The smooth appearance of ihis layer can be explained by the fact that copper tungstate melts around 1000~ and at 900~ (0.9 TH), one would expect appreciable atormc mobility and d i ~ i o n of the material in the coating layer, explaining its smooth appearance. Finally, more active WO photocatalysts did not exhibit tfi3 smooth surface appearance of the coptx~ doped material.

5. C(~CTUSIONS We have reproduced the results reported in the literature for both methane photolysis and catalytic photolysis of water. In experiments that combine elements of both systems,

416 methane and water have been converted to methanol, hydrogen, and acetic acid by a doped semiconductor p h o t ~ y s t at t ~ of--94~ and atmospheric pressure. Conversion of methane and the production of methanol are augmented by the addition of hydrogen peroxide, consistent with the postulated mechanism that proposes hydroxyl radical as an intermediate in the reaction sequence. The use of a UV filter or alternate visible light source is the next step to ensure that no UV induced reactions are occtrfing in parallel with the photocatalytic process. 6. ACKNOWI EDGMENT We would like to acknowledge the technical assistance of Richard R. Anderson, John Balmas, J. Rodney Diehl, E I ~ A. Frommell, Neil Johnson, and Joseph P. Tamilia. 7. DISCLAIMER Reference in this report to any specific commercial product, process, or service is to facilitate understanding and does not necessarily imply its endorsement or favoring by the United States ~ e n t of Energy. ~ C E S

.

K. O g t ~ M. Kataoka, J Mol. Ca., 43 (1988) 371-379. K. Ogtwa, C.T. Migita, M. Fujita, Ind Eng. Chem. Res., 27 (1988) 1387-1390. M. Ashokkumm', P. Marmhamuthu, j Mat. Sci. Lea., 24 (1988) 2135-2139. P. Maruthamuthu, M. Ashokkaamar, Int. J Hya~ogen Energy, 14 (4) (1989) 275-277. P. Mmaathamuthu, M. Ashokkaamm',K. Cmrtmath~ E. Subranmnian, M.V.C. Sastri, Int. J Hy&~ggenEnergy, 14 (8)(1989) 525-528. C.E. Taylor, R.P. Noceti, J.1L IYEste, Proceedings of the 15th Coal Liquefaction Gas Conversion Contractor's Review Meeting, Pittsburgh, PA, (1994) 777-783.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

417

Role of A- a n d B - C a t i o n s in C a t a l y t i c P r o p e r t y of S u b s t i t u t e d H e x a a l u m i n a t e (ABA111019_~) for H i g h T e m p e r a t u r e C o m b u s t i o n Koichi Eguchi, Hiroshi Inoue, Koshi Sekizawa, and Hiromichi. Arai Graduate School of Engineering Sciences, Kyushu University 6-1 Kasugakoen, Kasuga, Fukuoka 816 Japan

Cation-substituted hexaaluminate compounds, ABAlllO19.a, w e r e investigated for application to high temperature catalytic combustion. Two series of modifications of the compounds was made by cation substitution; substitution of large cations in the mirror plane with lanthanides ions, and substitution of transition metals for A1 site in the spinel block. In a series of AMnAlllO19-a (A = La, Pr, Sin, and Nd), surface area and catalytic activity increased with an increase in ionic radius of lanthanides. La 3§ is superior as the large cation in the mirror plane of the hexaaluminate to other tri-valent cations with small ionic radii. The catalytic activities of LaBAlllO19.a (B = Cr, Mn, Fe, Co, Ni, and Cu) were enhanced when Mn and Cu were employed as the B-site substituents. Although Mn and Cu were also effective substituents for enhancing catalytic activity in Ba-based hexaaluminate compounds, their activity was low as compared with the La-based catalysts. These results indicate that the redox cycle of transition metal in hexaaluminate lattice and catalytic activity appears to be affected sensitively with the electronic or structural effect of large cation in the mirror plane.

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

Catalytic combustion has attracted attention from the viewpoint of the protection of atmospheric environment [1-3]. As compared with conventional flame combustion method, emission of nitrogen oxides and unburned hydrocarbon can be significantly diminished and high energy efficiency can be achieved by using catalytic combustion method. For application of catalysts to high temperature combustion system above 1000 ~ it is desirable for the development of thermally stable catalyst materials. Conventional heterogeneous catalysts are immediately deteriorated at high operation temperatures (> 1200~ Further difficulty for the combustion catalyst is the wide range of reaction temperatures exposed to the catalysts. The

418 catalysts are requested to cover several different kinetic processes and/or their transient regions. Combustion is initiated as kinetically controlled surface reaction, since the catalyst temperature is low in the inlet zone. Mass transfer process controls the reaction rate when exothermic surface reaction reached to a certain conversion level. A further increase in t e m p e r a t u r e initiates gas-phase reaction and catalyst temperature reaches the maximum (>1200~ Some hexaaluminate compounds have been reported as a new material for high temperature catalytic combustion, since they possess excellent thermal stability in m a i n t a i n i n g large surface areas above 1300~ [4-6]. Partial substitution of some transition metals for A1 significantly promotes catalytic reaction due to high reduction-oxidation activity of transition metals, and Sr0.sLa0.2MnA111019~ is one successful design for the catalyst of high t e m p e r a t u r e combustor [6]. It m e a n s t h a t catalytic and thermal behavior of hexaaluminate is greatly influenced by composing cationic species. In this study, the cationic composition was optimized to attain high catalytic activity and thermal resistance by clarifying roles of A- and B- cations in substituted hexaaluminate catalysts (ABA111019-a).

2. E x p e r i m e n t a l

ABAlllO19.a (A = Sr, Ba, La, Pr, Nd, Sin, and Gd; B = Cr, Mn, Fe, Co, Ni, and Cu) samples were prepared by hydrolysis of metal alkoxides. Calculated amounts of Sr or Ba metal and Al(OC3H7)3 were vigorously stirred in 2-propanol at 80~ for 3 h. After dissolution was complete, an aqueous solution of other metal n i t r a t e s was added to this alcoholic solution. The resultant gel was dried and calcined at 1200~ for 5 h in air. Crystalline phase and surface area of calcined samples were characterized by XRD (RIGAKU, RINT-1400) and BET method. An analytical electron microscope (JEOL, JEM-2000FX) was used for estimating crystal morphology and composition. Catalytic combustion of methane over hexaaluminate catalysts was examined in a conventional flow reactor at atmospheric pressure. A gaseous mixture of CH4 (1 vol%) and air (99 vol%) was supplied at a space velocity of 48000 h -1. Methane conversion in the effluent gas was analyzed by on-line gas chromatography.

3. R e s u l t s a n d D i s c u s s i o n 3.1. Effect of A- c a t i o n in s u b s t i t u t e d h e x a a l u m i n a t e c o m p o u n d s AMnAlllO19-a (A = La, Pr, Nd, Sm, and Gd) samples were used to investigate the dependence of the catalytic property on A cation, which occupies the large cationic

419 site in the m i r r o r p l a n e of m a g n e t o p l u m b i t e - t y p e c r y s t a l s t r u c t u r e . F o r the AMnA111019.~ samples, except for A = Gd, the hexaaluminate phase was formed as a m a i n crystal phase, but a small a m o u n t of AA103 was formed as an i m p u r i t y phase. The amount of AA103 increased with a decrease in ionic radius of A. The cationic size of Gd in GdMnAlnO19_a sample was too small to lead to h e x a a l u m i n a t e formation, but resulted in phase separation of GdA1Os and a-A12Os. Surface area of AMnAlnO19.a (A = La, Pr, Nd, and Sin) samples increased monotonously with an increase in ionic radius of A (Figure 1). T10% is the temp e r a t u r e at which m e t h a n e conversion reaches 10%, and was lowered with an increase in ionic radius of A (Figure 1). This result suggests t h a t the sintering resistance of samples are enhanced as the ionic radius of A becomes large. It is, therefore, a p p a r e n t t h a t La s§ is superior as the large cation in the mirror plane of the hexaaluminate in achieving both high t h e r m a l resistance and high activity to other tri-valent A cations with small ionic radii. Every hexaaluminate compounds crystallizes in ~-alumina or magnetoplumbite structure (Figure 2). Both of these s t r u c t u r e s consist of alternative stacking of a spinel block and a mirror plane along the c direction [7]. Closed packing layer of oxygen is located in the spinel block; however, the packing in the mirror plane and along the c direction is relatively loose. These two crystal structures, being different in the coordination in the mirror plane, are regarded as hexagonal layered structures. The anisotropic nature of this crystal strongly affects the kinetics of crystal growth and diffusion. Iyi et al. studied the h e x a a l u m i n a t e

Fig. 2 Crystal structures of hexaaluminate compounds.

420 Table 1 Surface areas, crystal phases, and catalytic activities of Srz.xAxMnAlzlO19-a (A = La, Ce, Pr, Nd, Sin, and Gd). 0

0.2

0.4

0.6

1.0

-

-

H, P

0.98

19.7 b)

La

H c)

520 d) 755 e) 13.4

10.7

-

-

0.93

20.9 H,P 20.8 520 780 H,A, SA2. 17.8 555 780 H, P 520 780

14.9 H,P 520 770

17.7 H,P,A 530 770

12.5 H,P,A 520 810

0.91

18.6 H, P 500 770

16.9 H, P,A 520 780

11.1 H, P,A 520 770

0.91

Sm

17.6 H,P 520 770

16.8 H,P 520 780

13.7 H,P,A 510 780

6.5 H,P,A 570 870

0.89

Gd

17.2 H, P 530 820

15.2 H, P 540 780

13.9 H, P,A 520 780

5.3 Ia, A - -

0.87

Nd

H, F

485 760

H, F,A

Ce

9

17.5

rMa)

a) Ratio of ionic radii between Sr and A in Srl.xAxMnAI11019. b) Surface area (rn2 g-l) after calcination at 1200~ c) Crystal phase; H= hexaaluminate, P= perovskite, A= a-alumina, F= fluorite, SA2= SrAI407. d,e) Temperatures at which conversions of methane are 10% and 90%, respectively. Reaction condition; CH4, 1 vol%; air, 99 vol%; spece velocity, 48 000 h -1.

crystal structure in detail [8]. The thickness of mirror plane in the hexaaluminate is determined by the ionic radius, valence number, and number of elements in the mirror plane. The thickness of spinel blocks is also affected by the concentration of Frenkel defect due to excess charge in the mirror plane. In the previous study, we found t h a t Sr0.sLa0.2MnAlllO19-a calcined at 1300~ possesses the prominent catalytic activity and thermal stability as a result of partial substitution [6]. Therefore, the cationic substitution by lanthanide ions was further investigated for SrMnAlllO19-based h e x a a l u m i n a t e s for enhancing catalytic activity and/or thermal stability. Crystal phases, surface areas, and catalytic activities of Srl.xAxMnAlllO19~ (A = La, Ce, Pr, Nd, Sin, and Gd) catalysts are summarized in Table 1. La, Pr, and Nd are excellent dopants for enhancing catalytic activity and thermal stability as compared with SrMnAl11019-a catalyst. In Srl-xCexMnA111019~, CeO2 was deposited as impurity crystalline phase at every Ce concentration, since Ce is stable at the quad-valent state. Although, the ions with smaller ionic radius, e.g., Gd, lead to the formation of GdA103 and a-A1203, hexaaluminate structure was obtained at the range of 0<x_~<0.6. Lanthanide ions except

421

for Ce resulted in partial perovskite-type X=O o~r e, oxides (AA103) formation. Intensities of i i o i o their crystalline phase increased with an 9 o i .o .~, o ! , increase in extent of substitution and aX=0.2 ! o I i I A1208 f o r m e d at t h e r a n g e of x>0.6. These i m p u r i t y phases deteriorate the .,~ ; ~ 9 oi ~ o catalytic activity due to reduction of suro i eo foe 9 face areas. To elucidate the effect of l a n t h a n i d e ions on the catalytic properties of hexaI _ Ill II o ,t o aluminate, Srl-xPrxMnAl11Ols~ system was chosen as optimum substituent due 20 30 40 to its ionic radius and valence of Pr. The 2e / deg crystal structure of Srl.xPrxMnAlnO19-~ calcined at 1400~ are shown in Figure 3. T h e s a m p l e s consisted of the two Fig. 3 X-ray diffraction patterns of Srl.xPrxMnAlllO19-a calcined at 1400~ hexaaluminate phases at 0<_x~<0.4. Analytical electron microscopic analysis indicated t h a t these two h e x a a l u m i n a t e 5.7 phases are different in their contents of 0 0 0 Mn ions. To clarify the effect of the Mn c5.6 concentration in SrMnyAl12.yO19~(0l. Continuous shift of difFig. 4 Lattice constant of SrMnyfraction line was hardly observable in All2.yOl9-a. this region, but relative intensities of these phases were changed. It indicates t h a t concentration of Mn ions in the h e x a a l u m i n a t e lattice changes not continuously but relative amounts of two phases with low and high Mn contents depend strongly on the composition. A single hexaaluminate phase was obtained at 0.4<x
1

'

L

i

1

'

1

~

_

A

z

1

422 gradually to 5.608A. Pr ions are soluble by replacing Sr site in SrMnAlllOa9_~ host lattice. The perovskite phase is unequilibrium but sometimes formed as an impurity intermediate. Approximate composition of individual h e x a a l u m i n a t e particles was m e a s u r e d by an analytical electron microscope to elucidate the correlation between hexaalum i n a t e phases (Mn-rich and -poor) and Pr contents in the lattice (Figure 6). Although relative concentration of Sr to A1 was a constant at x=0, the composition of the particles in the sample was separated into two groups from their Sr/A1 ratios and Mn contents. P r ions are preferentially doped in the particles with small Mn/ A1 ratios at x=0.2. With an increase in Pr concentration, the number of Pr-poor 5.62 , , , , particles decreased and their Mn/A1 ratios approached to those of Pr-rich particles at x=0.4. The substitution of Mn 2§ t-- 5 . 6 0 _------Ofor A13§ site compensates excess charge tof P r 3§ in the hexaaluminate lattice and _.e leads to an increase of Mn in the hexa-~X 5.58 aluminate particle. It is considered that ease of reduction-oxidation cycle was influenced by the compensation and gave 5.56 rise to increase in catalytic activity for 0.0 0.2 0.4 0.6 0.8 1.0 m e t h a n e combustion. x in S r l - x P r x M n A I l ~ O ~ 9 Fig. 5 Lattice constant of Srl.xPrxMnAl11019-~. 0.2

x=0.4

x=0.2

x--O

9 o O0 "c -0 . 1

o 0.0

0.0

'

'

0.1

o'

Sr / AI

0.2

0.3 0.0

9

0.1

Sr / AI

0.2

0.3 0.0

0.1

Sr / AI

0.2

Fig. 6 Chemical composition of Srl_xPrxMnAlllO19.a. Q , Pr-poor particle.

0.3

423

3.2. E f f e c t o f B - c a t i o n i n s u b s t i t u t e d h e x a a l u m i n a t e compounds The effect of Al-site substitution on the phase and catalytic activity was investigated for the sample of BaBA]llO19.a and LaBAlllO19~ ( B = Cr, Mn, Fe, Co, Ni, and Cu). In a series of BaBA]11019-a, a single phase of h e x a a l u m i n a t e was obtained except for the case of B = Cr and Fe. Small amounts of impurities were observed in the case ofB - Cr (BaCrO4) and B = Fe (BaA1204 and a-A1203) even after calcination at 1200 ~ for 5 h. It appears to be difficult to form hexaaluminate with stable trivalent ions, such as Cr and Fe, as compared with cli-valent cations; Cu, Mn, and Ni. Catalytic activities and surface areas of these hexaaluminate compounds are summarized in Figure 7. A planar crystal morphology resulted from anisotropic crystal growth is effective in retaining high specific surface area even at 1200~ [5-6]. When BaAl12019-a with the low catalytic activity was submitted to the reaction, combustion appears to be initiated by the radical formation at the surface of the catalyst and progress t h r o u g h a chain reaction in the gas phase. This is characterized by the high initiation temperature and the steep rise in conversion. Substitution with transition metals significantly lowered the ignition temperature. The sequence of the catalytic activity, T10%, for methane oxidation was Mn > Cu > Fe > Ni = Cr = Co. However, complete combustion by BaMnAlllO19.a and BaCllA]llO19.a required higher t e m p e r a t u r e s than by B a A l l 2 0 1 9 and BaCoAlllO19-a. This is explained t h a t a mass transfer limitation levels off a high conversion region over Mn- or Cu- samples, whereas a gas phase reaction due to a radical formation brought about rapid propagation of methane oxidation near the surface over neat- or Co- samples. Catalytic activity for methane combustion of LaBAl11019-a was shown in Figure 8. Catalytic activity of LaBAlllO19~ 100 , ~-y ~F~ Surface area 9 M /m2g. 1 o a sample was also greatly dependent on 8o Cr 17.1 the t r a n s i t i o n metal species. In these -o Mn 21.9 a cA Fe 14.9 s a m p l e s , the catalytic activity of the .o 60 9 Co 13.3 ~ 9 Ni 16.4 sample with B = Cu demonstrated ex> 9 Cu 13.7 9A i o 40 cellent activity in the whole temperature o i /x region. A l t h o u g h combustion on the ~ 20 QI 9 sample with B = Mn started from lowest 0 200 600 800 1000 4O0 t e m p e r a t u r e a m o n g the substituented Temperature / ~ samples, LaCuA1110,9~ sample demonstrated excellent catalytic activity among any hexaaluminate-related combustion catalysts expressed by La- and Fig. 7 Catalytic combustion of m e t h a n e over BaBAlllO19-a (B = Cr, Mn, Fe, Co, Ba-BAlllOz9-a samples. Ni, and Cu). Reaction condition; CI-I4, T h e c a t a l y t i c a c t i v i t i e s (Tlo~) of 1 vol%; air, 99 vol%; space velocity, BaBAlllO19-a and LaBAlzlO19-a are plot- 48 000h 1.

424 ted against the enthalpies of reduction of the transition metal (B) oxides in Figure 9. The e n t h a l p y is calculated for the reduction of Cu e+ to Cu § and B 3§ to B 2§ for other cations. The relation shows a volcano-shaped correlation. This indicates that the oxidation-reduction cycle in basically operative for this reaction. The metaloxygen bond is most suitable for Cu and Mn, which is locating near the summit. Although the C u - s u b s t i t u t e d La-hexalOO Surface area M /rn2 g. 1 I 0 ~ 0 aluminate exhibited prominent activity, 80 m Cr 21.0 0 A~ o Mn 17.5 the activity of Cu was significantly low-tA Fe 14.2 A O 60 ered in Ba-based host lattice. In the se"~ 9 Co 19.6 0 O 9 Ni 25.4 ries of BaBA111019~ (B = Cr, Mn, Co, Ni, + 40 9 Cu 12.3 a n d Cu), t h e c a t a l y t i c a c t i v i t y of 0 A El& BaCuA111019~ was lower t h a n t h a t of ~ 20 BaMnAlI1019~ in the conversion level beo ~8~ ~; , 200 400 600 800 1000 low 50 %. The difference in the catalytic Temperature / ~ activity between La- and Ba-based hexaa l u m i n a t e s is suggested to r e s u l t from the valence of A cation or from the struc- Fig. 8 Catalytic combustion of m e t h a n e tural difference between magnetoplum- over LaBAII1OI9.a (B = Cr, Mn, Fe, Co, bite and ~ - a l u m i n a types. T h e redox Ni, and Cu). Reaction condition is the cycle of t r a n s i t i o n metal in hexaalumi- same in Figure 7 nate lattice and catalytic activity are affected sensitively with the electronic or Ni CoCu Mn Fe Cr 400 . . . . . . , structural effect of A site cation. o o The difference in catalytic activity be5O0 A A tween the La- and the Ba-based hexaoo o o a l u m i n a t e s r e s u l t s from the following "~ 6 0 0 o reasons; the first difference is the valence ~= A 7O0 of cation in the mirror plane between tria T10%(BaBAIl1019) v a l e n t l a n t h a n u m ion a n d d i - v a l e n t 0 T10%(LaBAIl1019) 800 barium ion. The second is the crystal 0 ~o 15o 200 -it/-/# / kcal g-l-O atom structure between magnetoplumbite and ~-alumina, which are different in the coordination of ions and concentration of Fig. 9 Relation between catalytic activFrenkel-type defect in mirror plane. The ity (T10%) of Ba- and La-based hexaaluredox cycle of transition metal in hexa- minates and enthalpy of reduction of the aluminate lattice, which closely related transiton metal (B) oxides per tool metal (AHr~). with catalytic activity, is affected sensiAHr~d = AHdBOt.s)-AH~BO) (B = Cr, Mn, tively with these two factors. Fe, Co, and Ni) or AH~CuO)-AHf(CuO0.5). !

425

Conclusions Modifications of hexaaluminate compounds were investigated for high temperature catalytic combustion. In a series of AMnAlllO19-a (A = La, Pr, Sm, and Nd), surface area and catalytic activity increased with an increase in ionic radius of lanthanides. La 3§ is superior as the large cation in the mirror plane of the hexaaluminate to other tri-valent cations with small ionic radii. LaCuAlllO19-a demonstrated the highest catalytic activity in LaBAlllO19-a and BaBA111019.~ (M = Cr, Mn, Fe, Co, Ni, and Cu) samples, and have the highest catalytic activity in a series of hexaaluminate catalysts so far investigated. These results indicate that the redox cycle of transition metal in hexaaluminate lattice and catalytic activity appears to be affected sensitively with the electronic or structural effect of large cation in the mirror plane.

Acknowledgment Present work was partially supported by the Grant-in-Aid for Scientific Research from the Ministry of Education, Japan.

References 1. 2. 3. 4. 5.

D. L. Trimm, Appl. Catal., 7 (1983) 249. L .D. Pfefferle and W. C. Pfefferle, Catal. Rev., 29 (1987) 219. R. Prasad, L. A. Kennedy, and E.Ruckenstein., Catal. Rev., 26.(1984) 1. M. Machida, K. Eguchi, and H. Arai, J. Catal., 123 (1990) 477. M. Machida, T.Shiomitsu, K. Eguchi, H. Arai, and Y. Shimizu, J. Solid State Chem., 95, (1991) 220. 6. M. Machida, T. Shiomitsu, FL Eguchi, H. Haneda, and H. Arai, J. Mater. Chem., 2, (1992) 455. 7. A. L. N. Stevels and J. M. P. J. Verstegen, J. Lumin., 14, (1976) 207. 8. N. Iyi, S. Takekawa, and S. Kimura, J. Solid State Chem., 83, (1989) 8.

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J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis -40th Anniversary Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

427

Au/Metal Oxides for Low T e m p e r a t u r e C O Oxidation Girish Srinivas, John Wright, C.-S. Bai, and Ron Cook TDA Research, Inc. 12345 W. 52 nd Avenue, Wheat Ridge, CO 80033, USA

Abstract Room temperature CO oxidation has been investigated on a series of Au/metal oxide catalysts at conditions typical of spacecraft atmospheres; CO = 50 ppm, CO 2 = 7,000 ppm, H20 -- 40% (RH) at 250C, balance = air, and gas hourly space velocities of 7,000- 60,000 hr1. The addition of Au increases the room temperature CO oxidation activity of the metal oxides dramatically. All the Au/metal oxides deactivate during the CO oxidation reaction, especially in the presence of CO 2 in the feed. The stability of the Au/metal oxide catalysts decreases in the following order: TiO2 > Fe203 > NiO > Co304. The stability appears to decrease with an increase in the basicity of the metal oxides. In situ FTIR of CO adsorption on Au/TiO 2 at 25~ indicates the formation of adsorbed CO, carboxylate, and carbonate species on the catalyst surface.

1. INTRODUCTION Low temperature oxidation is an attractive option for removing trace quantities of CO from ambient air in enclosed atmospheres such as submarines and spacecraft on long duration missions, and for industrial applications such as automotive cold start, ammonia synthesis, fuel cells, and CO 2 lasers. Traditionally, noble metals such as Pt and Pd supported on AI203 have been the catalysts of choice because of their high activity; however, they generally need to be operated at temperatures higher than 100~ Recent work [1-3] has shown that small particles of Au supported on various metal oxides such as C0304, Fe203, NiO, and TiO 2 are highly active for CO oxidation at low temperatures. The small size of Au particles and the method of preparation of the catalysts are both crucial to the high activity of the catalysts at low temperatures [1]. The objective of this study was to develop a low temperature CO oxidation catalyst that continually removes low concentrations of CO from the atmospheres of space stations. CO is a major contaminant in spacecraft environments. Since

428

power is scarce aboard spacecrafts, using a room temperature CO oxidation catalyst instead of conventional Pt or Pd supported on AI203 would result in substantial savings in energy. Thus, we studied catalysts composed of small particles of Au supported on various metal oxides as room temperature CO oxidation catalysts.

2. Experimental 2.1. Catalyst Preparation The activity of the Au/metal oxide catalysts is extremely sensitive to the method of preparation. The Au/metal oxide catalysts were prepared by the co-precipitating method [1]. During the course of this study, we have determined that the activity and the stability of the catalyst for room temperature CO oxidation were a function of Ph of the solution, temperature of precipitation, aging temperature and time, catalyst wash procedure, and calcination. In order to obtain reproducibility of catalyst synthesis, a semi-automated catalyst synthesis apparatus was used to prepare the catalyst. The catalysts were synthesized by co-precipitating a solution of hydrogen tetrachloroaurate (HAuCI4.XH20) and the metal nitrate using sodium carbonate (Na2CO3). The sodium carbonate was added dropwise to the solution of hydrogen tetrachloroaurate and metal nitrate at 60~ using an automated syringe pump. The precipitate was aged for 1 hour at 600C. The precipitate was then washed thoroughly with water three times to remove residual chlorine and sodium, dried in air overnight, and calcined at 350~ for 8 hours. These conditions were optimized to increase the low temperature performance of the catalysts. The Au/TiO 2 catalysts were made by precipitating Au particles on TiO 2 (p-25) similarly. 2.2. Reaction Studies The catalysts were tested for their CO oxidation activity in an automated microreactor apparatus. The catalysts were tested at space velocities of 7,000 60,000 hr 1. A small quantity of catalyst (typically 0.1 - 0.5 g.) was supported on a frit in a quartz microreactor. The composition of the gases to the inlet of the reactor was controlled by mass flow controllers and was: CO = 50 ppm, CO2 = 0, or 7,000 ppm, H20 = 40% relative humidity (at 25~ balance air. These conditions are typical of conditions found in spacecraft cabin atmospheres. The temperature of the catalyst bed was measured with a thermocouple placed half way into the catalyst bed, and controlled using a temperature controller. The inlet and outlet CO/CO2 concentrations were measured by non-dispersive infrared (NDIR) monitors. 2.3. In Situ FTIR Studies The in situ FTIR studies were conducted to monitor the formation and stability of surface species during CO chemisorption on the Au/metal oxide catalysts, using a Nicolet Magna IR 550 Fourier-Transform IR spectrometer equipped with a DTGS detector operating at a resolution of 4 cm 1. The IR reactor cell consisted of a central 314 in. hollow stainless steel tubing attached to circular stainless steel flanges. The reactor cell was sealed at the ends using CaF 2 windows with Viton O-

429

rings. The powdered catalyst was pressed into a wafer and placed in the reactor cell. Inlet and outlet lines were used to contact the pellet with the reactants. A thermocouple in contact with the pellet was used to monitor the temperature of the catalyst. The gaseous volume in the reactor was kept to a minimum by inserting two CaF 2 rods on either side of the pellet to minimize the intensity of the gas-phase bands in the IR spectra. The reactor could be heated with a heating mantle, capable of raising the temperature of the pellet to 500~ The IR spectra were obtained as a function of time and catalyst temperature. 3. RESULTS AND DISCUSSION CO oxidation tests on Au supported on various metal oxides were undertaken at low CO concentrations, where the adiabatic temperature rise in the bed is negligible. Since CO oxidation is highly exothermic, when high CO concentrations are present in the feed ( ~ 1%), and at high conversions, the adiabatic temperature rise in the catalyst bed due to the heat of reaction may be as high as 100~ Therefore, it is important to monitor the catalyst bed temperature when high CO concentrations are present in the feed.

100

c O Conversion

80 A

o<

CO - 50 ppm COe-0

H=O -

60

CO e selectivity

RH(2S~

25 ~ C balance air GHSV - 7,000 hr"1

40 20

0

2

4

6

8

Time (hr) Figure 1

CO oxidation on 5%Au/Co304 at a space velocity of 7,000 hr-1.

Figure 1 shows the CO oxidation on a 5%Au/Co304 catalyst at 25~ and a space velocity of 7,000 hr -1 in the absence of CO 2 in the feed. The CO conversion remains

430 at 100% throughout the 7 hour run. During the first hour, the CO 2 selectivity, {CO 2 formed/CO reacted} remained at 0. The CO 2 selectivity gradually increased and reached 100% only after 5 hours of reaction. This indicates that the catalyst adsorbed all the CO 2 that is formed from the CO during the first hour of reaction. This data also suggests that operating at 100% CO conversion could only provide transient data, and provide no information on the deactivation of the catalyst. As Figure 1 shows, the catalyst has a high enough activity at the low space velocity of 7,000 hr -~ to convert all the CO in the feed. If the catalyst is deactivating, it would be impossible to determine the extent of deactivation until the CO conversion decreased below 100%. In order to determine the extent of deactivation of the catalyst, if any, we performed the remaining tests at space velocities higher than 7,000 hr -1 and CO conversions less than 100%.

100 A

o~

s0 CO - 50 ppm CO==0

c:

0 m 60 0

a m

H=O - ~

s_..

0 o

0 0

RH(2S~

25 ~ C balance air GHSV - 60,000 hr"1

> e~

40

20

=

0

I

2

I

I

4

I

I

6

i

8

Time (hr) Figure 2 CO conversion on 5%Au/C0304 at a space velocity of 60,000 hr "1. Figure 2 shows the CO conversion as a function of time on stream in the absence of CO 2 in the feed on the 5%Au/Co304 at a space velocity of 60,000 hr 1. In contrast to the data shown in Figure 1, the catalyst showed an initial CO conversion of about 80% and showed considerable deactivation over eight hours. Figure 3 shows the CO conversion as a function of time at 25~ in the presence of CO 2 for C0304, TiO2, 5%Au/C0304, and 1%Au/TiO 2 catalysts at a space velocity of 60,000 hr -1. The data clearly shows that the addition of Au to the metal oxides

431

resulted in increasing the activity of the catalyst for CO oxidation at 25~ The data also showed that the initial activity and stability of the catalysts were different for different metal oxides.

100 i

Ii,l~r~am...

'

o<

.

- - ~ ' . "~1l ~&..

I

, . J' . , ~ . ~ .

9- - r , ~ , , m ~

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', ~ 1 ~ - ~ , , ~ . . , .

. . -.. ., ,., - - , , , . , . = ~ .

-

= "~O 60 -9 >

L i, "-

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i

0

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~

l Aurno=

CO - 50 ppm CO2 - 7000 ppm H=O - 40% RH(25~ 25 ~ C

---~.~_._

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~

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baJanoe Mr

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r-~e-li

n

n

i-i

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I-I

10

i

r-'!

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15

20

Time (hr) Figure 3

Effect of Au on the CO oxidation activity of metal oxides.

CO oxidation on 1%Au supported on various metal oxide catalysts was carried out to determine the effect of metal oxide on the activity and stability of the catalysts during room temperature CO oxidation. Figure 4 shows the CO conversion as a function of time on stream on 1%Au supported on various metal oxides such as Co304, Fe203, NiO, ZrO 2, and TiO 2. All the catalysts showed high initial CO conversions. The stability of the catalysts decreased in the following order: TiO 2 > ZrO 2 > NiO > Fe203 > C0304. The stability of the catalysts appears to decrease with increasing basicity of the metal. All of the Au/metal oxide catalysts deactivate quickly, under the conditions shown in Figure 4. In addition, the deactivation of the Au/metal oxide catalysts appears to be enhanced in the presence of CO 2. In support of the theory that increased basicity of the metal oxides leads to lower stability, we carried out CO 2 temperature programmed desorption experiments on the various catalysts. The CO 2 TPD data also confirmed that an increase in the basicity of the metal oxides leads to an increase in the amount of CO 2 adsorption on the catalysts. In situ FTIR studies of CO adsorption on a 1% Au/TiO 2 have identified various surface species on the catalysts. Figure 5 shows the in situ FTIR spectra of CO

432

100 L A

0 to L-0 c 0 o

8o r

1%Au/TIO e

\

1%Au/ZrO e

60

~"~'"~"~"~'~~--.----:---~W..

1%Au/NIO C O - 50 ppm

40

O0 e - 7000 ppm ~ 0 - 40% (eu4) T-25~

0 o

20 0

Figure 4

~

1%AulCosO4 0

balance air

~--. 1%Au/FeeO. ,

I

5

,

,,,i .........

10

GHSV -- eo,ooo hr"~ i

...... i

15 Time (hr)

i

20

I

I

25

,

80

CO oxidation on 1%Au on various metal oxides.

chemisorption on a 1%Au/TiO 2 catalyst. The spectrum prior to introduction of CO over the catalyst is shown for reference. Upon introduction of 1%CO in N2 over the catalyst, gas phase CO species appear at 2180 and 2130 cm 1. Flowing CO over the catalyst over time resulted in the development and increase in the intensity of bands in the 1700-1300 cm 1. The bands in the 1700-1300 cm 1 could be attributed to carboxylate and carbonate species. The uneven variation of the bands at 2180 and 2130 cm 1 could be due to the formation of adsorbed CO species on TiO 2 and on Au. These bands could be overlapping with the gas phase CO bands. Haruta et al. have attributed CO adsorbed on TP§ and on Au at 2183 and 2110 cm 1 [1]. Upon flushing the catalyst with N2 following introduction of CO (as shown in Figure 5), the gas phase and adsorbed CO bands disappeared. There was also a decrease in the intensity of the carboxylate and carbonate bands, suggesting that some of these species were weakly adsorbed at 25~ Further study is required to identify and determine the stability of all the adsorbed species. 4. CONCLUSION Au/metai oxides are active for low temperature CO oxidation. The activity of the catalysts is very sensitive to catalyst preparation. All the Au/metal oxides tested for room temperature CO oxidation deactivated substantially with time. The deactivation

433

0.04 -

g

o~o=""~176176,,,=~oo~

o~o=""~176176

t'~ I~ ~ i . .

0.08 o c .Q o 0.02.o <

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,

60 min

!

I,,,,

9

40 mln

~t

20 min

0.01 -

10 min

"I~O/N=

.

0.00

~ . . 2800

5 rain l . . . 2500 2200

. . 1900

Wavenumber Figure 5

. 1600

1800

'

t --0

{om"1}

In situ FTIR spectra of CO adsorption on 1%Au/mio 2.

increased in the presence of CO2 in the feed. The stability of the catalyst decreased with an increase in the basicity of the metal oxide. In situ FTIR studies suggest that adsorbed carboxylate or carbonate species on the catalyst surface could be inhibiting the room temperature CO oxidation reaction. A C KNOWLE DG EM ENTS The authors gratefully acknowledge support from NASA Marshall Space Flight Center under Contract No. NAS8-39344.

REFERENCES o

.

.

M. Haruta, S. Tsubota, T. Kobayashi, H. Kageyama, M.J. Genet and B. Delmon, J. Catal., 144 (1993) 175. S.D. Gardner, G.B. Hoflund, B.T. Upchurch, D.R. Schryer, E.J. Kielin, E. J. and J. Schryer, J. Catal., 129 (1991) 114. S.D. Lin, M. Bollinger, and M.A. Vannice, Catai. Lett. 17 (1983) 245.

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J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

435

P h o t o c a t a l y t i c O x i d a t i o n of Air C o n t a m i n a n t s b y C h l o r i n e (CI) or H y d r o x y l (OH) R a d i c a l s or H o l e s (h+): M e c h a n i s t i c Correlations Olga d'Hennezel and David F. Ollis D e p a r t m e n t of Chemical Engineering, Box 7905, N o r t h Carolina S t a t e University, Raleigh, NC 27695-7905, USA* We recently demonstrated that photocatalyzed destruction rates of low quantum efficiency contaminant compounds in air can be promoted substantially by addition of a high quantum efficiency contaminant, trichloroethylene (TCE), in a single pass fixed bed illuminated catalyst, using a residence time of several milliseconds [1-3]. Perchloroethylene (PCE) and trichloropropene (TCP) were also shown to promote contaminant conversion[2]. These results establish a novel potential process approach to cost-effective photocatalytic air treatment for contaminant removal. The mechanism(s) by which these photocatalyzed oxidations are initiated remain uncertain. Early proposals have included involvement of either the photo-produced holes (h § arising directly from semiconductor photo-excitation, or the (presumed) derivative hydroxyl radical (OH) which was argued to arise from the hole oxidation of adsorbed hydroxyls (h § + O H - - - > OH*). Recent subambient studies [4] with physisorbed chloromethane and oxygen suggest the dioxygen anion ( 0 2 ) as a key active species, and the photocatalytic high efficiency chain destruction of TCE is argued to be initiated by chlorine radicals (C1) [5]. The chlorine-enhanced photocatalytic destruction of air c o n ~ m i n a n t s has been proposed [1, 2, 6] to depend upon reactions initiated by chlorine radicals. No systematic studies of a number of compounds have yet appeared to discover correlations suggestive of mechanism. This paper presents the fractional conversions and reaction rates measured under reference conditions (50 mg contaminants/m 3) in air at 7% relative humidity (1000 mg]m 3 H20), for 18 compounds including representatives of the important contaminant classes of alcohols, ethers, alkanes, chloroethenes, chloroalkanes, and aromatics. Plots of these conversions and rates vs. hydroxyl radical and chlorine radical r a t e constants, vs. the reactant coverage (dark conditions), and vs. the product of rate constant times coverage are constructed to discern which of the proposed mechanistic suggestions appear dominant.

* The authors t h a n k the National Renewal Energy Laboratories for support of this research (Subcontract number XCK-5-14318-02).

436 1. I N T R O D U C T I O N From an economic point of view, Miller and Fox [7] indicate that photocat~ytic air treatment is commercially attractive, vs. activated carbon or incineration, for conversion of high quantum efficiency reactants such as trichloroethylene or methanol/ethanol in contaminated air streams. In contrast, photocatalyst process economics for compounds exldbiting lower apparent quantum efliciencies (below 20-30%) such as the aromatics, present in air vented from gasolinecontaminated soil (benzene, toluene, xylenes), and from paint drying operations (xylenes, odor compounds, etc.) compare unfavorably to incinerative and carbon adsorption processing. Since most candidate air contaminants demonstrated to be destroyed by photocatalytic air treatment exb_ibit lower quantum efficiencies, the breadth of the commercial future of photocatalytic applications depends centrally on finding approaches to raise to the order of 100% the a p p a r e n t quantum yield of pertinent air contaminants. Chlorine-enhancement may offer a partial solution. The addition of the chlorinated olefin TCE, PCE, or TCP to air/contaminant mixtures has recently been demonstrated to increase q u a n t u m yields substantially [1, 2, 6]. We recently have extended this achievement [3], to demonstrate TCE-driven high q u a n t u m efficiency conversions at a reference feed concentration of 50 mg contaminant/m 3 air not only for toluene but also for other aromatics such as ethylbenzene and m-xylene, as well as the volatile oxygenates 2-butanone, acetaldehyde, butyraldehyde, 1-butanol, methyl acrylate, methyl-ter-butyl-ether (MTBE), 1,4 dioxane, and an alkane, hexane. Not all prospective contaminants respond positively to TCE addition: a conventional, m u t u a l competitive inhibition was observed for acetone, methanol, methylene chloride, chloroform, and 1,1,1 trichloroethane, and the benzene rate was altogether tmaffected.

1.1. E n h a n c e m e n t M e c h e . i s m Berman and Dong [6] discovered this rate enhancement, labeled TCE a rate sensitizer, and suggested that "the function of the added sensitizer is to provide the radicals required to initiate chain propagated destruction of the pollutant". Nimlos et al subsequently proposed a specific chain mechanism [5] for the high quantum yield conversion of TCE itself, in which the initial photo-induced step, involving a photo-produced hole (h +) or OH- reaction with TCE, leads to the formation of a chlorine radical, which then reacts with TCE in the first step of a conventional chain oxidation. To rationalize the newly observed TCE enhancement of toluene conversion, Luo and Ollis [1] proposed a chain transfer step involving hydrogen abstraction from toluene by chlorine radical, and Sauer et al. [2] subsequently argued t h a t PCE and TCP cottld provide photocatalyst rate enhancement by the same chain transfer mechanism. 1.2. Rate d e p e n d e n c e a r g u m e n t There are two distinctly different mechanisms for a surface reaction between two spedes [8], for example toluene (T) and an active surface species (#). In the Langmuir-Hinshelwood (LH) mechanism, reaction occurs between toluene and the active surface species when both are adsorbed on the catalyst surface. If this step is the slow initiation step, the rate is proportional to the product of the coverages of toluene and the active site species: rLH = k CI)TcI)#

(1)

437 In the Rideal-Eley (RE) alternative mechanism, reaction occurs between an adsorbed active species a n d a gas p h a s e molecule (T). The r a t e is now proportional to: rl~ = k PT (I)#

(2)

The possible active species are: OH 9 radicals, the photo-produced holes (h § as suggested by Draper and Fox [9], the surface oxygen vacancies or anions (02") suggested by Lu et al. [4], and chlorine radicals (Clo) when chloroolefins (e. g. TCE) are p r e s e n t [1-3, 5, 6]. We m a y anticipate several possible behaviors for plots of photocatalytic rate vs. kinetic variable: (1) if hydroxyl or chlorine radical are involved, the surface rate at fixed gas phase concentration should correlate with koH or kc1 (Rideal-Eley), or koHCW and kclCT (Langmuir-Hinshelwood). These surface rate constants are unknown, b u t we m a y use the b i m o l e c u l a r gas phase values to e s t i m a t e expected r e l a t i v e reactivities. (2) if holes or vacancies are involved, and their concentration is constant, the surface rate should correlate with kh+ and ko2. (Rideal-Eley) or kh+ ~W and ko2-CT (Langmuir-Hinshelwood). These rate constants are u n k n o w n and have no gas phase analog. Since rate was observed to vary appreciably with 4~W, we consider instead a simpler plot of surface rate vs. ~T (which would hold if hole or anion vacancy reactivities were independent of contaminant structure). The present paper tests the assumed original and enhancement mechanisms with rates and conversions for a broad range of c o n ~ m i n a n t s measured under a fixed mass concentration (50 mg/m 3) feed condition. The plots compared are reaction rates vs. (1) dark adsorption, CW, (2) second order rate constant for (OHo) (TCE absent) or (Clo) (TCE present), and (3) the product of these gas phase second order rate constant times the reactant dark coverage. Where a second order gas phase rate constant was not available, we estimated its value from correlations of k cl vs. koH for the same class of compounds.

2. E X P E R I M E N T A L The near-IYV illuminated t i t a n i u m dioxide (anatase) powder flow reactor, as well as the gas chromatographic methods for analysis of effluents, has been described in detail earlier [2]. The procedures followed were substantially those of Luo and Ollis [1], and Sauer et al. [2]. 3. R E S U L T S We have screened 18 compounds in single contaminant air feeds, and co-fed in air with TCE [3]. For each run, the initial inlet concentration is first measured. Then, we allow a "dark" period during which the contaminated air feed passes

438 through the bed w i t h o u t illumination until the outlet concentration equals the inlet. This p r e t r e a t m e n t is required because both the catalyst and the fritted glass h a v e considerable surface areas. Consequently, a n appreciable time, varying w i t h each pollutant from m i n u t e s to hours, is needed to r e a c h gas-solid equilibrium. For each compound we m e a s u r e d the specific d a r k adsorption on TiO2 (molecule cm-2cat). D a t a collected from the l i t e r a t u r e include gas phase second order rate constants for the contaminant reaction with hydroxyl radicals (koH, cm 3 molecule "1 s "1), and with chlorine atoms (kcl, cm 3 molecule -1 s "1). The m a x i m u m reaction r a t e s and conversions reported here r e p r e s e n t d a t a at early times before any deactivation bec~me i m p o r t a n t [3]. The m a x i m u m reaction rate is calculated following: r= Fl0w rate ....,Initial pollutant concentration,N, ...Maximum,conversion Illuminated catalyst surface 100 Pollutant molecular weight Flow rate = 0.828 cm3/s I l l u m i n a t e d c a t a l y s t s u r f a c e = 0 . 2 6 8 " 1 0 4 cm 2. Teichner et al. [10] have shown that 99% of light absorption occurred within approximately 5 ~ n depth of TiO2 powder.

Initial pollutant concentration = 50 mg]m 3 N = Avogadro's n u m b e r

3.1. koH / k c ! correlations in homogeneous photochemistry ..... . . . . . . . .

I

. . . . . 9 ' ' '~*~ ~ r ~ ----o---

10

V

.

"10

..U"

B~S,~

--- "=

~,0+

....

_

-L) m{9

/

O

E

o

/

."

03

E o

v

1 0 "11 o

, ,,

,

.

i

l

10 -~a

i l ill

+

_

!

t

10 "12 k

OH

(cm 3 molecule1

!

!

i

t

I iI

L

|

n-alcohols

-9- - e-

n-alkanes -

n-ethers

- - A - -

chloroethenes

- - .-

-

aldehydes

---=-

-

1-chloroalkanes

--+---

aromatics

wA..-

ketone

10 "~1 s'l

)

Figure 1: kct vs. koH. Second order gas phase rate constants for the reaction of C1 atoms vs. the corresponding OH radicals rate constants for the reactions with a. n-alkanes [11] b. n-alcohols [12] c. n-ethers [12] d. chloroethenes [13] and e. 1-chloroalkanes [14]. Figure 1 present kcl versus koH values for families of compounds w h e r e both rate constants were found in the l i t e r a t u r e for n-alcohols, n-alkanes, n-ethers, chloroethenes, 1-chloroalkanes, aromatics, aldehydes, and ketones. This graph demonstrates that: 9 kci is always greater t h a n koH. 9 on average, kc1 is g r e a t e r t h e n koH by b e t w e e n one to two orders of magnitude. 9 the t r e n d b e t w e e n koH and kcl for each f a m i l y of compounds can be reasonably represented by the linear correlations shown.

439

For n-alkanes, n-alcohols, 1-chloroalkanes, n-ethers, and chloroethenes, the carbon chain length influences the reactivity, and the clear linear correlations indicate that the attack mechanism of these pollutants by OH 9 or CI 9 radicals occurs via the same pathway. However, such correlations do not hold true for aromatics, ketones, and aldehydes, for reasons discussed in our previous paper [3]. We also estimated missing values of kcl by analogy: for ethylbenzene, we take kc1 = 1.5e-10 cm 3 molecule "~ s -~ , greater than that for m-xylene, but smaller than the 2.0e-10 cm 3 molecule -~ s -~ value for very reactive compounds. Also we estimate a similar value for butyraldehyde; kcl = le-10 cm 3 molecule -~ s -~, only 10% larger than kcl of acet~dehyde to remain consistent with the equivalent koH value. 3.2. S p e c i f i c d a r k a d s o r p t i o n t r e n d s Figure 2a and 2b present the contaminant m a x i m u m conversions and corresponding initial rates plotted vs. the dark adsorption. These figures reveal the general trend that as the dark adsorption increases, both the conversion and the initial rate rise. This trend holds in both the presence (Figure 2b) and absence of chlorine (TCE) (Figure 2a); the adsorption of the pollutant in the dark is unchanged in the presence of TCE, since the dark adsorption of TCE is very low. TCE rate promotion is responsible for the rate increases, often to nearly 100% conversion, seen in Figure 2b. 100

.....

+

~ + +

+

§

#::

4 * 1 012

+

r

r

.....

~40

-

0

+

20 -

m

3o

"

O

1012

60

0

40

1011

§ o ,

....

,.l

1

013

conversion _ ~ ,..= = rate , , ...... 1 . . . . . . . 4 * 1 09 1 0 TM

1 015

Dark a d s o r p t i o n ( m o l e c u l e c m "2

)

cal

$"

3

20

O O

101~

0

,.,,,

o

,,,,,,,.. ,,,,...

.

O O

O

ITI

-1

O o...

e-

1 012

~ 1013

"::

...,..,

3 o_.

O

> E

0

,~

o

L_

; ;

+

.

0

80

+'+ #+

-

m 60

=

..4

--L

v

~O

o~

,~. ~..

o~'80

100

+ o

O e-eE

w

,

1 012

.......

1

~m O

conversion ~. rate I

013

,

,,jl,,,I

1 0

I

TM

3

t,,,,r~

10

9

1 015

Dark adsorption ( m o l e c u l e c m 2

) cat

Figure 2a: Maximum single pass conversion and initial rate in the absence of TCE for each compound vs. dark adsorption. (+ conversion, o, * rates; filled circles i n d i c a t e compounds for which conversion is less than 95%).

Figure 2b" M a x i m u m single pass conversion and initial rate in the presence of TCE for each compound vs. dark adsorption. (Same symbols as (2a)).

440

3.3. Rate constant

trends

(Rideal-Eley)

Figure 3a presents the maximtlm conversion and the initial rate of each compound versus the literature second order rate constant for contaminant reaction with hydroxyl radicals. The conversion increases with increasing values of koH, but with considerable scatter. For chlorine enhanced conversion and initial rate vs. the second order chlorine rate constant (presence of TCE) (Figure 3b), the trends are clear: conversion and initial rate increase with increasing kcl values. For values ofkcl above 2e-l l cm 3 molecule -t s -1, the enhanced conversion for all corresponding compounds is greater than 80% per pass. 100

~80

-

§ §

v

-

60

..,,.

+

,--80

..~,,.

-

§

++~:.~

100

4 " 1 0 1 2

o. 4 0

+ 9

0 toO_

+

9

0

0

0 0

-

20

9

%

I1)

c

+

o

9

--z

o

+

3

0

-

0

o0. .

:t - ' 2 0

0

O 0

"1o 0

o

0

e-

, , J,..l

s i .J,,,J

, , ,,-J

OH

(cm 3 molecule" 1 s 1)

Figure 3a: Maximum conversion in a single pass reactor and the initial rate (no TCE) vs. literature second order h y d r o x y l r a t e c o n s t a n t s . (Same symbols as (2a)). 3.4. (Rate constant Hinshelwood)

* specific

_

m_.,.

-

..,

3o

t101~

.~=

. -

+ o

conversion ~,=

i rate= ,2-J 10 g

.......

(D c~ 0

~

,

3

' ~ .L

1 0 "13 1 0 "121 0 "111 0 "1~

10 "14 10 "1310 "1210 "11 k

5" .,.,.

.,...,,

L ,,,d . . . . . . . d . . . . . . .

4 " 1 0 "9

Q-

:"

:

.

_+

r

+

-

W

, , ,,,t.[

%

+

0 IID

0

- 1 0 1 2

~40

3

conversion - - - e - - - rate

W ::~

o60

0

0

~,.

+"

.

~. o

0

1 013 m -

k

ci

(cm 3

molecule

1 s " 1)

F i g u r e 3b: Maximum enhanced conversion in a single pass and the enhanced initial rate (TCE added) vs. literature second order chlorine rate constants. (Same symbols as (2a)). dark

adsorption)

trends

(Langmuir-

The conversion and initial rate in the presence and absence of TCE versus the product of the second order rate constant and the dark adsorption appear in Figure 4a,b. Figure 4a shows considerable scatter in the data, revealing only broad, general trends between conversion or rate and hydroxyl second order rate constant. However, the plots of enhanced conversion and initial rate vs. the corresponding chlorine second order rate constant multiplied by the dark adsorption data are smoother (Figure 4b).

441

,- : ~ - 4 -

100 §

o~" 8 0

+

v

+

.

.=.

+ + "

c'-

6O

+

W

o60

o~_

20

:: +

+

~1013

m

"

W

. .

..~

+

0

c40

~0

9~ ~ _ o

+

0

+

@-

?o

o_'

1011

.

9

~"

3

.

>e - 2 0

0

o

.

e,-

>

0

"~

0

~-

+

101~

_

o

0

0

+

. _

o

conversion =

* Dark

"1

101

adsorption

+

e-e'-

3

9

,r

(cm 3 cm 2

cat

s "1)

Figure 4a: M a x i m u m conversion in a single pass and the initial rate in the absence of TCE vs. the product of second order hydroxyl rate constants time the dark adsorption. (Same symbols as (2a)).

k

CI

* dark

=

rate

"

~i,=.i ~,,,"J .... ,,I .... ,,,I , ,,,,,I , ,,'~

1 O" 1

103

"

conversion

o

w

rate

...... ~ ..... .1 ,, ....J ..... .! .... .,I ,,, 4 " 1 0 10

OH

o~ ,

~ 0

e-

k

~,-

_ 1012 _.~

0

++

~

0

o~ L_

~80

3

0

o

_., _,~.. W

O)

+

c, 4 0

+~++

100

1012

1 01

adsorption

1

0 3

1

109 05

-

(cm 3 cm 2

~. v

cat

s

-1)

F i g u r e 4b" Enhanced maximum conversion in a single pass and the enhanced initial rate (TCE added) vs. the product of second order chlorine rate constant time the dark adsorption. (Same symbols as (2a)).

4. D I S C U S S I O N 4.1 In t h e a b s e n c e o f (CI) TCE In the absence of TCE and chlorine, the possible active species are holes (h§ anion vacancies, or anions ( 0 2 ) , and hydroxyl radicals (OH 9 At constant illumination and oxygen concentration, we may expect h § and O2 concentrations to be approximately constant, and the dark adsorption to be a dominant variable. Ifkh+, or ko2- does not vary appreciably with the c o n ~ m i n a n t structure, the rate would depend clearly on the contaminant coverage as shown in Figure 2a, and the reaction would therefore occur via Langmuir-Hinshelwood m e c h a n i s m . (Note: only rates with conversions below 95% are correlated here (filled circles), as the 100% conversion data contains no kinetic information). This rate vs. r LH plot is smoother than those for koH or koH @T, suggesting that non-OH species (holes, anion vacancies, or O2) are the active species reacting with an adsorbed c o n ~ m i n a n t . 4.2 In t h e p r e s e n c e o f (CI) TCE The rate of gas phase reaction of pollutants with chlorine radicals is 10 to 100 times faster than with hydroxyl radicals, and we [1, 2] have proposed C1 9 as the surface active species responsible for the rate enhancement observed on addition

442 of TCE, PCE or TCP. The rate vs. kcl or kc1 (]~Tcorrelations developed in Figures 3b and 4b, respectively verify this involvement, but a choice between the LH and RE mechanism is not clearly evident from the available data. 5. C O N C L U S I O N 9 We can correlate our experimental conversions and rates with the extent of (dark) cont3minant adsorption (~), and the literature homogeneous second order rate constants for OH 9 and C1 9 radicals, and the product of the rate constant times coverage. 9 In the absence of chlorine atoms, our results suggest that hydroxyl radical is not the reactive site; other alternatives not tested here would be a hole (h§ an oxygen vacancy, or a dioxygen anion ( 0 2 ) as recommended by Fox [9] and Yates [4] respectively. 9 In the presence of chlorine atoms, the chlorine radical appears to be the active surface species. It is not possible from our limited data to establish whether most reaction occur via Langmuir-Hinshelwood or Rideal-Eley mechanisms. 9 However, none of the correlations a t t e m p t e d display r a t e v a r i a t i o n proportional to reactant coverage to the first power. F u r t h e r work under conditions of lower conversion per pass will allow inclusion in these correlations of all compounds examined here and should assist in resolving current mechanistic uncertainties. REFERENCES

1. Y. Luo and D.F. Ollis, J. Catal., (1995) Submitted. 2. M.L. Sauer, M.A. Hale, and D.F. Ollis, J. Photochem. Photobiol. A: Chem., 88 (1995) 169. 3. O. d'Hennezel and D.F. Ollis, J. Cat~., (1996) Submitted. 4. G. Lu, A. Linsebigler, and J. J. T. Yates, J. Phys. Chem., 99 (1995) 7626. 5. M.R. Nimlos, W.A. Jacoby, D.M. Blake, and T.A. Milne, Environ. Sci. Technol., 27 (1993) 732. 6. E. Berman and J. Dong, in "The Third International Symposium Chemical Oxidation: Technology for the Nineties" (W.W. Eckenfelder, A.R. Bowers, and J.A. Roth, Eds.), p. 183. Technomic Publishers, Chicago, 1993. 7. R. Miller and R. Fox, in "Photocatalytic Purification and Treatment of Water and Air" (D.F. Ollis, and H. A1-Ekabi, Eds.), p. 573. Elsevier, Amsterdam, London, New York, Tokyo, 1993. 8. M.G. White, "Heterogeneous Catalysis" (N.R. Amundson, Ed.), p. 202, Prentice Hall, Englewood Cliffs, 1990. 9. R.B. Draper and M.A. Fox, Langmuir, 6 (1990) 1396. 10. M. Formenti, F. Juillet, P. Merideau, and S.J. Teichner, Chem Tech 1, Nov ( 1971) 680. 11. S.M. Aschmann and R. Atkinson, Int. J. Chem. Kinet., 27 (1995) 613. 12. L. Nelson, O. Rattigan, R. Neavyn, and H. Sidebottom, Int. J. Chem. Kinet., 22 (1990) 1111. 13. R. Atkinson and S.M. Aschmann, Int. J. Chem. Kinet., 19 (1987) 1097. 14. T.J. Wallington, L.M. Skewes, and W.O. Siegl, J. Phys. Chem., 93 (1989) 3649.

j.w. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

443

Partial o x i d a t i o n o f m e t h a n e to synthesis gas o v e r R u / T i O 2 catalysts

Y. Boucouvalas, Z.L. Zhang, A. M. Efstathiou and X. E. Verykios Department of Chemical Engineering

and Institute of Chemical Engineering and High

Temperature Processes, University of Patras, GR-26500 Patras, GREECE The catalytic partial oxidation of methane to synthesis gas is investigated over Group VIII metal catalysts. It is shown that while all catalysts promote methane combustion followed by reforming with H 2 0 and CO 2, the Ru/TiO 2 catalyst, to a large extent, promotes the direct formation of synthesis gas. The existence of the direct reaction route is probed by steady-state isotopic transient experiments. The extent of the direct route is found to be sensitive to modifications of the TiO 2 carrier. FTIR and XANES studies indicate that the unique performance of the Ru/TiO 2 catalyst is related to its high resistance to oxidation, which renders high selectivity to synthesis gas in the presence of oxygen. 1. INTRODUCTION The catalytic partial oxidation of methane for the production of synthesis gas is an interesting alternative to steam reforming which is currently practiced in industry [1]. Significant research efforts have been exerted worldwide in recent years to develop a viable process based on the partial oxidation route [2-9]. This process would offer many advantages over steam reforming, namely: (a) the formation of a suitable H2/CO ratio for use in the Fischer-Tropsch synthesis network, (b) the requirement of less energy input due to its exothermic nature, (c) high activity and selectivity for synthesis gas formation. Concerning the reaction pathway, two routes have been proposed: the sequence of total oxidation of methane, followed by reforming of the unconverted methane with CO 2 and H20 (designated as indirect scheme), and the direct partial oxidation of methane to synthesis gas without the experience of CO 2 and H20 as reaction intermediates. The results obtained by Schmidt and his co-workers [4, 5] indicate that the direct reaction scheme may be followed in a monolith reactor when an extremely short contact time is employed at temperatures in the neighborhood of 1000~ However, the majority of previous studies over numerous types of catalysts show that the partial oxidation of methane follows the indirect reaction scheme, which is supported by the observation that a sharp temperature spike occurs near the entrance of the catalyst bed, and that essentially zero CO and H 2 selectivity is obtained at low methane conversions (<25%) where oxygen is not fully consumed [2, 3]. A major problem encountered

444 when this reaction pathway is followed is highly uneven axial temperature distribution within fixed bed reactors due to large quantities of heat produced by combustion reaction(s) at the entrance of the catalyst bed, which could bring the local temperature well above 1000~ This results in several undesirable consequences, i.e. catalyst deactivation due to sintering, danger of explosion, and demand for special high-temperature materials of construction. On the other hand, the direct partial oxidation is of high industrial interest since the problems associated with catalyst deactivation and steep temperature profiles along the catalyst bed can be largely eliminated. However, several stringent requirements are imposed on the catalyst which must be able to effectively retard the deep oxidation of CO and H 2 which are thermodynamically more favorable. Thus, it is highly improbable that the direct reaction scheme can be followed over a wide range of catalysts. 2. EXPERIMENTAL Supported metal catalysts (M=Rh, Ru, Pd, Ir, and Ni) were prepared by the method of incipient wetness impregnation, using metal chloride and/or metal nitrate as the precursor compound for the metal [10]. The apparatus employed for kinetic measurements consists of a flow measuring system (MKS thermal mass flow meters and control valves), a heated quartz tube (4 mm, i.d.) reactor and an on-line gas chromatograph. Reactor operating conditions with respect to total feed flow rate and average catalyst particle size were defined so as to eliminate intraparticle and interphase transport resistances [9]. Transient isotopic labelling experiments were performed on a switching apparatus which has been described earlier [11]. Chemical analysis of the gases was done by an on-line mass spectrometer (Fisons, SXP Elite 300H) equipped with a fast response inlet capillary system. Calibration of the mass spectrometer signal was performed based on mixtures of known composition. An FTIR (Nicolet 740) spectrometer equipped with a DRIFT cell was employed to study CO adsorption and the surface species formed under reaction conditions. X-ray absorption near edge structure (XANES) experiments were conducted in Hasylab/Hamburg, concentrating on the L and K edges of Ru crystallites, following different pretreatments. The catalyst particle was finely powderized and the thickness and homogeneity of the sample film were carefully checked in order to eliminate the "thickness effect". 3. RESULTS AND DISCUSSION 3.1 Kinetic Studies

3.1.1

Unique performance of the Ru/TiO2 catalyst

Controlled elimination of mass and heat transport resistances is an important prerequisite for obtaining intrinsic kinetic parameters of the fast exothermic reaction of partial oxidation of methane to synthesis gas. It has been demonstrated that under conditions of strong transport limitations erroneous conclusions concerning the reaction scheme can be derived [7-9]. It was determined in this laboratory that transport limitations are practically absent over a wide range of operating conditions if one portion of the catalyst (< 40 ktm) is diluted with -5 portions of an

445

100 80

750 ~ zx A -"AZXZ~ZxK" ZX

6O

9

7

J

80 60 40

40

.550_~ . . . . . r

2O

20

550~d 6oo~ 0

20

(a) 40

5~176176 ~

(b)

_r 6O

XC H4 (%)

600~ . . . . . . . . . . . . . . . . . 10 20 30 40

50

60

70

XCHa (%)

Figure 1. Influence of methane conversion on a) CO selectivity and b) H 2 selectivity over various catalysts under lean feed conditions, CH4/O2/He=4/2/94 vol.% (zx,,t, 9 : Ru~iO2; I"I:Ru/AI203;

x : Ru/ZrO2;

7:Pd/TiO2;

9 :Rh/AI203;

o :Ir/Al203;

+: Ni/AI203).

V:Pd/A1203; *: Rh/La203;

,: Rh/SiO2;

inert solid (e.g. 0t-A1203), the feed gas is heavily diluted with He, up to 94 vol.% and the total feed flow rate is kept above 200 ml/min [9]. Unless mentioned otherwise, these reaction conditions were applied in the present study. Detailed studies were carried out for assessing the intrinsic differences between the unique Ru/TiO 2 and conventional catalysts which include Ni, Rh, Pd and Ir catalysts, as well as Ru catalysts supported on carriers other than TiO 2. CO and H 2 selectivities are shown in Figure 1 as a function of methane conversion, which was varied by changing contact time. Attention is focused in the low methane conversion region in which oxygen conversion is less than 100%. Due to the presence of oxygen in the reaction mixture, the metal surface is anticipated to be oxidized or partially oxidized, which renders the rate of reforming reactions negligibly slow as compared to that of oxidation reactions. In this range, any synthesis gas formation is mainly attributed to the direct partial oxidation of methane. It is shown in Fig. 1(a) that over Rh, Pd and Ni as well as over Ru on carriers other than TiO 2, CO selectivities are essentially zero when methane conversion is lower than ca. 25%. Selectivity to synthesis gas was found to increase rapidly at higher methane conversions. This observation can be explained evoking the indirect reaction route, i.e. methane is first deeply oxidized to H 2 0 and CO 2, and, as oxygen is completely consumed, corresponding to ca. 25% of methane conversion (CH4/O2=2), reforming of the unconverted methane with H 2 0 and CO 2 occurs to produce synthesis gas. The absence of gas phase CO and H 2 at low methane conversions suggests that CO is not a primary product of the reaction. CO and H 2 selectivities obtained over the Ru/TiO 2 catalyst at variable methane conversion

446 are also shown in Fig. 1. It is evident that significantly high CO and H 2 selectivities are obtained over the Ru/TiO 2 catalyst at methane conversions approaching zero. This suggests that over this catalyst CO and H 2 are formed as primary products, not as secondary products from CO 2 and H 2 0 reforming. Increasing reaction temperature from 550 to 750~

results in a

sharp increase in CO and H 2 selectivities over the Ru/TiO 2 catalyst at the low methane conversion region, indicating that higher temperatures favor the direct reaction route.

3.1.2 Influence of modification of the TiO2 carrier In order to assess the performance of the Ru/TiO 2 catalyst under realistic conditions of operation, undilute CH4/O 2 (=2) feed mixtures were also used. It was observed that the axial reactor temperature profile was relatively flat when the Ru/TiO 2 catalyst was used and no temperature runaway conditions were developing. The unique behavior of the Ru/TiO 2 catalyst, which is capable of catalyzing the direct partial oxidation route to a significant extent, may be attributed to some type of interaction between the Ru crystallites and the TiO 2 carrier. The nature of this interaction is not clear at this moment. However, as Figure 2 illustrates, the Ru-TiO 2 interaction can be enhanced in the direction which favors the direct partial oxidation of methane, or in the opposite direction, by modification of the TiO 2 carrier. Thus, doping TiO 2 with small quantities (- 1 wt%) of ZnO, CaO or Li20 results in significant enhancement of CO selectivity at low methane conversion while it does not affect H 2 selectivity to a significant extent (Fig. 2b). A significant reduction of CO selectivity was observed upon doping TiO 2 with 1 wt% of WO 3 (see Table 2). A spectacular enhancement in H 2 selectivity is observed when the TiO2 carrier is doped with 1 wt% La203 (Fig. 2b). In this case, at methane

"//

9()

100 90 ~()

&l

~0

-

o~0

A /

---

o j' A

G

- O ~ c D

-

70 6O [3

9

70

j

9

5O 9

(a)

4O

(b)

o (

I

)

20

,,,

1

!

3o

40

I .......

50

60

70

80

I0

Xc:t~ (%)

20

30

40

50

60

70

g0

XcI In ('7o)

Figure 2. Effects of altervalent cation doping of TiO 2 on a) CO selectivity and b) H 2 selectivity under concentrated feed conditions, CH4/O2=66.7/33.3 vol.% (TiO 2 doped with lwt% of ,," ZnO; A" CaO" o" La203" El " Li20;

I " undoped).

447 conversions less than ca. 30%, H2 selectivity is close to 100% and drops suddenly to 80% at slightly higher methane conversions due to oxidation of H2 to H20. Upon full consumption of oxygen, H 2 selectivity increases significantly, due to the occurrence of reforming reactions. Qualitatively similar results were also obtained when dilute feed mixtures were used. 3.2 Catalyst characterization and Mechanistic Studies 3.2.1 Steady.State isotopic tracing experiments The following isotopic labeling experiment was performed in order to quantify the contribution of the direct and indirect reaction routes to CO formation: After steady-state reaction with CH4/O2/He was achieved, an abrupt switch of the feed from 12CH4/O2/H e to an isotopic mixture of 13CH4/1602/12 C 1802/H e was made, in which the partial pressures of 13CH 4 and 1602 were kept exactly the same as in the ordinary CH4/O2/He mixture, so as not to disturb the steady-state condition. However, 12C 1802 was added to the isotopic mixture in an amount corresponding to approximately 10-15% of the CO 2 produced during reaction of the CH4/O2/He mixture. The purpose was to measure the production of 13C 180 due to reforming of 13CH4 with 12C1802 only (indirect reaction scheme) under steady-state conditions of the working catalyst surface. Figure 3 shows the transient responses of 13C180 and 12C180 obtained following the second and third switch, respectively, of the following gas delivery sequence: CH4/O2/He (550~

lh) --> 13CH4/O2/C1802/He (550~

(550~ t) which was applied over the Ru/TiO 2 catalyst. It is noted that the added

crt~/o:e 113~/ctSo2 ! ctt4/oeme 500

.

.

.

.

.

.

.

.

.

5 rain) --->CH4/O2/He

.

.

.

12 C 1 8 0 2 (870 ppm) was sufficient to

.

produce a very good 1 3 C 1 8 0 signal to noise ratio in the mass spectrometer under the reaction conditions applied. After approximately 3 min on stream in the

RunaOz T=550oC

375I =

.~.

isotopic mixture of 13CH4/O2/C 1802/He, a r

steady-state value in the rate of 13C 18 0 formation is obtained (Fig. 3). This value is used to estimate the relative contribution of the CO2-reforming route to the overall

125

0

Figure 3. Transient isotopic labelling experiment

production of CO (direct + indirect routes). Proper analysis of this result, taking into consideration scrambling of isotopic oxygen atoms between the CO 2 and 0 2 molecules

for estimating the relative contribution of direct and indirect reaction routes over the Ru/TiO 2 catalyst,

as well as exchange reactions, can give an upper value of the concentration of CO which is derived from the reforming route.

0

5

Time (rain)

10

15

448

If the concentration of 13 C 18 0 produced under 13CH4/1602 / 12 C 1802/H e treatment is Y13C180 (ppm), that of CO and CO 2 produced under CH4/O2/He treatment are YCo and YCO2, respectively, and the concentration of 12C 1802 in the 13CH4/1602/12C 1802/H e feed mixture is yclso~, f then the extent of participation of the reforming route to the formation of CO under CH~O2/He reaction conditions is given by: Z = ( 2 ) ( a ,xy'3clSO y ~ .... x 100(%) 9

[1]

f where ot = Yco2/Yclso2. In Eq. [1 ] the factor of 2 arises from the stoichiometry of the CH4/CO 2 reforming reaction. Table 1 summarizes the results obtained over the Ru/TiO 2 catalyst. The experiments were conducted at 550~ at low and high CH 4 conversions. It is observed that only 8% of the CO produced over the Ru~iO 2 catalyst is through the reforming route, at low CH 4 conversion. As CH 4 conversion is increased to 52%, the Z value increases to 51.2%, indicating that the contribution of the CO 2 and H20 reforming reaction to the production of synthesis gas becomes significant. In the case of Ru/A120 3 catalyst, a similar isotopic experiment performed at 570~ resulted in a value of Z=62% at a low methane conversion. Therefore, it becomes evident from the present isotopic results that the Ru/TiO 2 catalyst largely promotes the direct partial oxidation scheme at low CH 4 conversions as opposed to the Ru/AI20 3 catalyst which promotes the indirect reaction scheme. Table 1 Results of isotopic labeling experiments conducted over the Ru/TiO2 catalyst. T(~ W/F (gmin/cm3)

550 1.9 10-5

550 3.6 10-5

XCH4 (%)

10

52

Sco (%) Yco (ppm) Yco2 (ppm)

13.7 555 3500

62.7 16140 9605

Y13c18o (ppm) Z(%)

5 8

375 51.2

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

. . . . . ....

3.2.2 FTIR studies The nature of surface sites on the reduced 0.5 wt.% Ru/A1203 and 0.5 wt.% Ru/TiO 2

catalysts was probed by FTIR spectroscopy of adsorbed CO. Four adsorbed CO bands were

449 observed over the Ru/A120 3 catalyst as shown in Figure 4. The bands at 2077 and 1857 cm" 1 are assigned to linear and bridged CO respectively, on Ru crystallites. The two bands at 2138 and 2015 cm -1, which were found to appear and develop simultaneously, are ascribed to the dicarbonyl species adsorbed on Ru + sites, i.e. Ru+(CO)2 . Similarly, four adsorbed CO bands were observed over the Ru/TiO 2 catalyst, at 2060, 1850, 2120 and 2014 cm" 1 (Fig. 4). The first two bands correspond to the linear and bridged CO, respectively, while the other two bands are due to the Ru+(CO)2 species. A striking difference in the population of oxidized (Ru +) surface sites between the two catalysts is observed. Apparently, the Ru/AI20 3 catalyst contains a much larger number of Ru + surface sites, as compared to the RuffiO 2 catalyst. The surface species formed on the working catalyst surface were also studies by in situ FTIR. Only one adsorbed CO band at ca 1989 cm-1 was observed on the RuffiO 2 catalyst, presumably due to the bridged adsorbed CO, which tended to be stable after 6 min of reaction time. In contrast, two adsorbed CO bands at ca. 2064 and ca. 1963 cm" 1 were observed over the Ru/A120 3 catalyst, assigned as the 2077

linearly and b r i d g e d CO s p e c i e s , respectively. The intensity of these two bands was found to continuously decrease with time on stream, while a broad band at

i

around 1 7 0 0 - 2 0 0 0 c m ' l gradually developed. This broad band remained intact after Ar purge at 550~ suggesting that it may correspond to formation of a new surface structure different than the metallic Ru structure, probably an oxidized one. The surface state of the spent catalysts was also studied by FTIR of adsorbed CO

/~ 1t

T |0q

II

"o ]120~5

a Ru/AlzO=l(!resh) b Ru/TiOz(fresh)~ c Ru/AlzO~(~pent~ d Ru/TiO2(spent)

2138! k /~ ) 2060 ~

a~

1857

O

2130

following Ar purge at 550~ and cooling to room temperature. Two strong and broad bands were observed at ca. 2130 and 2072 cm -! over the Ru/A120 3 catalyst, assigned as Ru~

and Ru+(CO)2, respectively.

z2b0

2o8o

....

....i 9 6 o

Wavenumber

...................

/

ii2o

e m -l

Over the RuffiO 2 catalyst the Ru+(CO)2 band at ca 2122 c m l is much weaker,while the Ru~

band at 2065 cm" 1 is dominant

(Fig. 4). In both cases there is a shoulder band at lower frequencies of the band at

Figure 4. FTIR spectra of CO adsorbed on fresh and spent RuffiO 2 and Ru/AI20 3 catalysts at room temperature.

2060-2070 cm -I. Apparently, the spent Ru/AI20 3 catalyst contains a significantly larger

450 number of Ru + sites than the spent Ru/TiO 2, indicating that the former catalyst has a higher tendency to be oxidized, as compared to the latter one.

3.2.3 X-ray Absorption Near Edge Structure (XANES) Studies As demonstrated by FTIR (Fig. 4), the unique catalytic performance (i.e. high selectivity to synthesis gas at low methane conversions) of the Ru/TiO 2 catalyst may be related to its surface state which is highly resistant to oxidation under reaction conditions. It was also demonstrated that doping of the TiO 2 carrier with small quantities of altervalent cations can influence significantly selectivity to synthesis gas formation in the presence of oxygen (Fig. 2). In order to determine whether doping of the carrier affects the oxidation affinity of the Ru crystallites, the oxidation state of Ru dispersed on undoped, Zn +2 doped and W6+-doped TiO 2, following different pretreatments, was studied by XANES. During the test period, the X-ray beam intensity was found not to be stable in the LII I edge and thus, the LII edge region was used. Figure 5 shows the LII edge spectra obtained over the three catalysts after reduction in pure H 2 at 300~

for l h followed by oxidation in air at 300~

for lh. The white lines which are seen as sharp absorption peaks immediately above the threshold (i.e. the near continuum) are related to the d-electron vacancies of the absorbing atom. Therefore, the larger the area under the white line, the higher the degree of oxidation of Ru. As shown in Fig. 5, the area under the white line of the Ru/TiO 2 (Zn 2+) catalyst is the smallest. Although the height of the white line of Ru/TiO 2 (undoped) is approximately the same as that of Ru/TiO 2 (W6+), the width of the former is much narrower than that of the latter. The area of the white lines is found to increase in the order: RuffiO 2 (Zn 2+) < Ru/TiO 2 (undoped)
Figure 5. LII edge spectra of Ru dispersed

K-edge spectra of the Ru catalysts were also studied, following different pretreatments. Two sharp bands at ca. 22127.7 and 22152.3 eV (not shown) were observed in the K-edge of the reduced Ru catalysts, as well as Ru metal foil. Upon oxidation, the second band at ca. 22152.3 eV d e c r e a s e d significantly and shifted towards lower energies, while the first band did not change significantly in terms of intensity and position. Therefore, the intensity ratio and separation of these two bands in the Ru K edge region can serve as a relative indication

on undoped, Zn 2+- and W6+-doped TiO 2

of the extent of oxidation, i.e. the higher the

carriers, following exposure to air at 300~ for lh.

intensity ratio of the first to the second band, and the smaller the value of the separation

~.ca rll ~I/b ~II

4 .~..4

~3

...o

0.5~ Ru/T;02

(o) undoped 6 (,b) 0. I o t . ~ W ; (c) 0.3 ot.~; Zn "

.J..4

o Z

0

'20-i0

.... 6 ' 1'13 ........2"0 3"0 "' 4'0 P h o t o n Energy (ev)

451 between the two band maxima, the higher is the extent of oxidation of Ru. The results show that upon oxidation of the reduced catalysts in air at 300~ for 1h, the extent of reduction in the intensity ratio follows the order: Ru/TiO 2 (Zn 2+) < Ru/TiO 2 (undoped) < Ru/TiO 2 (W6+). The separation between the two band maxima is reduced from ca. 24.6 eV (reduced state) to 14.3, 21.8 and 23.2 eV for the Ru/TiO 2 (W6+), Ru/TiO 2 (undoped) and Ru/TiO 2 (Zn 2+) catalysts, respectively. Clearly, both the L- and K-edge X-ray absorption studies demonstrate that the RuffiO 2 (Zn 2+) has the highest resistance to oxidation, whereas the Ru/TiO 2 (W 6+) has the lowest resistance to oxidation. The Ru/TiO 2 (undoped) catalyst lies in between. The relative resistance of the Ru catalysts to oxidation is shown qualitatively in Table 2 along with the selectivity to synthesis gas obtained over the same catalysts at low methane conversion (CH4/O2/He=4/2/94 vol.%, T=750~

It is apparent that a good correlation between

oxidation resistance of the catalysts and selectivity to syngas formation exists, in agreement with the conclusion derived from the FTIR studies. Table 2 Correlation between oxidation resistance and selectivity at low methane conversion. Catalyst Oxidation tendency Selectivity %

Ru/TiO 2 (W 6+) low 40

R u ~ i O 2 (undoped) ~

resistance to oxidation ~ 52

R u ~ i O 2 (Zn 2+) high 67

3.3 Long -term stability and practical considerations of the RufHO 2 catalyst The long-term stability of the Ru/TiO 2 catalyst was studied under various reaction conditions and the spent catalysts were characterized for assessing the reasons of deactivation. It was observed that the rate exhibits a rapid reduction at the initial several hours of reaction, followed by a slow and continuous deactivation. Analysis of the spent catalyst, by H 2 adsorption after removing surface carbon, mainly due to metal sintering, while the occurrence of the SMSI phenomenon at conditions prevail. In order to avoid these

showed that the initial rapid reduction of activity is continuous and slow deactivation is related to the the later part of the catalyst bed, where reducing processes which lead to catalyst deactivation, TiO 2

was pre-sintered at temperatures higher than 900~ for several hours. It was found that at oxygen conversions less than 100%, Ru supported on sintered (futile) TiO 2 exhibits a stable performance for a period of at least 100 h, at a temperature of 800~ It is estimated that the heat produced at the entrance of the catalyst bed is reduced by approximately 80% when the present Ru/TiO 2 catalyst is employed, due to suppression of combustion reactions. It should be stated that the quantity of heat produced depends on conversion and selectivity. Thus, regardless of the reaction scheme which is followed, if high

452 selectivity to synthesis gas is achieved at low CH 4 conversions (Xo2 < 100%), the quantity of heat released near the entrance of the catalyst bed will be small. Approximately 70% of methane in the feed (CH4/O2=2) can be converted to synthesis gas at oxygen conversions less than 100%. The remaining methane can be further converted to synthesis gas either by subsequent reforming reactions or by partial oxidation upon introduction of additional amounts of oxygen. 4. CONCLUSIONS The following conclusions can be drawn from the results of the present investigation: 1. The Ru/TiO 2 catalyst is unique, among many catalysts investigated, in promoting the direct partial oxidation of methane. 2. The extent of the direct partial route can be enhanced by appropriate doping of the TiO 2 carrier with cations of lower valence (i.e. Li +, Ca 2+, Zn 2+, La3+). 3. The unique ability of the RuffiO 2 catalyst is probably related to the enhanced resistance of the metal surface to be oxidized. ACKNOWLEDGEMENT Financial support by the Commission of the European Community (contract no. JOU2-0073) is gratefully acknowledged.

REFERENCF~ 1. J.R. Rostrup-Nielsen, "Catalytic Steam Reforming", Catalysis, Science and Technology, vol. 5, p.1, J.R. Anderson and M. Boudart, (Eds.), Springer, Berlin, 1984. 2. A.T. Ashcroft et al., Nature 344 (1990) 319. 3. D.Dissanayake, M.P. Rosynek, K.C.C. Kharas and J.H. Lunsford, J. Catal. 132 (1991) 117. 4. D.A. Hickman and L.D. Schmidt, J. Catal. 138 (1992) 267. 5. P.M. Torniainen, X. Chu, and L.D. Schmidt, J. Catal. 146 (1994) 1. 6. H. Matsumoto, and S. Tanabe, J. Chem. Soc., Faraday Trans. 90 (1994) 3001. 7. D. Dissanayake, M.P. Rosynek, K.C.C. Kharas, and J.H. Lunsford, J. Phys. Chem. 97 (1993) 3644. 8. Y. Chang and H. Heinemann, Catal. Lett. 21 (1993) 215. 9. Y. Boucouvalas, Z.L. Zhang, and X.E. Verykios, Catal. Lett. 27 (1994) 131. 10.T. Ioannides and X.E. Verykios, J. Catal. 145 (1994) 479. 11 .A.M. Efstathiou, D. Papageorgiou and X.E. Verykios, J. Catal. 141 (1993) 612.

J.W. Hightower, W.N. Delgass, E. lglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis -40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

453

The Oxidative Transformation of Methane over the Nickel-based Catalysts Modified by Alkali Metal Oxide and Rare Earth Metal Oxide Qing Miao, Guoxing Xiong*, Shishan Sheng, Wei Cui, Xiexian Guo (State Key Laboratory of Catalysis, Dalian Institute of Chemical Physics, Chinese Academy of Sciences, P.O.Box 110, Dalian 116023, P.R.China)

Abstract Two completely different behaviors of the oxidative transformation of methane were performed over the nickel-based catalysts because of the different modifications by alkali metal oxide and rare earth metal oxide and the different interactions between nickel and supports, and two types of catalysts, namely the LiNiLaOx catalyst with a good Oxidative Coupling of Methane (OCM) performance and the LiNiLaOx/Al203 supported catalyst with an excellent performance of the Partial Oxidation of Methane to Syngas (POM) reaction were obtained. Several techniques, such as flow-reaction, pulse-reaction, XRD, H2-TPK , XPS, TPO, and TG, etc., were employed to investigate the relation among the preparation and composition of catalysts, the structures of catalysts and the catalytic performances, especially effects of each component, the active phases and their precursors, the redox behaviors and the states of nickel present in those nickel-based catalysts. The effects of acid-base properties on the states of nickel present and on the directions of the oxidative transformation of methane, the interaction between nickel and other components and the deposition of surface carbon over catalysts were studied. The types of active centers, the modes of the activation of methane and the reaction mechanisms were discussed in detail

1.

INTRODUCTION

The large m o u n t s of natural gas (mainly methane) found worldwide have led to extensive research program~ in the area of the direct conversion of methane [1-3]. The oxidative transformation of methane (OTM) is an important route for the effective utilization of the abundant natural gas resources. How to increase catalyst activity is a common problem on the activation of methane. The oxidation of methane over transition metal oxides is always high active, but its main product is CO2, namely the product of deep oxidation. It is because transition metal oxides have high oxidative activity. So, they were usually used as the main component of catalysts for the complete oxidation of alkane[4]. The strong oxidative activity of CI-I4 over transition metal oxides such as NiO indicates that the activation of C-H bond over transition metal oxides is much easier than that over alkaline earth metal oxides and rare earth metal oxides. Furthermore, the activation of C-H bond is the key step of OTM reaction. It is the reason that we use transition metal oxides as the main component of the OTM catalysts. However, we have to realize that the selectivity of OTM over transition metal oxides is poor.

454 We expected to control the direction of OTM reaction over NiO by surface modification, namely making use of the interaction between NiO and other components to beget a synergistic effect. In this paper, two completely different behaviors of the oxidative transformation of methane were performed over the nickel-based catalysts because of the different modifications by alkali metal oxide and rare earth metal oxide and the different interactions between nickel and supports. Furthermore, the two completely different reactions were related with the acid-base properties of catalysts and the states of nickel present.

2.

EXPERIMENT

2.1. Preparation of Catalysts The LiNiLaOx catalysts were prepared by impregnation method, namely impregnating appropriate amounts ofLiNO 3 and Ni(NO3) 2 on La20 3 for 24 hr., and dried at 393K and then calcined in air at 823-1173K for 4 hours. The LiNiLaOx supported catalysts were also prepared by impregnation method, namely impregnating appropriate amounts of LiNO3, Ni(NO3) 2 and La(NO3) 3 on different supports (AI20 3, SiO2, MgO) for 24 hr., and dried at 393K and then calcined in air at 823-1073K for 4 hours. 2.2. Test of Catalytic Performance Catalysts were tested by the flow-reaction in an atmosphere pressure fixed-bed microreactor. Products of the reaction were analyzed by Gas Chromatography using a TCD detector. The conversion of methane, the selectivity and yield of the products were calculated on the basis of carbon numbers of the methane reacted. Pulse reactions were performed in a fixed-bed microreactor that was changed by combining the reactor and chromatography column on-line. Products of the reaction were analyzed by Gas Chromatography using a TCD detector. Weight of catalysts is 0.1 g. 2.3. Characterizations of Catalysts H2-TPR profiles were recorded by an on-line computer at a programmed temperature velocity of 14 K/rain in 5% H2/Ar flow after ~mples were pretreated in Ar flow at 873K for 30 rain When the programmed temperature reached 1173K, it was stopped and held for 30 rain. XRD characterization was performed with a Riguku D/Max-RB X-ray diffractometer using a copper target at 40kV x 100mA and scanning speed of 8 degree/rain.

3.

RESULTS. AND DISCUSSION

3.1 Effects of Li content on the catalytic behaviors and structures of LiNiLaOx catalysts The dependence of performance of LiNiLaOx catalysts on Li content at 1073K was shown in Fig. 1. When Li/Ni mole ratio was 0, the relatively acidic LaNiOx had the highest CH 4 conversion(92.0%), but no C 2 yielded. The products were CO, CO 2 and H 2, and CO selectivity was 98.3%~ It is not an OCM catalyst but a good catalyst for partial oxidation of methane(POM). With Li content and the basic property of LiNiLaOx catalysts increasing, CH4 conversion and CO selectivity decreased, but there was still no C 2 formed until Li/Ni mole ratio was 0.4. There was a tumpoint of catalytic behavior between 0.2 and 0.4 (Li/Ni mole

455 ratio). Beyond Li/Ni mole ratio of 0.4, CH4 conversion, C 2 selectivity and C2-/C 2 increased with Li content increasing, and CO selectivity was almost zero. There was an optimum C 2 yield at the Li/Ni mole ratio of 1.6, namely CH 4 conversion was 34.7%, and C 2 selectivity was 60.1% and CO selectivity was 0. Those indicate that Li content and the acid-base properties of catalysts result in the radical change of the behaviors of the oxidative transformation of methane over LiNiLaOx catalysts. CPS 20K

100

O L)

80

.~

60

Li/Ni

1.5

1

o4

2.0

40 0.5

20

1.6 1.2

m r~

0

.= 0

u =, u == .= 0.1 0 . 2 0 . 4 0 . 6 0 . g

== w == ~, 1.0 1.2 1.6 2 . 0

Li Coat~ (U/tqi)

1.0 0.8

C H 4 C o n y . C 2 S d C O Sol C 2 - - / C 2 --o-

--o-

~

Fig. l Dependenceof performance of catalysts on Li content at I073K PCH4:Po2:PHe--3:I:4, W/F=0.083 g-h/l

0.6 0.4 0.2

___ 0.1 The XRD tests were performed to . 0.0 determine the crystal phases of those 1 samples(see Fig.2). When Li/Ni mole ratio 20 40 60 65 20 was 0, the dominate crystal phase was Fig.2 XRD spectra of LiNiLaOx catalysts with LaNiO 3. With Li addition, the peak different Li content intensity of crystal phases increased and the dominate crystal phase was changed to La2Nil_yLiyO4. x (d: 0.2823, 0.2652, 0.1563 nm). When Li/Ni mole ratio was 0.4, its dominate phase was La2Nil_yLiyO4_x. Beyond the Li/Ni mole ratio of 0.8, the crystal phases of LiNiLaOx catalysts were almost completely La2Nil.yLiyO4.~. Other results[5] demonstrate that the La2Nil.yI_2yO4.x crystal phase with oxygen vacancy is the active phase of the LiNiLaOx catalyst for OCM reaction and its lattice oxygen may be the active oxygen species. CH 4 was activated by the active oxygen species and produced CH 3- flee radial. Yu and his co-workers also think that the La2Nil.yLiyO4_x crystal phase is the active phase of LiNiLaOx catalysts [6]. It is shown that the Li addition affects the formation and structure of crystal phases of LiNiLaOx catalysts. It is the difference of crystal phases that results in the different behaviors of the oxidative transformation of methane. 3.2. Effects of Li c o n t e n t on the redox ability and the state of nickel p r e s e n t in

LiNiLaOx catalysts The states of nickel present in the different LiNiLaOx catalysts with differenf Li content were investigated by H2-TPR techniques (see Fig. 3). With Li content increasing, the states of LiNiLaOx catalysts present were changed. The more Li content and the stronger the basic

456 property of LiNiLaOx catalysts, the more difficult to be reduced by H 2. An obvious high temperature reducing peak appeared according to the H2-TPR profiles of LiNiLaOx (Li/Ni=l.6) catalyst. That high temperature reducing peak is attn'buted to the existence of fixed-form NiO, which may be relate to the La2Nil_yLiyO4_~.crystal phase.

I

I

1 1 ; 3 1 ; 3 1

473 573 673 773 8 973 10 1173 Temperature (K) Fig.3 H2-TPR profiles of LiNiLaOx catalysts with different Li content (a) LaNiOx; (b) Li/Ni=0.2; (c) Li/Ni=0.4; (d) Li/Ni=l.6

I

I

l

I

i

J

t

U

473 573 673 773 873 973 1073 1173 Temperature (K) Fig.4 H2-TPRprofiles of LaNiOx catalyst after 0 2 pretreatmmt (a)without 0 2 pretreatment after 0 2 pr~reatment for 30 mm. at (b) 323K; (c)423K; (d)573K; (e) 773K

Because of the co-existence of reductive agent CI-I4 and oxidizing agent 02, there was an oxidation-reduction reaction over the surface of L~iLaOx catalysts. In order to imitate the redox process in the oxidative transformation of methane and compare the redox ability of those catalysts, the LaNiOx and LiNiLaOx(Li/Ni= 1.6) catalysts under 0 2 pretreatment for 30 rain at different temperature were investigated by H2-TPR technique. The procedure of the experiment is that the routine TPR characterization was performed at first and the samples were completely reduced, and then those reduced samples under 0 2 pretreatment for 30 rain. at different temperature were investigated by TPR technique with the routine procedure. The results show that the reduced LaNiOx catalyst under 0 2 pretreatment can produce crystal NiO that is easily reduced(see Fig.4). The reduced LiNiLaOx catalyst under the same 02 treatment can produce the well-dispersed NiO and fixed-form NiO that are more difficult to be reduced than crystal NiO(see Fig.5). Those indicate that the LaNiOx catalyst is more easily reduced under the same redox reaction condition than the LiNiLaOx catalyst(Li/Ni= 1.6). To observe the state of nickel present in worlcing catalysts under reaction condition, the LaNiOx and LiNiLaOx catalysts after reaction were investigated by H2-TPR techniques. The H2-TPR results of the working LaNiOx and LiNiLaOx catalysts(see Fig.6) indicated that the nickel in the worlcing LaNiOx sample existed in the reduced nickel, but the nickel in the working LiNiLaOx sample existed in the oxidized nickel and had two kinds of existing states, namely well-dispersed NiO and fixed-form NiO. The acidic property favors keeping the reduced nickel and the basic property favors keeping the oxidized nickel The reduced nickel is necessary for the partial oxidation of methane to syngas(POM) and the oxidized nickel is

457 necessary for the oxidative coupling of methane(OCM). The effects of Li20 on the performance of the LiNiLaOx catalysts include two aspects: (1) the addition of Li20 enhances the basic property of the catalyst and makes nickel difficult to be reduced, and keeps it in the oxidized state; (2) the addition of Li20 favors the formation of La2Nil.yLiyO4_~ crystal phase and oxygen vacancy, and enhances the mobility of oxygen anion and C 2 selectivity.

CI

e

d c

~b

b I

1

I

I

I

I

I

t

473 573 673 773 873 973 1073 1173 Temperature (K) Fig.5 H2-TPR profiles of LiNiLaOx (Li/Ni=I.6) catalyst atter 0 2 pretreatment (a)without 0 2 pretreatment after 0 2 pretreatmm3t for 30 mm. at (b) 323K; (c)423K; (d)573K; (e) 773K

cl 473 573 673 773 873 973 1073 1173 Temperature (K) Fig.6 H2-TPR profiles ofLiNiLaOx (Li~i=l.6) and LaNiOx catalysts after reaction (a) LaNiOx; (b) LiNSLaOx(Li/Ni=l.6) I

!

t

l

,1

I

!

f

3.3. Effects of supports on the of LiNiLaOx supported catalysts When the LiNiLaOx sample was supported on the supports with different add-base properties, different behaviors of OTM were performed (see Table 1). When the L~iLaOx sample was supported on the acidic support A1203, it was changed to an excellent POM catalyst. However, when it was supported on the basic support MgO or much weaker acidic support SiO2, both CO and C 2 were produced over those two supported nickePbased catalysts. It was shown that the acid-base properties of the nickel-based catalysts resulted in the radical change of the behaviors of the OTM reaction. XRD determination of those nickelbased catalysts indicated that NiO was well-dispersed on the A1203 support [7]. It also was shown by the TPR results that nickel existed in a uniform state over the Al203 support and the high temperature reducing peak was due to the formation of NiAI204, which was assured by XPS results [8]. Those indicate that there exists a strong interaction between LiNiLaOx and AI203, and the strong interaction results in the formation of NiAI204. This metal-support interaction improves the dispersity of active component nickel and the high dispersion of the nickel over AI203 support is beneficial to obtain an excellent performance for POM reaction. Ross [9] also pointed out that the unreduced NiO/AI203 catalysts used in the steam reforming reaction contained surface nickel aluminate phases which, on reduction, gave monodispersed nickel atoms closely associated with alumina sites in addition to metallic crystallites arising from the reduction of nickel oxide, and the monodispersed nickel atoms probably participated in the CH 4 + H20 reaction. Geus and his co-workers [ 10] studied the adsorption of methane

458 on silica-supported nickel catalysts using a low-field magnetic method and infrared spectroscopy. In the entire temperature range (303K
~. Ni3C (" surface nickel-carbide") + 4Ni-H

It was observed that per unit surface area, ~qrnoll nickel crystallites were more reactive toward methane than were large crystallites. Table 1 Performance and XKD determination of the different nickel-based,,, catalysts Sample CH 4 Conversion C 2 Selectivity CO Selectivity XRD .......... (%) ........ (%) (%) ,, determination LiNiLaOx LiNiLaOx/Al20 3 LiNiLaOx/SiO 2 LiNiLaOx/MgO

43.0 91.5 36.8 41.9

47.4 0.0 57.1 24.1

3.5 99.2 38.2 19.2

La2Ni]-yLiyO4-k Y-AI20 3 SiO2(Li), NiO MgO, MgNiO,NiO La2Nil_yLiyO4.~.

CH4 " 0 2 = 2 " 1, GHSV = 2.7 x 104 Vkg.hr, reaction temperature = 1123K The redox behaviors and the states of nickel present in the working LiNiLaOx/Al20 3 catalysts were investigated by H2-TPR techniques. The result of LiNiLaOx/Al20 3 (Fig.7, Fig.8) is similar to that of LaNiOx catalysts (see Fig.4, Fig.6) and contrary to that of LiNiLaOx (see Fig.5, Fig.6). The results indicated that the reduced LiNiLaOx/Al20 3 sample under 0 2 pretreatment also produced crystal NiO that was easily reduced (Fig.7), and the nickel of the working LiNiLaOx/Al20 3 catalyst existed in the reduced nickel (Fig.8). In a word, the acidic property favors keeping the reduced nickel and the basic property favors keeping the oxidized nickel The reduced nickel is necessary for POM reaction and the oxidized nickel is necessary for OCM reaction.

d

C b

/

/~

--J -J i ..i.. , , , i i , 473 573 673 773 873 973 1073 1173 Temperature (K) Fig7 H2-TPR profiles of LiNiLaOx/A120 3 cat~yst after 0 2 pretreatment (a)without 0 2 pretreatm~nt after 02 pretreatmerrt for 30 mm. at (b) 323K; (c) 423K; (d) 573K; (e) 773K Qi

9

L

I

i

........

i ......

l

-

I

=

713

I

b

473 573 673 773 873 973 10 1173 Temperature (K) Fig.8 H2-TPR profiles of LiNiLaOx/A120 3 and LiNiLaOx catalysts after reaction (a) LiNiLaOx; (b) LiNiLaOx/Al203

459 3.4. POM performance of LiNiLaOrJAI203 catalysts The problems of the loss or sintering of nickel and of the deposition of surface carbon over the catalyst are the main factors that result in the deactivation of the NiO/AI203 catalyst[ll,12]. The adoption of rare earth metal oxide and alkali metal oxide successfidly solved those problems. The LiNiLaOx/A1203 catalyst has a lower nickel content, excellent POM performance (CH 4 conversion of 94.8% and CO selectivity of 98.1%), excellent stability and carbon-deposition resistance. Under the reaction condition that the reaction temperature is 1123K and CH4/O 2 ratio is 1.96, and GHSV is 2.7 x 104 l/kg-hr, the catalyst passed 50 hr. liR test and kept its excellent performance (CH 4 conversion > 96.5% and CO selectivity > 95%). Yu et al also lind that the adoption of rare earth metal oxides can improve the stability of NiO/A1203 catalysts for the steam reforming reaction of methane[ 13]. The most important advantage ofLiNiLaOx/A1203 is its carbon-deposition resistance. TG, TPO and XPS results of ~ _topics after the reactions during 10 hr. indicate that the carbondeposition resistant ability of LiNiLaOx/A1203 is much better than that of NiO/A1203 [8]. According to the TG results, the weight of surface carbon over NiO/AI203 is 12 percent of its net weight. However, the weight of surface carbon over LiNiLaOx/A1203 is only 0.2 percent (nearly to none) of its net weight. Because the basic oxides, such as alkali metal oxides and rare earth metal oxides, modify the stronger acidic centers of the AI203 support that favor the carbon deposition[ 14,15], the trend of carbon-deposition over the NiO/A1203 catalyst can be inh~ited. It can be concluded that the adoption of rare earth metal oxide and alkali metal oxide improve the carbon-deposition resistance of the supported nickel-based catalysts. 3.5. Study of the active site and reaction mechanism of POM reaction The two mechanLqrns proposed to account for the partial oxidation of methane to syngas may be designated as the IPO (Indirect Partial Oxidation) mechanism and the DPO (Direct Partial Oxidation) mechanism. The IPO mechanL,ma was proposed by Prette et al [16] and Lunsford et al [12]. They think CO and H 2 are the products of indirect reaction, the overall reaction of the POM reaction is composed of three different reactions CH 4 + 2 0 2 CH 4 + C O 2 CH 4 + H 2 0

-~ CO 2 + 2 H 2 0 -~ 2 C O + 2 H 2 -~ C O + 3 H 2

(1) (2) (3)

However, Schmidt [17] and Lapszewiez [18] claimed that CO and H 2 are the primary products of POM reaction. The dissociative adsorption of CH 4 on LiNiLaOx/Al203 produces the adsorbed hydrogen( [H] ad) and the surface carbon([C]ad). The surface carbon reacts with surface oxygen species to produce CO, and the adsorbed hydrogen was desorbed to form H 2. Although there are two different proposals about the POM reaction mechanism, the reduced metal (Ni, Co, Fe, Pt, Pd, Rh, etc) is generally thought to be the active center of POM reaction [12,16-18]. According to the above-mentioned results (see Fig.6, Fig.8), the reduced nickel is also the active center of the LiNiLaOx/A1203 catalysts for POM reaction. In order to study the mechanism of POM reaction, a series of pulse reactions were performed. The results of CH 4 pulse reaction on the fresh LiNiLaOx/Al203 sample were shown in Fig.9. The products of the CH 4 pulse reaction were H2, CO and CO 2. With pulse

460 number increasing, CH4 conversion reached a maximum value and then decreased drastically. CO selectivity increased with pulse number increasing and reached 100% after the third CH 4 pulse. Those indicated that the NiO on the surface of L~iLaOx/A120 3 san~le was gradually reduced and the adsorbed oxygen species over the sample surface were consumed by CH 4 and H 2 that was the product of the CH4 pulse reaction with CH 4 pulse number increasing. So, CO selectivity was improved to 100%. Because CH 4 can dissociativelly adsorb on the nickel to produce the surface carbon [10,19,20], the carbon-deposition over the ~ c e of the ~n~ple was aggravated with CH 4 pulse number increasing and the active centers of the ~ l e gradually deactivated. Therefore, CH 4 conversion and H 2 yield decreased rapidly. To demonstrate the formation of the surface carbon during the CH4 pulse reaction, the 0 2 pulse reaction was performed immediately after the CH 4 pulse reaction. CO was produced at the selectivity of 96.6% and 0 2 was completely consumed at the first O 2 pulse, but only a little of CO 2 was produced at the second 0 2 pulse. Those finding indicated that the surface carbon was formed during the CH 4 pulse reaction and the surface carbon reacted with relatively low concentration O2(compared with the mount of surface carbon) to produce CO. At the second 0 2 pulse, because the amount of the sm'face carbon was much lower than that of 0 2, the main product was CO 2. Those demonstrate that CH 4 dissociatively adsorbs on the LiNiLaOx/Al20 3 catalyst and produces the reaction intermediate, surface carbon. CO selectivity depends on the relative concentration of oxygen. The surface carbon reacts with relatively low concentration 0 2 to produce CO and with relatively high concentration 0 2 to produce CO 2. 60

120 5:100 ~o

80

40

60

30

40

20 . ~

~

120

~ .o

I00

80

-u

0

~8

o= 9

~ 2o ~ r..)

0

10 ~ -0

1

2

3

4

5

6

7

C H 4 Pulse number

Fig.9

C H 4 Conv. C O Sel. H 2 Yield -41.+ --o--

Results of CH4 pulse reaction on the LiNiLaOx/Al203 catalyst at 1123K

r~ 0 r~ 0

60 40 20

o

-.1

v

i-

.-

2

3

4

.......,., em 5

n~ -. 6

0 2 I ~ s e number 0 2 CO2 CO -.-

-El-

-o-

Fig. 10 Restflts of 02 pulse reaction on the LiNtLaOx/Al203 catalyst at 1123K

To distinguish the different actions of different oxygen species over the LiNiLaOx/A1203 on the CO.selectivity, the sample was pretreated by 5% H2/Ar flow at 1123K for 0.5 hr. ARer the pretreatment, the NiO was reduced to the reduced nickel and the surface adsorbed oxygen species was completely consumed, and then the CH 4 pulse reaction was performed. The results are different from the results of CH 4 pulse reaction on the fresh ~mple. The

461 products of the CH 4 pulse reaction are only CO and H2, and no CO 2 appears. It can be inferred ,-e. 100 from those finding that there exist two kinds of 0 8O 40 oxygen species over the LiNiLaOx/Al20 3 catalyst, namely the surface lattice oxygen ([O]lattice) that 30 ''~ ~ 60 is difficult to be depleted by H 2 and the surface 20 .~ adsorbed oxygen([O2]ad) that is easy to be depleted by H 2. The surface carbon that is the "~" 20 product of the dissociative adsorption of CH 4 -0 0 reacts with the surface lattice oxygen on the 1 2 3 4 5 CH4 Pulse number catalyst to produce CO and reacts with surface adsorbed oxygen on the catalyst to produce CO 2. CH4 Cony. CO Sel. H2 Yield The CO selectivity depends on the relative Fig. 11 Results of CH4 pulse reaction on concentration of the two surface active oxygen the reduced LiNiLaOx/Al203 catalyst species. Furthermore, the relative concentration at 1123K of the two surface active oxygen species is determined by 0 2 concentration in gas phase, the velocity of the formation of the surface adsorbed oxygen and the transformation from the surface adsorbed oxygen species to surface lattice oxygen species. The study on the mechanism of the formation of different oxygen species is under way. 120

In a word, POM and OCM reactions proceed at different active sites with different reaction intermediates and reaction mechanisms over the nickel-based catalysts proposed as follows: LiNiLaOx i

I.aNiOx LiNiLaOx/AhO3 i2 under reaction surface catalyst condition T [o,~],~ ~ co2 oxygen species on N_~ [C],o + ~ ~ catalyst surface C2 ~

CH3 9 L ~ i i - y L i y O 4 & p h a s e

CH4

[O]m

4[I-I]~ ~ oxidative coupling of methane

~

CO

214_2

partial oxidation of methane to syngas

4. CONCLUSIONS Two completely different behaviors of oxidative transformation of methane, namely the Oxidative Coupling of Methane to C 2 Hydrocarbons(OCM) and the Partial Oxidation of Methane to Syngas(POM), were performed and related over the nickel-based catalysts due to different modification and different supports. It is concluded that the acidic property favors keeping thereduced nickel and the reduced nickel is necessary for POM reaction, and the basic property favors keeping the oxidized nickel and the oxidized nickel is necessary for OCM reaction. POM and OCM reactions proceed at different active sites caused by different

462 configuration with different reaction intermediates and reaction mechanisms over the nickelbased catalysts. The NiO/A1203 catalysts modified by alkali metal oxide and rare-earth metal oxide have excellent POM reaction performance, excellent stability and carbon-deposition resistance.

ACKNOWI~DGEMENT The financial support of the National Natural Science Foundation of China is gratefully acknowledged.

REFERENCES

[ 1] 1L Pitchai, K.Klier, CataL Rev.- ScL Eng., 28(1) (1986) 13 [2] S.C.Tsang, J.B.Claridge, M.L.I-I_C~een, CataL Today, 23 (1995) 3 [3] A.MMaitra, AppL CataL, A104 (1993) 11 [4] T.Seiyama, CataLRev.-SckEng., 34(4) (1992) 281 [5] Qing Miao, Guoxing Xiong, J.NatuLral Gas Chemistry (submit) [6] S.-Y. Yuan, B.-C. Liu, and Z.-L. Yu, J.Natural Gas Chemistry, 2 (1992) 166 [7] Y. Xie and Y. Tang, Adv. in CataL, 37 (1990) 1 [8] Q Miao, Ph.D. dissertation, Dalian Institute of Chemical Physics, Dalian, China, 1995 [9] J.1LH.Ross, M.C.F.Steel and A.Zeini-Isfahani, J.CataL, 52 (1978) 280 [10] E.G.M.Kuijpers, A.K.Breedijk and J.W.Genus, J.Catal., 72 (1981) 210 [l l] P.M.Tomiainen, X.Chu and UD.Schmidt, J.Catal., 146 (1994) 1 [12]DDissanayake, M.P.Rosynek, K.C.C.Kharas and J.I-LLunsford, J.Catal., 132(1991) 117 [13]J.-Y.Cheng, I-L-G.Yu, J.Natural Gas Chem., No.2 (1992) 129 [14] D.L.Trimm, Design of Industrial Catalysts, Elsevier Scientific Publi~ing Company, Amgerdam-Oxford-New York, 1980. [15] S.1LdeMiguel, O.A.Scelza and A.A.Gastro, Topics in Catalysis, 1 (1994) 87 [16] M.Prettre, C.Eichner, and M.Perrin, Trans. Faraday Soc., 43 (1946) 335 [17] D.A.Hickman and L.D.Schmidt, Science, 259 (1993) 343 [18] J.A.Lapszewicz, X.-Z.J'mng, Prepr.-Am. Chem. Soc. Div. Pet. Chem.,38(4) (1993) 815 [19] C.KembaU, Advan. Catal, 11 (1959) 223 [20] E.G.M.Kuijpers, J.W.Jansen, A.J.Van Dil]en, and J.W.Geus, J. CataL, 72(1981) 75

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis -40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

463

Design of stable catalysts for methane-carbon dioxide reforming J.A. Lercher, J.H. Bitter, W. Hally, W. Niessen, and K. Seshan Faculty of Chemical Technology, Catalytic Processes and Materials Group, University of Twente, P.O. Box 217, 7500 AE, Enschede, The Netherlands

Abstract

The activity and stability of catalysts for methane-carbon dioxide reforming depend subtly upon the support and the active metal. Methane decomposes to carbon and hydrogen, forming carbon on the oxide support and the metal. Carbon on the metal is reactive and can be oxidized to CO by oxygen from dissociatively adsorbed CO2. For noble metals this reaction is fast, leading to low coke accumulation on the metal particles. The rate of carbon formation on the support is proportional to the concentration of Lewis acid sites. This carbon is non reactive and may cover the Pt particles causing catalyst deactivation. Hence, the combination of Pt with a support low in acid sites, such as ZrOz is well suited for long term stable operation. For non-noble metals such as Ni, the rate of CH~ dissociation exceeds the rate of oxidation drastically and carbon forms rapidly on the metal in the form of filaments. The rate of carbon filament formation is proportional to the particle size of Ni. Below a critical Ni particle size (d<2 nm), formation of carbon slowed down dramatically. Well dispersed Ni supported on ZrO2 is thus a viable alternative to the noble metal based materials.

I. INTRODUCTION Carbon dioxide reforming of methane to produce synthesis gas, i.e., a mixture of carbon monoxide and hydrogen (CO2+CH4 ~ 2CO+2H2; AH~ = 261.0 kJ/mol) has attracted substantial interest [ 1-4]. The reaction route is well suited to produce CO rich syngas or very pure carbon monoxide for the synthesis of bulk chemicals such as acetic acid, dimethyl ether and alcohols via the oxoalcohols synthesis [5-7]. More significantly, for acetic acid manufacture carbon dioxide reforming is estimated to have economical advantages over other syngas production routes [6]. The reaction contains similar elementary reaction steps as in steam reforming ( H 2 0 + C H 4 C O + 3 H 2 , AH~ = +206.2 kJ/mol) [8,9], but the absence of water and the higher C/H ratio in the reactant feed favors coke formation [ 10]. Minimization of coking rates is therefore one of the key aspects for designing a stable catalyst for the reaction. Coke forms readily via methane decomposition (CH4 ~ C + 2 H 2, AH~ = +74.9 kJ/mol) and CO disproportionation (2CO ~ C +CO~, AH~ =-172.4 kJ/mol) [ 11,12]. Options to reduce the coke build up are (i) the addition of water (coupling with steam reforming) [ 10], (ii) the addition of oxygen (coupling with partial oxidation) [7], or (iii) the use of catalysts which minimize the rate of coking [1 ]. We will report here on the last aspect, i.e., on the successful design of high temperature Pt and

464 Ni catalysts using Z r O 2 as a unique support that seems crucial to minimize coking under reaction conditions applied for CH4/CO2 reforming. For two successfully developed catalysts, (Pt and Ni on ZrO2) the present contribution outlines the sequence of the elementary steps and the catalytic chemistry of the active metal and the support in order to explain catalysts activity and stability.

2. EXPERIMENTAL

2.1 Catalytic materials and reactants Titania was obtained from Degussa (P25) and consisted of a mixture of anatase and rutile. ZrO2 (RC 100, Daichi Kigenso, Japan) contained only the monoclinic phase, y-A1203 was obtained from Akzo-Nobel Chemicals (type 000-3AQ). The oxides were pressed into pellets and crushed to 0.30.6mm grains. These grains were then calcined at 1125K for 15 hours in a stream of dry flowing air (30 ml/min). The catalysts were typically prepared by impregnation of the calcined supports with aqueous 0.1 molar solutions of H2PtCl6 or Ni(NOs)2. Subsequently, the catalyst precursors were dried at 395K for 8 hours and calcined at 925K for 15 hours in flowing air (30 ml/min). 2.2 Catalytic experiments Catalyst testing was performed in fixed bed (300 mg) continuous flow reactors (using a C~ to analyze the products). The feed had a composition of CH4/CO_JHe =I:I :2 and was passed over the catalyst with a space velocity of 28000 hx (GHSV). Prior to testing, the catalysts were typically reduced m situ with 5% H2 in N2 for one hour at 1125K in the case of Pt catalysts and at 875K in the case of Ni/ZrO2. Coke formed on the catalyst was determined by combustion in oxygen and measuring the amount of carbon oxides generated [ 13,14]. 2.3 Catalyst characterization The reaction on the catalyst surface was followed by in situ i.r. spectroscopy using a Bruker IFS88 FTIR spectrometer for the characterisation of sorbed species and mass spectroscopy for the analysis of gas phase. The state of Pt was further investigated by m situ X-ray absorption spectroscopy (Daresbury, UK, beamline 9.1, transmission mode, Si(220) monochromator, Pt-L m edge). Details of catalyst characterisation techniques are reported elsewhere [ 13,14].

3. RESULTS

3.1 Platina based catalysts The characteristics of the supported Pt catalysts used in this study are compiled in Table l. The activity of the catalysts during a typical time-on-stream experiment are shown in Fig 1. While Pt/ZrO 2 catalyst showed high stability (less than 5% loss in conversion during 180 hours time on stream), Pt/y-Al203 lost all catalytic activity after 3 hours and Ptfrio2 deactivated gradually during 40 hours. The supports alone were catalytically inactive. The amount of coke formed as a function of the number of turnovers is shown in Fig.2. The steeper slopes of these curves for Pt/y-Al203 and Pt/TiO2 indicate the higher selectivity of Pt/yA1203 and Pt/TiO2 to form coke than Pt/ZrO2. Hydrogen chemisorption capacity decreased markedly after some time on stream (see Table 2), but could be completely restored by oxidative treatment.

465

I

"~~

100.00

9zro2 .... 9T~O2

E

'm

1

A A~~ ' ~ . ~ .~. . . . A

0

10

=_

20

,

,,

30

9

40

10.00

~ ;

1.~

~. r

0.10

;

O.Ol

9AI203

@

r~l

o

0.01

i

0.1

t

1

!

10

'

100

COz TON I m a e s )

Time (hours)

Figure 2 Amounts of carbon formed as a function of the total amount of CO: converted for different Pt catalysts during C02/CH4 reforming at 875K

Figure 1 Stability of different supported Pt catalysts for COJCH4 reforming at 875K

4T _

10-i

I

8~-

i i

6

3t m -

~2T 1+

0

2~

5

10

15

Amount of C formed ('10"~moles.g "1 c a t a l y s t )

Figure 3 Catalytic activity of Pt/y-Al:O~ as function of the carbon deposited on the catalyst (875K)

20

lOOO

Temperature (K)

Figure 4 Temperature programmed desorption of pyridine ; pyridine adsorbed at 4 75K, heating rate 20K/min

466 Table 1 Physicochemical properties of Pt catalysts investigated Metal loading (wt%)

BET Area (m2.g1)

Pt dispersion"

y-Al203

0.5+0.2

112

35

TiO2

0.5+0.2

5

25

ZrO 2

0.5_+0.1

16

33

Support

(%)

* catalysts were reduced at 1125K, Pt/H=I While these observations link coking, blocking of active sites and catalytic activity, the results do not indicate to what extent the support contributes to coking and if such coke may physically block the access to the metal particles. To probe for the individual role of the support in coking, Table 2 Hydrogen chemisorption capacity and catalytic activity Catalyst

State

Pt dispersion (%)

IRate o f C O 2 convn. ( l 0 -S. mol/g, s)

Pt/'y-A1203

fresh used 2 regenerated 3

35 8 35

3.4 0.3 3.2

33 3.2 9 3.1 regenerated 3 32 3.2 1at 875 K; 2after testing at 875K for 25 hrs; 3coke burned offin air at 675K

Pt/ZrO 2

fresh

used 2

the weight of coke deposited after 25 hours on stream was compared for the supports and catalysts (see Table 3). It can be seen that the level of coking on Pt/y-Al203 does not exceed that of the support alone. On the other two catalysts coke is reduced by the presence of Pt. In general, the extent ofcoking decreased in the order A1203>>TiO:>ZrO2 irrespective of the presence of Pt. When the amount of coke formed as a function of time on stream is compared to the decrease in catalytic activity (see Fig. 3), two regimes of deactivation can be noticed for the strongly deactivating catalysts, i.e., a slow initial deactivation which is followed by a rapid loss of activity. This first phase is characteristic of a slow transformation of the reactive carbon into less reactive coke. The second phase is attributed to carbon formed on the support which accumulates there and rapidly covers the Pt particles when its amount reaches a critical value causing the sudden decay of catalytic activity. This dearly indicates that the support is able to convert a fraction of methane but contributes primarily to coking. To investigate the role of acid sites in the conversion of methane to coke and hydrogen, the acid sites of the catalysts were characterized by sorption and temperature programmed desorption (t.p.d.) of pyridine. T.p.d. of pyridine (see Fig. 4) suggest a higher

467 concentration of stronger acid sites on Pt/y--ml203 (presence of a second peak around 650 K) than on Pt/ZrO2. I.r. spectra of adsorbed pyridine evidenced only Lewis acid sites on both catalysts. 3.2 Individual reaction steps on the Pt/ZrO2 The question remaining now to be addressed is the role of reactants and catalyst phases in the network of reactions leading to CO, H 2 and carbon. For that purpose pulse reaction/titration experiments using the individual reactants mostly in sequence were applied. Pulsing of methane at 875K over prereduced Pt/ZrO2 yielded twice the number of moles of hydrogen in the gas phase indicating quantitative dissociation of methane. Fig. 5 shows the evolution of hydrogen as detected by mass spectrometry during pulsing of methane. As ZrO2 was virtually inactive for methane dissociation, we conclude that Pt is the main catalytically active component. In contrast to ZrO2, y-AI203 was rather active in methane decomposition. This higher activity of y-Al203 as compared to ZrO2 parallels the higher concentration of Lewis acid sites on Pt/y-Al203 and indicates that Lewis acid sites are affiliated with the catalytic activity of the Table 3 support to dissociate methane. Coke deposition at 875K, after 25 hrs Subsequent admission of oxygen in pulses time on stream indicated that carbon deposited by methane decomposition could be removed quantitatively Catalyst Amount of coke by oxidation. The carbon remaining on the (10 .6 . mol/[) catalysts could also be quantitatively removed in y-Al203 56 the presence of Pt by CO2. CO was the only Pt/y-A1203 59 reaction product. TiO2 14 In order to investigate whether CO2 reacts in Pt/TiO2 6 a concerted way with surface carbon or whether it dissociates first to CO and adsorbed oxygen ZrO 2 9 and the adsorbed oxygen reacts, infrared Pt/ZrO2 1 spectroscopy, pulse reactor studies and XANES measurements were used. The i.r. spectrum of a prereduced (lhour at 675K in 5%H2/N2) Pt/ZrO2 catalyst in contact with CO2 at 775K is shown in Fig. 6. The spectrum shows the presence of linearly bound CO on Pt at 2053 cm-~ [ 15]. Additionally, bands of carbonate type species appeared in the region between 1375 and 1540 cm ~. Over pure supports (in the absence of Pt) the CO band was not seen, but peaks in the carbonate region were observed. The presence of CO on Pt (seen by the i.r. bands at 2053 cm~ ) indicates the dissociation of CO2. This is further confirmed by the evolution of CO (detected by mass spectroscopy) during pulsing CO 2 over Pt/ZrO 2 and in situ XANES measurements. The XANES of the Pt L-rn edge in the presence of reactant gases at 775K is shown in Fig. 7. The comparison of the Pt white lines in the presence of CO2 and O2 indicates that Pt is (partially) oxidized in the presence of CO2 in all catalysts. This indicates rapid dissociation of CO2 into CO and atomic oxygen as outlined above. The extent of the change of the Pt white line led us to conclude that not only the top most layer of the particles is involved, but also a considerable fraction of the bulk atoms takes part in the reaction. This is even more surprising as one notes how rapidly the white line changes between the oxidized state and a state characteristic of fully reduced Pt when one switches between CO2 and oxygen containing atmosphere on the one hand and hydrogen and/or methane containing

468

=i

+!

== =~

oc

s ~ .05

200

.==

j 4

I

5

6

7

8

<

9

5

0 21300

o L~

1800

1600

1400

o

10 15 Time [mini

Wavenumber [cm-1]

0

Pulse number

Figure 5 Hydrogen evolution during methane pulsing over Pt/ZrO:; 875K

3~

Figure 6 Lr. spectra of Pt/ZrO: during contact with CO, at 775K

T

I t

CO~

o~

~o4o(ll

r 75

11550

+

115~

11570

11580

11590

Energy (eV)

Figure 7 XANES spectra of Pt/ZrO: in different gas atmospheres at 775K

11600

100

hi6 --

0

+

~

~

+

10

20

30

40

Time (hours)

Figure 8 Activity of lwt% Ni/ZrO: (Nil) and

6wt% Ni/Zr02 (Ni6)for CO:/CH4 reforming

469 atmosphere on the other hand.

3.3 Ni/ZrO2 catalysts Characteristic physicx>-chemic~ properties of two typical catalysts are compiled in Table 4 [ 16]. Fig. 8 shows the activity of two Nrt/ZrO, catalysts as a function of time. It can be seen that the catalyst containing the higher loading of Ni deactivated most rapidly, while the one with the low Ni loading was remarkably stable. The nature of coking in the case of Ni based catalysts is different firomthat of the noble metal based catalysts. Transmission electron microscopy (TEM) of the 6Ni/ZrOz sample after one hour on stream (Fig. 9) shows the presence of filamentous carbon. The carbon filaments grew at the interface between Ni and ZrOz. The presence of the Ni particle at the tip of the filament was checked by EDAX during the TEM measurement. However, filamentous carbon was not observed to be formed for 1Ni/ZrO2 after 50 hours time on stream. Table 4 Physicochemical characteristics Ni/ZrO2 catalysts Catalyst

BET area (m2/g)

Ni content (wt%)

Average Ni particle size (nm)

1Ni/ZrO z

26

1.1

2

6Ni/ZrO 2

19

6.2

9

4. DISCUSSIONS Steady state and non steady state kinetic measurements suggest that methane carbon dioxide reforming p r o s in sequential steps combining dissociation and surface reaction of methane and CO2. During admission of pulses of methane on the supported Pt catalysts and on the oxide supports, methane decomposes into hydrogen and surface carbon. The amount of CH4 converted per pulse decreases drastically after the third pulse (this corresponds to about 2-3 molecules of CH4 converted per Pt atom) indicating that the reaction stops when Pt is covered with (reactive) carbon. CO2 is also concluded to dissociate under reaction conditions generating CO and adsorbed

470 oxygen atoms. As nearly all of the carbon on Pt formed via methane decomposition can be oxidized in a concerted way by CO2 or by the adsorbed oxygen produced from CO2 (see Table 4) we conclude that the carbon on the surface is a reactive intermediate in the reforming reaction to yield CO. The overall reaction sequence can be summarized in the following elementary steps. CH4

+

-~

~

C-~

+ 2H 2

CO 2

+

2-~

~

CO-~

+ O-~

C~

+ O-~

~

CO~

+-~

-~

2CO

+ 2-~

~

2CO

+ 2H 2

2CO-~ CH4

+

CO 2

Infrared spectroscopy and XAS measurements support this model of the reaction sequence. I.r. spectroscopy shows that CO is formed and adsorbed on Pt upon contacting the catalyst with CO2 at 875K. In situ XAS measurements show that a large fraction of Pt (and not only the surface atoms) are oxidized by CO2 at 775K indicating the presence of atomic oxygen on/in the Pt particles. Subsequent admission of methane completely reduced Pt in such a catalyst. Oxidation and reduction of Pt are rapid. It is important to emphasize that XAS measurements demonstrate that these processes are similar for all catalysts investigated indicating that the support does not induce a particular surface chemistry on Pt. This is in accordance with the similar intrinsic activity of Pt sites (TOF, in Fig. 1) of all catalysts. The carbonate species seen on the supports irrespective of the presence of Pt do not seem to play a key role in the catalytic cycle. Let us now use the sequence of elementary steps to explain the activity loss for some of the catalysts. The combination of hydrogen chemisorption and catalytic measurements indicate that blocking of Pt by coke rather than sintering causes the severe deactivation observed in the case of Pt/y-AI203. The loss in hydrogen chemisorption capacity of the catalysts after use (Table 2) is attributed mainly to carbon formed by methane decomposition on Pt and impeding further access. Since this coke on Pt is a reactive intermediate, Pt/ZrO2 continues to maintain its stable activity with time on stream. Coke formation on these catalysts occurs mainly via methane decomposition. Deactivation as a function of coke content (see Fig. 3 for Pt/y-AI203) seems to involve two processes, i.e., a slow initial one caused by coke formed from methane on Pt that is non reactive towards CO2 (see Table 3). In parallel, carbon also accumulates on the support and given the ratio between the support surface and metal surface area at a certain level begins to physically block Pt deactivating the catalyst rapidly. The coke deposited on the support very close to the Pt- support interface could be playing an important role in this process. Using the t.p.d, and i.r. spectra of adsorbed pyridine we conclude that high concentrations of Lewis acid sites (e.g., present on y-AI203) are responsible for high rates of coking. Such coke on the support is difficult to be removed by CO2 as pure supports showed no catalytic activity. Therefore, coke formed on the support causes deactivation when it covers Pt. This corresponds to the rapid decay of activity seen, for example, for Pt/y-A1203 (Fig.3). The small continuous loss in activity (for all catalysts) may be due to part of the coke formed from methane on Pt undergoing aging which makes it non reactive. Thus, we conclude that it is the uniquely low concentration of weak Lewis acid sites on the ZrO2 support that seems to be indispensable for minimizing coking in Pt/ZrO2. Addition of Pt could be reducing the number of Lewis sites on ZrO2

471 further and explains why P t / Z r O 2 fOrmS even less coke than ZrO2 (Table 3). In contrast to the Pt catalysts discussed above, Ni based catalysts (i.e., also when supported on ZrO2) usually form coke at such a rapid rate that most fixed bed reactors are completely blocked after a few minutes time on stream (see Fig. 8) [ 16]. The coke formed with the Ni catalysts is filamentous. The Ni particle remaining at the tip of the filament hardly deactivates as the coke formed on its stuface seems to be transported through the metal particle into the carbon fibre, but the drastic increase in volume causes reactor plugging and prevents use of the still active catalyst (see Fig. 8). The TEM photographs indicate that the carbon filaments have similar diameters to those of the Ni particles. As the metal particle size decreases the filament diameter should also decrease. It has been shown that the surface energy of thinner filaments is larger and hence the filaments are less stable (11, 17-18). Also the proportion of the Ni(111 ) planes, which readily cause carbon formation, is lower in smaller Ni particles (19). Therefore, even though the reasons are diverse, in practice the carbon filament formation ceases with catalysts containing smaller Ni particles. Consequently, well dispersed Ni catalysts prepared by deposition precipitation of Ni (average metal particle size below 2-3 nm) were stable for 50 hours on stream and exhibited no filamentous coke [ 16].

5. CONCLUSIONS Methane reforming with carbon dioxide proceeds in a complex sequence of reaction steps involving the dissociative adsorption/reaction of methane and CO2 at metal sites. Hydrogen is generated during methane dissociation. In the second set of reactions CO2 dissociates into CO and adsorbed oxygen. The reaction between the surface bound carbon (from methane dissociation) and the adsorbed oxygen (from CO2 dissociation ) yields carbon monoxide. A stable catalyst can only be achieved if the two sets of reactions are balanced. Two ways to design such a stable catalysts are described. The support and the active metal play an important role in ensuring catalyst stability. For a noble metal such as Pt, which itself has a low tendency to form coke [ 1], the minimization of the concentration of acid sites on the support is most critical. ZrO2 is unique in that respect as it combines strong anchoring of the Pt metal particle (important for high temperature operation) and very low concentrations of (Lewis) acid sites (important for minimizing carbon formation on the support). Excellent results were obtained with a 0.Swt%Pt/ZrO2 catalyst which proved to be well suited for upscaling [20]. For non-noble metal catalysts, such as oxide supported Ni, also the metal function has to be adapted to minimize coke formation. This results from the fact that the rate of methane decomposition over Ni catalysts is enhanced relative to CO2 dissociation leading to a significantly higher concentration of carbon at the Ni particle. This triggers transport of the carbon through the Ni particle and formation of (hollow) carbon filaments carrying the metal particle at the top. It was found that reducing the size of the Ni particles slows down the overall rate of formation of the carbon filament. Thus, preparing a Ni/ZrO 2 catalyst with a particle size below 2nm prevented the generation of filamentous coke and allowed stable operation for prolonged periods.

Acknowledgments We gratefully acknowledge the support for this work from the JOULE II Programme (Energy from fossil fuels: Hydrocarbons, contract no. JOU2-CT92-0073) and Human Capital and Mobility Programme (contract no. ERB4001GT941163) of the European Union. The authors are indebted to M. Englisch and A. Jentys for valuable discussions on XANES of Pt/ZrO:.

472 6. REFERENCES

1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20..

A.T. Ashcrott, A_K. Cheetham, M.L.H. Green and P.D.F. Vernon, Nature, 225 (1991) 352. J.T. Richardson and S.A. Paripatyadar, Appl. Catal., 61 (1990) 293. K. Seshan and J.A. Lercher, in J. Paul and C. Pradier, (Eds.) "Carbon dioxide: Environmental issues", The Royal Soc. Chem., Cambridge, 1994, p 16. A. B hattacharya and V.W. Chang, Int. Conf. on Catalyst Deactivation, Oostende, Oct., 1994, Stud. Surf. Sei Catal. 88 (1994) 207. G. Kurz and S. Teuner, Erdol. Kohle, 43(5) (1990) 171. P.F. van den Oosterkamp, Q. Chen, J.A.S. Overwater, J.R.H. Ross and A.N.J. van Keulen, Meeting of"Large Chemical Plants", Antwerp, Belgium, Oct., 1995. Gas Process Handbook '92, Hydocarbon Processing, 90, (1992). I.M. Bodrov, L.O. Apel'baum, Kinet. Katal. 8 (1967) 379. J.R. Rostrup Nielsen, J. Catal., 144, (1993) 38. N.R.Udengaard, J.H. Bak Hansen, D.C. Hanson, J.A. Stal, Oil Gas J., 90 (1992) 62. J.R. Rostrup-Nielsen, J.Catal., 27 (1972) 343. R,T.K. Baker, M.A. Barber, P.S. Harris, F.S. Feates and R.J. Waite, J.Catal., 26 (1972) 51. K. Seshan, H.W. ten Barge, W. Hally, A.N.J. van Keulen, J.R.H. Ross, Stud. Surf. Sci. Catal. 81 (1994)285. J.H. Bitter, W. Hally, K. Seshan, J.G. van Ommen and J.A. Lercher, Catal. Today, accepted for publication (1995). L.H. Little, Infrared Spectra of Adsorbed Species,, Academic Press, New York, 1966, p 54. W. Hally, H.J. Bitter, K. Seshan, J.R.H. Ross, and J.A. Lercher, Int. Conf. on Catalyst Deactivation, Oostende, Oct., 1994, Stud. Surf. Sci. Catal. 88 (1994) 167. I. Alstrup, J. Catal., 109 (1988) 241. P.K. de Boks, A_J.H.M. Kock, E. BoeUeerd, W. Klop, J.W. Geus, J. Catal., 96 (1985) 454. C.H. Bartholomew, Cata. Rev. Sc. & Engg., 24, 67 (1982) K. Seshan, P.D.L. Mercera, E. Xue, J.R.H. Ross, German Patent, P43 13 673.7 (1993), US Patent Appl. PCT/DE94/00513 dated (1994)

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

473

C o m p a r i s o n o f perovskite and hexaaluminate-type catalysts for C O / H 2 - f u e l e d gas turbine combustors. C. Cristiani', G. Groppi', P. Forzatti', E. Tronconi", G. Busca b and M. Daturi b "Dipartimento di Chimica Industriale e Ing. Chimica, Politecnico di Milano, Piazza Leonardo da Vinci, 20133 Milano- Italy Istituto di Chimica, Facolta di Ingegneria, Universitb. di Genova Fiera del Mare, Padiglione D, P.le Kennedy- 10136 Genova- Italy b

In this work the results of catalytic activity tests in CI-Ia, CO and H2 combustion over perovskite (LaCoOs, LaM/IO3 and LaFeOs) and hexaaluminate-type (BaMn~llOlg, Sro.sLao.2MnA/~Olg, and BaFeAl~O19) systems are compared in order to investigate the potential of such materials as catalysts for syngas fueled combustors for gas turbines. Perovskites-type catalysts are shown to be the most active systems in the combustion of all the investigated fuels but to suffer from thermal stability problems that constrain their use in high temperature applications. Mn-substituted hexaaluminates have been shown to be more active by orders of magnitude in CO-H2 combustion than in CI-h combustion. Scale up of the activity data by mathematical modelling has demonstrated the potential of such catalysts in meeting the operating requirements of syngas fueled catalytic combustors. 1. INTRODUCTION Catalytic combustion represents a promising method for the effective combustion of lean fuel-air mixtures in gas turbines (GT) with minimum formation of NOx, CO and unburnt hydrocarbons (UHC). Different configurations of the catalytic combustor have been proposed. In the Catalytically Stabilized Thermal (CST) combustor a preheated and premixed fuel-air stream passes through a catalyst section where combustion proceeds till complete fuel consumption [ 1]. In this configuration the catalyst reaches temperatures corresponding to the full load firing temperature of the turbine, that is presently around 1533 K and that is expected to grow further in the near future [2]. The Hybrid Combustor aims at minimizing the problems associated with the thermal stability of the catalyst and with its resistance to thermal shocks. In this configuration the fuel feed is split into two parts: one is fed to the catalyst after being premixed with preheated air, the other is fed to a downstream homogeneous section [3]. The fuel/air ratio of the catalyst feed is constrained to prevent catalyst temperatures exceeding 1273 K. The hot gas leaving the catalyst improves the stability of the premixed combustion in the homogeneous section. Catalytica Inc. has recently developed a different configuration of a hybrid combustor in which all the fuel is fed to the catalyst temperature whose temperature is constrained by means of a proprietary catalyst design [4]. Downstream the catalyst the combustion is completed in a homogeneous section. Among the catalysts under investigation, noble metal based systems are the most active in combustion reactions, and specifically PdO based catalysts are best suited for the combustion

474 of methane, namely the typical fuel for gas turbines [5]. However operation at high temperature may cause deactivation due to sintering of the support and of the noble metal crystals and to the thermal reduction of active PdO into inactive Pd ~ [6]. Different possible solutions to the catalyst thermal stability problems are currently attempted [6]. Mixed oxide catalysts represent a valuable alternative to noble metals catalysts in view of their thermal stability. However a major drawback for the practical use of these materials in catalytic combustors is their limited activity, particularly in the case of GT fueled with natural gas. The limitations on catalytic activity can be overcome if more reactive fuels than methane are considered. CO/H2 mixtures from carbon and biomass gasification represent a perspective alternative to methane as gas turbine fuels in the near future in view of the increasing requirements for the utilization of different energy sources. Moreover the use of biofuel instead of fossil fuels may allow for the reduction of CO2 emissions to the atmosphere. Indeed CO/HE mixtures are more reactive than methane over most combustion catalysts. Accordingly, in this case, mixed oxide catalysts may provide a viable alternative to noble metal catalysts. Among mixed oxide catalysts both perovskite and hexaaluminate type systems seem most promising for this specific application. Lanthanum metalates exhibiting a distorted perovskite-like structure are reported as very active combustion catalysts [7]. Their activity is enhanced if part of the La ions in the structure are replaced with large divalent cations of alkaline-earth [7]. On the other hand M-substituted hexaaluminates (M--transition metal ion) have shown excellent thermal stability properties [8] which are strictly related to their peculiar layered structure resulting from the stacking of spinel-type blocks separated by planes containing large divalent (Ba, Sr) or trivalent (La) cations. The incorporation of transition metal ions into the structure by substitution of part of the Al ions is responsible for the catalytic activity in CH4 combustion while preserving the thermal stability of the material. Among transition metal ions (Mn, Fe, Co, Ni, Cr) Mn and Fe have been found to provide the highest combustion activity [8,9]. In this paper we attempt a preliminary investigation on the feasibility of catalytic combustion of CO/H2 mixtures over mixed oxide catalysts and a comparison in this respect of perovskite and hexaaluminate type catalysts. The catalysts have been characterized and tested in the combustion of CO, H2 and CH4 (as reference fuel). The catalytic tests have been carried out on powder materials and the results have been scaled up by means of a mathematical model of the catalyst section of the Hybrid Combustor.

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

2.1. Preparation of the catalysts. The hexaaluminate type catalysts BaMIL~llOI9, Sr0.sLao.2]LvinAlllOl9, and BaFeAlllOl9 have been prepared according to a coprecipitation route described elsewhere [10]. Briefly nitrates of the constituents (Aldrich A.C.S. reagents) except for AI were solved in hot water (T = 333 K) under vigorous stirring. After acidification with HNO3 (pH=l) Al('NO3)3"9H20 has been added and the resulting solution has been poured, under vigorous stirring at constant temperature (T=333 K), into another one containing (NH4)2CO3 in excess. The resulting slurry has been aged for about 3 hs at 333 K with pH values ranging from 7.5 to 8.0. Then it has been filtered and washed to remove nitrates and (NH4)2CO3 in excess. The precursors have been dried at 383 K overnight and finally calcined at different temperatures up to 1573 K for 10 hs. Analyses by atomic absorption of mother liquors and washing waters confirmed that the precipitation occurred in a quantitative way, while chemical analyses on the calcined samples showed that the actual ratios of the constituents corresponds to the nominal ones.

475 The perovskite-type catalysts LaCoO3 and LaMnOa, were prepared via a conventional coprecipitation method, starting from nitrates or acetates of the constituents. Stoichiometric amounts of La(NOa)3"6H20 and Co(NOa)E'6H20 or Mn(CH3COO)2"4H20 (Carlo Erba reagents) were solved in water under stirring. The precipitation of the catalysts precursors was obtained by addition of excess (NI-h)2CO3. The resulting slurry was aged overnight; then water was removed by evaporation and the resulting solid was dried at 383 K. Another perovskite catalyst LaFeO3 was prepared as aerogel through the supercritical drying technique using La(CHsCOO)3"6H20 and [CH3COCH=C(O-)CH3]3Fe (Aldrich) as starting reagents and methanol as solvent. The solution of the salts was hydrolyzed by addition of a stoichiometric amount of water and then dried at 523 K and 11,721 MPa, under supercritical conditions. The perovskite precursors were calcined at different temperatures up to 973 K. In the case of LaMnO3 a calcination at 13 73 K for 10 hs was also performed. In the following the samples are identified by a notation indicating atomic ratios and the calcination temperature: e.g. LalMnl-973 identifies perovskite LaMnO3 calcined at 973 K.

2.2. Characterization of the catalysts. XRD analyses have been performed with a Philips PW 1050-70 vertical goniometer using a Ni-filtered CuK~ radiation. Cell parameters have been calculated using a least-squares fitting routine. XRD measurements of synchrotron radiation have been collected at SERC laboratories at Daresbury. Surface area measurements have been obtained by nitrogen adsorption at 77K with a BET Fison 1900-Series apparatus. The skeletal FTIR/FTFIR spectra and the FTIR spectra of adsorbed molecules have been recorded by a Nicolet Magna 750 Fourier transform instrument. The adsorption experiments were performed on pressed disks of pure powders, after outgassing at 300-1070 K using piridine and CO2 as probe molecules. UVVis-Diffuse Reflectance spectra have been recorded with a Jasco-619 double beam spectrophotometer equipped with an integrating sphere and using BaSO4 as reference. SEM micrographs have been collected by a Jeol JSM 35C. EPR spectra have been collected by a Varian E-109 spectrometer. Zero point charge has been determined by mass titration. 2.3 Measurements of catalytic activity. Catalytic activity tests have been performed in a quartz microreactor (I.D.=0.8 cm) filled with 0.45 g of fine catalyst powders (dp=0.1 micron). The reactor has been fed with lean fuel/air mixtures (1.3% of CO, 1.3% of H2 and 1% of CFL in air respiectively) and has been operated at atmospheric pressure and with GHSV = 54000 Ncc/gcath- . The inlet and outlet gas compositions were determined by on-line Gas Chromatography. A 4 m column (I.D. =5mm) filled with Porapak QS was used to separate CI-I4, CO2 and H20 with He as carrier gas. Two molecular sieves (5 A) columns (I.D.=5 ram) 3m length, with He and Ar as carrier gases, were used for the separation and analysis of CO, N2, 02, CFL, and H2, N2, 02 respectively

3. RESULTS AND DISCUSSION

3.1. Physico-chemical characteristics of the catalysts. Extensive data on the characterization and the thermal evolution of the different catalysts have been reported elsewhere [9-14]. Phase composition, cell parameters and surface area of the final materials are summarized in Table 1. The XRD data indicate that for all the hexaaluminate-type samples the formation of the final phase begins at 1273 -1373 K and requires calcination temperatures of 1473-1573 K to be completed.

476 Table 1 Morpho!o~;ical and structural properties of the catal~,sts. surf. area Crystal Sample m2/g structure* I Illll

I

I

II

II

IIII III

calculated cell parameters* ao(A) bo(A) Co(A) Illllllllllll

Ill

Ill I Illl II

Ill

Ba1Mn1A111-1573

17

Ba-13-A1203 5.6401(3)

5.6401(3)

22.737(1)

BalFelAI11-1573

7

Ba-13-AI203 5.628(5)

5.628(5)

22.89(2)

Sr0.8La0.2Mnl All 1-1573

10

5.594(5)

22.09(2)

LalCol-973

NiP

5.594(5)

8

LalFel-973 LalMnl-973 I

II II

* For the dominant phase, ~ ,

hexagonal 5.454(2) 5.454(2) 13.135(9) perovskite 18 orthorombic 5.693(5) 7.857(4) 5.561(7) perovskite 7 hex. +orthor. 5.534(2)** 5.534(2)** 13.363(8)** perovskites ,,,,,,, magnetoplumlme, ** assuming a rombohedral structure III

_..

IIIIIIIIIIIIII

III

I

Upon calcination at 1573 K all the samples are monophasic and consist of a Ba~12019-type phase [JCPDS 26-135] with a Ba-13-Al203-1ike structure [12], in the case ofBaMnlAll 1 and BalFelAII 1, and of a SrAl~:O19-type phase with a magnetoplumbite-like structure [JCPDS 26-976], in the case of Sr0.SLa0.2MnlAll 1. The formation of a solid solution is confirmed by the expansion of the unit cell of BaAl12019 (a0=b0=5.607 ~ e0=22.900 /x.) and SrA112O19 (a0=b0=5.585 ~ c0=22.07/x) due to partial replacement of Al3+ ions by Fe*+.and Mn n§ With respect to the oxidation state of the transition metal ions in the structure a detailed characterization study by means of Rietveld structure refinement of the XRD measurement with synchrotron radiation and of EPR measurements provided evidence for the presence of both Mn 2§ and Mn 3§ ions in the structure of BalMnlAll 1-1573. Likewise the presence ofFe 2+ and Fe 3~ ions in BalFelAll 1-1573 is supported by UV-Vis-DR spectra. The peculiar layered-nature of Ba-13-AI203 and magnetoplumbite structures, generated by the repetition of layers of Al-O-containing spinel-type blocks intercalated by planes containing the largest cations, results in an excellent resistance to thermal sintering, so that surface areas of 17, 8 and 7 m2/g have been measured for BaMnlAI11, Sr0.SLa0.2MnlAI11 and BalFelAI11 calcined at 1573 K in line with previous literature results [8]. FTIR spectra of adsorbed probe molecules and zero point charge measurements have also shown a predominant basic character of the catalysts due to preferential exposure of large earth-alkali and rare-earth metal ions at the surface. In perovskite-type catalysts the formation of the final phase is completed already at 973 K. XRD and skeletal FTIR/FTFIR data for LalCol, LalMnl and LalFel calcined at 973 K evidence that only LalFel-973 is actually monophasic and consists of a perovskite-type phase with orthorombic structure. A perovskite type phase with hexagonal-rombohedral structure represents the main phase of Lal Co 1-973, but traces of Co304 and La2COs are also present. In the case of LalMnl-973 two phases have been detected both with perovskite-type structure, one orthorornbic and the other rombohedral. The calculated cell parameters of the dominant perovskite-type phase are reported in Table 1 for the three samples. The results compare well with those reported in the literature [JCPDS 37-1493, 32-484, 25-1060] which refer to similar samples prepared via solid state reaction. All the perovskite-type samples are markedly sintered

477 already at 973 K, as indicated by the relatively low BET surface areas reported in Table 1. However it appears that the preparation via the aerogel technique results in a significantly higher surface area (18 m2/g) as compared to the coprecipitation method (7-8 m2/g), in line with SEM photographs showing different crystal size dimensions for the different samples. FTIR measurements using pyridine and CO2 as probe molecules indicate that the perovskite systems possess a predominantly basic character.

3.2. Catalytic activity measurements The combustion reactivity of CH4 over the different catalysts has been preliminary investigated as reference activity tests. In Figure l a CI-I4 conversions are plotted versus reaction temperature for all the investigated systems except for Fe-containing ones (La1Fe1973 and BalFelAI11-1573). These latter however have exhibited the lowest activities in the oxidation of all the investigated fuels compared with the catalysts of the corresponding class. Moreover Fe-substituted hexaaluminates have been previously reported to deactivate under reaction conditions at high temperature (1173 K) [14]. Accordingly Fe-containing systems appear of minor interest within the scope of this work. For all the investigated systems the experimental curves of CH4 conversion compare well with literature data [8,15,16], in line with the evidence on phase composition and morphological properties described above. LalCol-973 and La1Mn1-973 exhibit similar activities, slightly higher than Sr0.8La0.2Mn1AI11-1573 that in turn is the most active Mnsubstituted hexaaluminate. Figure l a also shows that a marked deactivation has occurred in LalMnl catalyst upon calcination at 1373 K for 10 hs. La1Mn1-1373 exhibits lower activity than Mn-substituted hexaaluminates. Besides stability problems are evident for this system for reaction temperature above 973 K: no conversion enhancement has been obtained on increasing reaction temperature over this value. Similar apparent activation energies ranging from 21 to 24 kcal/mole have been derived from the experimental curves for all the investigated systems but La 1Mn 1-1373. Figures 1b and lc report the conversions of CO and H2 as functions of reaction temperature for the different catalysts. In line with general literature indications over all the investigated systems CO and H 2 oxidize at markedly lower temperatures than CI-h; T~0%ranging from 393 to 523 K and from 393 to 573 K have been observed for CO and H2 respectively, to be compared with 773-823 K required by CH4. With respect to CO oxidation an activity order similar to that described above for CI-I4 combustion has been obtained. A specific activity enhancement is observed for LalCol-973 that has provided a 10% conversion of CO already at 393 K, 60 K below the temperature required by La1Mnl-973. This behavior is in line with literature reports on CO oxidation over lanthanum metallates with perovskite structures [17] indicating LaCoO3 as the most active system. As in the case of CI-h combustion, calcination at 1373 K of LalMnl has resulted in a significant decrease of the catalytic activity. Indeed the activity of La1Mn1-1373 is similar to those ofMn-substituted hexaaluminates calcined at 1573 K. Differently from the results of CH4 combustion tests no stability problems have been evidenced under reaction conditions for La 1Mn 1-1373 possibly due to the low temperature range of CO oxidation experiments. Similar apparent activation energies have been calculated for all the investigated systems, ranging from 13 to 15 Kcal/mole, i.e. almost 10 Kcal/mole lower than those calculated for CH4 oxidation. In the case of H2 oxidation the two investigated classes of catalysts show different behaviors. Again perovskite type catalysts calcined at 973 K show higher combustion activity than hexaaluminates calcined at 1573 K, but characteristic values of apparent activation energy (5-7 Kcal/mole) have been calculated for perovskite catalysts that are markedly lower than

478 those observed for Sr0.SLa0.2MnIAIll-1573 and BalMnlAIll-1573 (15 Kcal/mole) and result in the smooth increment of H2 conversion with temperature shown in Figure 1c. JO0 90

a)

80 70 60

-'---4

50 40 30 20 ~0

~23 ~00 90

773

I

I,

i

823

873

473

573 TEMPERATURE (K)

1

/

923 973 ~023 ~073

TEMPERATURE (K)

~23

b)

80 70 60 50 40 30 20 ~0 73

1

6~3

773

IOO9o;- ~) 80 ~-

/~

7O

60

50

4O 3O

2O

I0 73

473

573 673 TEMPERATURE (K)

I

773

873 8"

Figure 1. Fuel conversion vs reaction temperature a) CI-h conv.; b) CO conv.; c) H2 conv. A LalMnl-973; T LalCol-973; I LalMnl-1373 I Sr0.8La0.2MnlAI11-1573; r-I BalMnlAI11-1573

For Mn-substituted hexaaluminates the 1-12 oxidation activity is markedly lower than that in CO oxidation: 10% conversion of H2 over these systems requires temperature almost 100 K higher than CO. Finally it is apparent in Figure 1c that calcination at 1373 K of LalMnl has c a u s ~ a marked deactivation, similar to that described for CO oxidation. In conclusion the catalytic activity data herein reported show that: i) CO and H2 exhibit catalytic combustion reactivities that are orders of magnitude higher than that of CI-L over both perovskites and hexaaluminate type systems; ii) perovskite-type catalysts calcined at 973 K are more active than hexaaluminates calcined at 1573 K in the combustion of all the investigated fuels, but the activity of perovskites markedly decreases upon calcination at 1373 K up to value quite similar to that of Sr0.8La0.2MnlAI11-1573 and BalMnl AI 11-1573. Besides stability problems under reaction atmosphere at temperatures higher than 973 K have been evidenced for LalMnl-1373 in CI-L combustion tests. Eventually such thermal stability problems restrain the use of perovskite as catalysts for high temperature applications. On the other hand the high reactivity of CO and H2 could afford the use of less active but more stable Mnsubstituted hexaaluminates. To better address this point a mathematical model analysis of the catalytic activity data has been performed. Procedure and results will be described in the next section.

479 3.3. Mathematical model analysis The experimental data described above indicate Mn-substituted hexaaluminates as the most promising materials among the investigated systems for high temperature catalytic combustion of syngas mixtures. However such data have been collected under a pure chemical regime, whereas the combustor performances are determined by both physical and chemical phenomena occurring in the monolith catalyst that are actually used to achieve GT constraints on pressure drops. Accordingly in order to make a preliminary assessment of the potential of such material in meeting the requirements of a Hybrid Combustor for syngas fueled gas turbines a scale up of the experimental data has been performed by means of mathematical model analysis, choosing Bal Mn 1AI 11-1573 as the reference catalyst. A single channel model accounting for variable distribution in the channels (assumed to be circular) on both axial and radial direction has been used for this purpose. The model was derived from a previous one, described in details in reference [18], simulating the catalyst section of a natural gas fueled Hybrid Combustor. It consists of enthalpy and mass balance equations for the gas phase derived under the assumptions of steady state conditions and laminar flow. It accounts for homogeneous and heterogeneous oxidation reactions, diffusion and convection of heat and mass in the gas phase, diffusion in the catalyst pores (through calculation of the effectiveness factor by simple analytical expressions) and gas property variation with temperature according to power law dependencies. The effects of solid conduction, radiation and hydrodynamical development have been neglected according to the results of a previous work in which the adequacy of multidimensional models of the catalyst section of a Hybrid Combustor fueled by CI-L was investigated [18]. Kinetic expressions for the heterogeneous reactions have been derived from fuel conversion data described above, assuming isothermal and plug flow behavior of the laboratory microreactor as well as a first order dependence on CO and H2 concentrations. No mixture effects have been included in such expressions according to the results of catalytic activity tests performed with a 1/1 CO-H2 fuel mixture at operating conditions identical to those described for experiments with pure fuels. Using such a mixture CO and H2 conversion curves obtained over BalMnlAI11-1573 have shown only minor differences from those obtained with pure CO and H2 reported in Figures l b and l c respectively. To account for the homogeneous reaction simple molecular kinetic expressions have been taken from the literature [ 19,20]. The catalyst section in Hybrid Combustors typically consists of 3-4 short monolith segments with small channel diameter. Such a configuration provides low pressure drops, improved thermal shock resistance and high gas-solid mass and heat transfer rates, thanks to the high specific geometric surface area and to the enhancement of transfer coefficients associated with multiple inlet effects. Accordingly a reactor configuration consisting of four ceramic monolith segments of 2.5 cm length each with circular channels of 1.2 mm diameter and an open frontal area of 0.68 has been assumed. Operating parameters for the simulations are reported in the figure captions. They have been chosen as representative of real conditions in gas turbine combustors except for gas velocities that, due to the laminar flow assumption, correspond to the lowest values of the operating range. Figure 2a shows the calculated axial profiles of catalyst temperature corresponding to different fuel compositions: pure CO, pure H2 and 1/1 CO-H2 mixture. Typical light-off profiles have been obtained with steep T-gradients confined to a narrow zone in which transition from the chemical to the diffusion controlled regime occurs. As a consequence of the higher oxidation reactivity of CO than H2 over the investigated catalyst, the fight-off is delayed on increasing the H2 content of the fuel. However while ignition of pure H2 is almost completely blown out and only occurs at the end of the last catalyst segment, the light-off of the CO/H2 mixture is only slightly delayed with respect to that of pure CO occurring at the

480 beginning of the first catalyst segment. Indeed in the adiabatic conditions of real combustor operation the catalyst heating due to the ignition of the CO fraction of the fuel immediately promotes the light-off of the whole fuel mixture even though mixture effects are not observed under the quasi-isothermal conditions of laboratory activity tests. The slight downstream shift observed with the 1/1 mixture is only due to the lower CO concentration resulting in a delayed ignition. T-profiles in Figure 2a also show that immediately after the ignition of He containing fuels the catalyst temperature can exceed the adiabatic reaction temperature. This phenomenon is typical of diffusion controlled chemical reactions in the presence of fast diffusing reactants such as He [ 1]. It arises from the steady state balance between heat production by reaction, controlled by the extremely fast He diffusion rate, and heat release from solid to gas phase that in turn is determined by the intrinsically slower heat diffusion. Noteworthy in model simulations the fuel inlet concentration has been fixed to achieve an adiabatic reaction temperature of 1273 K in line with literature indications on catalyst sustainable thermal stresses in Hybrid Combustors. Accordingly overtemperature phenomena causing additional thermal stresses may seriously injure the catalyst durability. The axial profiles of calculated average cupmix temperature of the 650 TEMPERATURE (K) gas phase plotted in Figure 2b are strictly related to the wall temperature profiles described just 450 f a ) S ADIABATIC TEMP~.ATURE ....." .... above. Downstream from the catalytic fight-off the gas phase is -'."'" i ."" .....---"-~ f"-"'"initially heated up due to heat 050 . . . . . release by the incandescent catalytic / ." ~/. , / / walls. Then, if the gas temperature / ." ,,.." / / B50 ~ ~--~ ~ / raises up to a sufficient level, i il.o . . i homogeneous combustion ignites in ~. 6so l 1 1 l the boundary of the catalyst wall 0 2 4 6 8 10 and rapidly propagates towards the AXIAL COORDINATE (CM) channel a x i s . This specifically 350 occurs in the case of Hz-containing b) fuel and is the main responsible for 250 the faster heating up observed in I_50 Figure 2b with respect to the case of pure CO. With the simple molecular 050 kinetics herein adopted neglecting 950 any complex kinetic interaction between homogeneous and 850 heterogeneous reactions (only /<~ ~ ' " 750 thermal interactions are considered), homogeneous combustion have 1 . 6500 8 10 been found responsible for about 2 4 6 AXIAL COORDINATE (CM) 50~ of heat production in case of He containing fuels, whereas only a Figure 2. Calculated temperature axial profiles. minor contribution has been a) Catalyst temperature; b) Gas average temperature calculated in case of pure CO. This Ti,=673 K, P=10 atm, v~ m/s, AT,d=-600 K is also due to the water free feed ..... pure CO; ...... CO-H2 mixture; ~ pure H2 that has been assumed in our simulation, indeed water is generally -

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481 recognized to promote strongly homogeneous combustion of CO. Noteworthy under assumptions similar to those reported in this work homogeneous combustion was found to be completely negligible in the catalyst section of a natural gas fueled Hybrid Combustor [ 18]. The effects of CO/H2 feed ratio and of the gas inlet temperature have been extensively investigated in order to check the capability of the catalyst to fast ignite syngas-air mixttwes thus allowing a gas phase heating up to a temperature high enough to effectively stabilize the premixed combustion in the downstream homogeneous section. A design target of 1073 K has been assumed for the gas outlet temperature in line with literature indications for natural gas fueled Hybrid Combustors. However such a temperature may be significantly lowered due to the higher reactivity in homogeneous combustion of syngas mixtures with respect to CI-h. Figure 3 shows the calculated gas outlet temperature as a function of H2/(I-I2+CO) ratio in the feed for different values of the gas inlet temperature. Characteristic curves have been obtained showing a smooth maximum at intermediate fuel compositions and a marked drop corresponding to H2 -rich fuels that become progressively evident on decreasing the gas inlet temperature. Such behavior originates from the following factors: i) the ignition of H2 -rich fuels is delayed by the lower heterogeneous reactivity of H2 but a certain fraction of CO in the fuel, increasing on decreasing the gas inlet temperature, can secure a fast light-off of the whole combustion mixture, ii) with the water free feed assumed in the simulations, gas phase combustion occurs to a significant extent only with substantial H2-contents in the fuel; iii) once light-off has occurred the high diffusion rate of H2 enhances the heterogeneous reaction rate of H2 oxidation in the diffusion controlled regime. Eventually for gas inlet GAS OUTLET TEMPERATURE (K) temperatures up to 673 K the 1300 catalyst is active enough to satisfy the design target on the gas outlet 1200 temperature for wide compositional li00 ranges of the fuel, including typical CO/H2 ratios of fuels from carbon 1000 and/or biomass gasification. On 906 further decreasing the gas inlet temperature the ignition of H2 rich 80C fuel is progressively blown out of 70C the catalyst channels thus preventing the achievement of the 60C O0 0.20 0.40 0.60 0.80 1.00 design target. H2/(CO+H2) FEED RATIO Simulations also provided values Figure 3. Calculated gas outlet temperatures. of the maximum catalyst P=10 arm, v~ m/s, T,a=1273 K temperature: in case of very rich A T~,=723 K, 9 Tm=673 K4 r-1 Tm=643 K, 9 Tt~=623 K H2-content in the fuel, they locally exceed the adiabatic reaction temperature even by two hundreds degrees. However the reliability of such results is still to be established. Indeed model predictions of maximum wall temperature are very sensitive to the kinetic parameters. Besides temperature peaks can be smoothed to a significant extent due to backward heat transmission by wall conduction which was herein neglected. Finally ovenemperature phenomena are related to the H2 diffusion rate, that in turn depends on the flow regime prevailing in the channels. Accordingly the extent of overtemperature would be markedly affected by changing the flow field from laminar to transitional and eventually to fully turbulent regime. However such results definitely make us aware of temperature peaks associated with catalytic combustion of H2 containing fuels.

482 4. CONCLUSIONS Catalytic activities in CO and H2 combustion of perovskite- and hexaaluminate-type systems have been compared to investigate the potential of such mixed oxides as catalysts for syngas fueled combustors for gas turbines. Test results indicated that even though perovskites show the maximum catalytic activity their lack of thermal stability restrains their use in high temperature applications. On the other hand Mn-substituted hexaaluminates have been demonstrated to retain high oxidation activity in CO-H2 combustion upon calcination at 1573 K. Through scale up by mathematical modelling, such activity levels have been shown to meet the operating requirements of Hybrid Combustors fueled by CO-H2 mixtures for gas inlet temperatures higher than 673 K and low reference gas velocities (~8 m/s). Such preliminary results indicate Mn-substituted hexaaluminates as potential materials for syngas fueled gas turbine applications. Simulation results have also evidenced that overtemperature phenomena due to H2-content in the fuel have to be carefully considered.

Acknowledgments This study was supported by MURST (Rome-Italy). The authors thanks prof. R.J. Willey of Northeastern University of Boston for the preparation of the samples via aerogel technique. 5. REFERENCES 1 L.D. Pfefferle and W.C. Pfefferle, Catal. Rev. - Sci. Eng., 29 (1987) 219 2 K.W. Beebe et al., ASME Paper No 95-GT-65, (1995) 3 T.Furuya et al., ASME Paper No 87-GT-99, (1987) 4 R.A. Dalla Betta et al. ASME Paper No 94-GT-260, (1994) 5 M.F.M.Zwinkels et al., Catal. Rev. - Sci. Eng., 35 (1993) 319 6 R. Fan'auto et al., Mat. Res. Soc. Symp. Proc., 344(1994) 101 7 Seiyama, T. Catal Rev Sci. Eng.,. 34 (1992) 281 8 M.Machida et al., J. Catal., 120 (1989) 377 9 G. Crroppi et al., Applied Catalysis, 104 (1993) 101 10. G. Groppi et al., J. Mat. Sci., 29 (1994) 3441. 11. M. Daturi et al., Chem. Mater., in press 12. G. Groppi et al, J. Solid State Chem., 114 (1995) 326 13. G. Busca et al., Catal Lett., 31 (1995) 65 14. G. Groppi et al., Actas XIV Simposio Iberoam. de Catalisis, 12-16 September, (1994) Concepcion (Chile), Vol II, p. 733 15. J.G. Mc Carty and H. Wise, Catal. Today, 8 (1990) 231. 16. H. Arai et al., Appl. Catal., 26 (1986) 265. 17. J.M.D. Tascon and L. Gonzalez Tejuca, React. Kinet Catal. Lett., 15 (1980) 185. 18. G.Groppi et al., A.I.Ch.E. Journal, 41 (1995) 2250. 19. T. Mitani and F.A. Williams, Combustion and Flame, 39 (1980) 169 20. Dryier, F. L., and I. Glassmann, 14th Symp. on Combustion, The Comb. Inst., (1972) 987

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis -40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

483

H y d r o c a r b o n activation and oxidation on transition metal m i x e d oxides: Ft-IR a n d f l o w r e a c t o r studies E. Finocchio a'b, R. J. Willey a, G. Ramis b, G. Busca b and V. LorenzeUi b aDepartment of Chemical Engineering, Northeastern University, Boston, Ma 02115, USA blstituto di Chimica, UniversitY, P.le J.F. Kennedy, I-16129 Genova, Italy

The gas-phase partial and total oxidation of hydrocarbons (light alkanes, light alkenes, alkylaromatics) over transition metal oxide catalysts has been investigated using FT-IR spectroscopy of the adsorbed species in an IR cell, and in a flow reactor. It is proposed that both total and partial oxidation occur at the expense of nucleophilic lattice oxygen species. CH activation gives rise to alkoxy groups that have an increasing carbocationic nature by increasing the possibility to delocalize the cationic charge. A good parallelism is found between our data and with those are found in solution with similar oxidizing species.

1. I N T R O D U C T I O N Transition metal oxides represent a prominent class of partial oxidation catalysts [1-3]. Nevertheless, materials belonging to this class are also active in catalytic combustion. Total oxidation processes for environmental protection are mostly carried out industrially on the much more expensive noble metal-based catalysts [4]. Total oxidation is directly related to partial oxidation, athough opposes to it. Thus, investigations on the mechanism of catalytic combustion by transition metal oxides can be useful both to avoid it in partial oxidation and to develop new cheaper materials for catalytic combustion processes. However, although some aspects of the selective oxidation mechanisms appear to be rather established, like the involvement of lattice catalyst oxygen (nucleophilic oxygen) in Mars-van Krevelen type redox cycles [5], others are still uncompletely clarified. Even less is known on the mechanism of total oxidation over transition metal oxides [ 1-4,6]. Previous studies revealed that the spinel Co304 is the most active among binary oxides in catalytic combustion [7], although it is not stable at high temperatures [8]. Other spinel-type mixed oxides, like chromites [9], and perovskite-type mixed oxides [4] are also very active and rather stable. Due to their high activity, these materials convert light hydrocarbons at low temperatures in oxidizing conditions. For this reason they can be successfully used in spectroscopic investigations of hydrocarbon oxidation [10-12]. We have summarized below recent results concerning spectroscopic / flow reactor investigations of hydrocarbons partial and total oxidation on different transition metal oxide catalysts. The aim of this study is to have more information on the mechanisms of the catalytic activity of transition metal oxides, to better establish selective and total oxidation ways at the catalyst surface, and to search for partial oxidation products from light alkane conversion.

484 2. EXPERIMENTAL

Transition-metal mixed oxides active in combustion catalysis have been prepared by two main procedures: i) classical coprecipitation / calcination procedures starting from metal nitrates and/or alkoxides; ii) preparation based on the supercritical drying of gels prepared from organic complexes (alkoxides, acetylacetonates or acetates), producing "aerogels". Details on the second preparation can be found in Ref. 13. Catalytic combustion experiments have been performed in a flow reactor operating below the lower explosion limits, using HC/OE/He mixtures. The product analysis was done by gas chromatography. FT-IR spectra have been recorded with a Nicolet Magna 750 instrument, using conventional IR cells connected with evacuation-gas manipulation apparatus. The powder was pressed into self-supporting disks, calcined in air at 773 K and outgassed at 773 K for 20 minutes before experiments. 3. RESULTS AND DISCUSSION

3.1. Characterization of the oxidized and reduced catalyst surfaces and the active oxygen species. In the cases of the selective oxidation reactions over metal oxide catalysts the so-called Mars-van Krevelen or redox mechanism [4], involving nucleophilic oxide ions 02, is widely accepted. A possible role of adsorbed electrophilic oxygen (molecularly adsorbed 02 and / or partially reduced oxygen species like 02, 022. or O) in complete oxidation has been proposed by Haber [2]. However, Satterfield [1] queried "whether surface chemisorbed oxygen plays any role" in catalytic oxidation. In order to have more information on the nature of the oxygen species active in partial and total oxidation we investigated the interaction of the hydrocarbons with the pre-oxidized surfaces of oxides where different types of surface oxygen species are formed. In particular we investigated p-type semiconductors like chromia, chromites and cobalt oxide C0304. Moreover, we studied n-type metal oxides like Fe203, metal ferrites and Cut-based catalysts. IR spectra give direct evidence concerning the state of oxygen on these oxides. In fact, oxygen is in excess on the surface of oxidized chromia and chromites, where high valency cations, essentially Cr 6+, give rise to chromate species (CrO42). These species are evident in the IR spectra, because they are responsible for typical Cr=O stretching bands in the region 1020-800 cm ~ [10,11,14], that are completely disappeared in reduced stoichiometric chromites. Instead, C0304 when fully oxidized is almost completely opaque to the IR radiation, but it becomes very transmitting by reduction with hydrocarbons and hydrogen [12,15]. According to the literature [16], part of Co 2§ in stoichiometric C0304 lose one electron in the presence of oxygen, so being oxidized to Co 3§ associated to excess oxygen in the form of additional interstitial 02. ions. Excess 02. can easily exchange one electron with Co 3+ giving rise to Co 2+ - O spe~es, i.e. electrophilic oxygen is easily generated, as reported in the literature [2]. In this way the empty orbitals of Co 3§ act as acceptor levels and holes in the O 2p valency band are easily generated, likely responsible for the absorption of the IR radiation. This picture finds support in the strong effect of the temperature in varying the IR light transmission by the oxidized sample, strongly increased at low temperature. On the contrary, n-type semiconductors tend to become opaque when oxygen deficient, as a result of chemical reduction or of simple outgassing at high temperature, as we observed for Fe203, ferrites and C u t , as well as for V205 and WO3 [17,18]. The optical behaviour of these

485 slightly reduced oxides can be attributed to the absorption by the additional electrons of the reduced cationic centers that at room temperature in part already lie in the conduction band and are, consequently, delocalized into the entire solid particle. According to the literature, the interaction of olefms with electrophilic and nucleophilic oxygen species is expected to give different products [2]. In all cases, including on Co30,~x (i.e. on a catalyst where electrophilic oxygen should be present at the surface) the products of propene oxidation are those we expect for interaction with nucleophilic (lattice) oxygen. This is shown in Fig. 1 where the FT-IR spectra of the adsorbed species arising from the interaction of propene and acrylic acid on Co30,~x at 373 K are compared. The spectra show the features of the same species, i.e. acrylate ions, that act as intermediates for the total oxidation of propene. This supports the conclusion that lattice oxygen can be involved in both partial and total oxidation over transition metal oxide catalysts, on both n- and p- semiconducting oxides.

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Fig. 1. Ft-IR spectra of the adsorbed species arising from the interaction of Co30~x with propene (upper spectrum) and acrylic acid (lower spectrum) both at 373 K. 3.2. Partial oxidation vs combustion catalysis. Shown in Fig. 2, left, are the light-off curves relative to the combustion of the C3 compounds propane, propene, acrolein and acetone on MgCr204. The oxidation of these organic compounds is substantially non-selective, although acrolein can be found from propene oxidation at low conversions (selectivity 34 % at 10 % conversion with C3He : 02 < 1/20 at 225 ~ Traces of oxygenates can be found for propane oxidation at 270 ~ in oxygen excess at 10-15 % conversion. To determine whether this catalyst is able to perform selective oxidation, we investigated a facile reaction on it, i.e. the oxidation of 2-propanol (Fig. 2, right). Acetone is produced from 2-propanol with high selectivity (> 90%) in the temperature range 250-300 ~ where conversion is up to 30 %. In these conditions the main byproduct is propene, produced by propanol dehydration. Above 300 ~ (the temperature at which acetone, when feeded, strats to be oxidized) the selectivity to acetone from 2-propanol falls to below 10 % and is zero at 350 ~ when conversion approaches 100 %. Thus MgCr204 can give rise to partial oxidation of organic compounds like alcohols at relatively low temperature.

486 However, in the same temperature range and 0 2 partial pressure total oxidation of acrolein and propene largely predominates. This can be taken as a further support that on transition metal oxide catalysts the same oxygen species (lattice oxygen) are involved in both partial and total oxidation. It seems valuable to remark the similarity of the behavior of the chromite catalysts, whose surface is covered by chromate species in oxidizing conditions, with that of chromate ions and chromic acid in solution, that are choice oxidizing agents for the very selective conversion of alcohols into carbonylic compounds [19]. The same reagents under more drastic conditions cause C-C bond oxidative cleavage and the deep oxidative destruction of organics. This allows to count the methyls of organic compounds because they, after oxidative chain clevage, give rise to acetic acid [20]. This corresponds to the formation of acetates that are detected as the most stable adsorbed species on chromite catalysts after alkane oxidation [10-12].

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Fig. 2. Left: catalytic oxidation of C3 organic compounds over MgCr204. Conversion of [] propane; A acetone; X acrolein; 9 propene. Right: catalytic oxidation of 2-propanol over MgCr204. 9 conversion of 2-propanol; selectivities to II acetone; A propene; X COx. 3.3. Surface oxidation pathways: the activation at the C-H bonds and the nature of "surface alkoxide species". The interaction of alkanes with the oxidized surfaces of combustion catalysts gives rise to oxygenate species. For the less reactive hydrocarbons (i.e. methane and ethane) the reaction is detected only at rather high temperature, and carboxylate species (acetates from ethane) or carbonate species are found as the adsorbed products. With more reactive hydrocarbons like propane and butane, carbonyl compounds (acetone and methyl-ethyl-ketone) have been observed at lower temperature on MgCrEO,~x. If the reactant is even more active, features typical of alkoxide species have been observed. This is shown in Fig. 3 where the spectrum of the adsorbed species arising from isobutane interaction with MgCrEO,~x at 423 K are compared with those obeserved after interaction of tert-butanol over the same catalyst in the same conditions. The bands observed in the region 1300-1000 cm ~ are typical of the C-O / CC stretchings of tert-butoxide species. The detection of these species parallels the detection of allyl-alkoxides from propene, over the same catalyst [10,11]. The rate of the oxidative cleavage of C-H bonds over this catalyst nearly follows the order: benzyl --- allyl > methine >

487 methylene > methyl > methane, and has been related either to the dissociation energy of the C-H bonds (inversely) [11] or to the electrostatic potential charge (the more negative the charge, the higher the cleavage rate) on hydrogen [21]. In any case, this trend is the same previously observed for the combustion rate on copper oxide catalysts [2,22], for which the first C-H cleavage step is thought to be rate-determining. Again we remark that the activity of chromite catalysts (whose surface is covered by chromates) closely corresponds to the activity of chromates in acid solution, that are able to oxidize tertiary alkanes to the corresponding tertialy alcohols and secondary alkanes to the corresponding ketones [20]. Interestingly, the oxidation of chiral tertiary alkanes by chromic acid in solution gives rise to the corresponding alcohols stereospecifically, with retention of the configuration [23]. ~.~-. J

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.

Fig. 3. FT-IR spectra of the adsorbed species arising from the interaction of (a) tert-butanol and (b) isobutane over a combustion catalyst (MgCr204) at 423 K, and from tert-butanol (373 K, c), isobutene (300 K, d) and isobutane (380 K, e) on a selective oxidation catalyst. In Fig. 3, on the left, the spectra of the adsorbed species arising from isobutane interaction with a molybdena-based catalyst that behaves quite selectively for the production of methacrylic acid from isobutane are compared with those arise from isobutene and tertbutanol adsorption on the same surface. In all three cases the typical features of tert-butoxy species are observed, that are stable up to near 400 K. According to our data, isobutane is activated by C-H bond oxidative cleavage giving rise to tert-butoxide that can decompose to isobutene and an OH group. Isobutene can later undergo aUylic oxidation on selective catalysts, while it can undergo oxidative C-C bond breaking to give one acetate and two formate species on the combustion catalysts. These data show that the C-H activation mode is similar on both partial and total oxidation catalysts. Starting from n-butane, 2-butoxides that rapidly convert to 2-butanone are found over MgCr204 [24]. However, the further oxidation of adsorbed 2-butanone only gives rise to the acetate species, while starting from n-butane, formate species are also observed. This can be explained assuming that sec-butoxides can partly isomerize to tert-butoxides before further oxidation. This implies that the C-O bond formed is partly ionic and the alkyl moiety has the

488 character of a carbenium ion. Similarly, the allyloxy species we observed to be formed by oxidation of propene can have the pronounced nature of a an ally1 cation, stabilized by the 1,3 charge delocalization. It is well known that symmetric aUyl species are involved in propene selective oxidation and ammoxidation to acrolein and acrylonitrile, and that no C-O bonds is formed irreversibly between the first and the second hydrogen abstraction in acrolein synthesis [ 1-3,25]. The picture we propose supports the idea that the adsorbed allylic species involved in propene oxidation is cationic in nature. A further support to the cationic nature of allyl species observed on the surface of chromite catalysts is given again by the comparison with the behavior of chromates in solution, that also give rise to aUylic oxidation and where transpositions of intermediate aUylic carbenium ions to more stable tertiary carbocations have been reported [20]. As for selective catalysts for acrolein production, the allyl species intermediates have been proposed to be radical-like, although allyloxy- species are assumed to be formed later [3,25]. In effect desorption of gas-phase allyl radicals has been observed from BiEMoO6 (a very selective catalyst) and, mostly, from Bi203, that is not a selective catalyst for acrolien production [26]. Over BiE(MoO4)3 allyl radicals were not formed while MoO3 (that allows acrolein production) acts a a sink for allyl radicals [26]. It seems quite reasonable to think that the allyl radicals formed on bismuth centers are intermediates favoring the formation of allyl carbeniurn ions / allyloxy- species on molybdenum. This picture also agrees with the detection of "isolated" benzyl species (with no C-O bonds) upon toluene and o-xylene selective oxidation over vanadia-titania selective oxidation catalysts [27], according to the very strong delocalization of the cationic charge over the aromatic ring. Due to this stability, the rate detrmining step in toluene oxidation is apparently shifted to the successive step, the reaction of benzyl species with surface oxide to give benzaldehyde [27]. In conclusion, the carbocationic character of the intermediate and its stability should decrease and the alkoxide character of the same species should increase following the sequence benzyl, allyl, saturated tertiary, secondary, primary.

3.4. Surface oxidation pathways: C3 hydrocarbons. IR studies have shown that propane and propene oxidation over a chromite catalyst follow two different paths constituted by partial oxidation steps that give rise to adsorbed carboxylate species; these species burn totally before leaving the surface [ 10]. The spectra of the adsorbed species arising from propene oxidation at r.t. over oxidized Co304, an even more active catalyst [ 12], are very similar to those arise from aUyl alcohol, acrolein and acrylic acid (see Fig. 1). They show that on Co304 the allylic methyl group of propene is attacked and oxidized to give the carboxylate ion acrylate. Also on MgCr204 acrylate species are formed from propene, although at higher temperature according to the weaker activity of this catalyst. Propane oxidation gives mainly rise to adsorbed acetone at low temperature on MgCr204, while on Coat4 propane gives rise to several different species, i.e. acetate, acrylate and propanoate species. On both surfaces, at the temperature at which the catalytic hydrocarbon conversion is sufficiently fast, the surface species disappear or are substituted by carbonate species, and CO2 gas is observed. In the following scheme, an oxidation pathway for propane and propene is proposed. This mechanism, that could be generalized to different transition metal oxide catalysts, implies that propene oxidation can follow the allylic oxidation way, or alternatively, the oxidation way at C2, through acetone. The latter easily gives rise to combustion, because it can give rise to enolization and C-C bond oxidative breaking. This is believed to be the main combustion way for propene over some catalysts, while for other catalysts acrolein overoxidation could

489 predominate. Consequently, propene combustion can be either successive or competitive to propene partial allylic oxidation. On the other hand, starting from propane the alkoxide intermediate can either decompose to propene or be oxidized to acetone and overoxidized later. So, the alkoxide evolution step can play the role of a "selectivity determining step" in the oxy-dehydrogenation of propane to propene, using an expression proposed by Kung and coworkers [28]. CH3CH2CH 3

CH3CH=CH2

~'4&

C~x~ /CH3 S

OH

c.

I

I

,c-2

.~~

"-

O

o 1

acetone

CH /+\ CH 2 CH 2

CH2=CH

C~x~ /CH3

CH31

C II O

oS C...~CH2

-..

f

O"

I

I

CH2=CH

~

~

CH ii 0

acrolein

,

I

acetic acid

,~

CH 3 C, O'O I

~

~

C Ox

.~~

CH2=CH ,C, OO ~

acrilic acid

Scheme I. A generalized pathway for C3 organics partial and total oxidation. 3.$. Surface oxidation pathways: C4 linear hydrocarbons. Similar experiments allowed us to propose a reaction pathway for n-butane oxidative conversion apparently common to selective oxidation catalysts like (VOhP2OT, that allows the production of maleic anhydride from butane; VEOs-TiO2 and MoO3-TiO2, that allow the production of acetic acid from butane and of maleic anhydride from butene; Mg3(VO4h that allows the oxy-dehydrogenation of butane to butene and butadiene; FeCrO3 that allows the oxy-dehydrogenation of butene to butadiene; and combustion catalysts [29]. The reaction network is substantially the same in all cases, but the different behavior of the catalysts is explained by the different rates of some alternative steps on the different surfaces. Again, a key step concerns the dehydration or the oxy-dehydrogenation of the 2-butoxide intermediate, playing the role of "selectivity determining step", where the oxy-dehydrogenation route diverges from the main combustion route. The low activity in oxidation and the significant basicity with absence of BrOnsted acidity typical of Mg vanadate are likely main factors favoring the production of butene from the 2-butoxy- species. On the contrary, the weak BrCnsted acidity of oxidized Mg-chromite is likely responsible for blocking of the 2-butoxide

490 species at the surface whilethe high oxidizing power of Cr 6+ causes its rapid oxydehydrogenation to the ketone and the following C-C bond breaking. The same mechanism proposed for the combustion catalyst Mg-chromite apply also to catalysts that allow significant yields in acetic acid from n-butane, like vanadia-titania, that accordingly also show a medium-high BrCnsted acidity. Being acetate ions intermediates in the combustion way, it is easily rationalized that the production of acetic acid is favored by the addition of steam in the reactant mixture and by adjusting the reaction conditions. The catalysts that allow the production of maleic anhydride from n-butane with high selectivity, like (VO)2P2OT, are characterized by a strong acidity, that, like a strong basicity, favors the decomposition of alkoxides to give the olefin and the diene. The catalysts that allow the production of maleic anhydride, either from n-butane or from butenes and butadiene, necessarily have particular sites that allow the insertion of oxygen atoms in the 1,4-position of butadiene. These sites are definitely absent on combustion catalysts. . . . . . . . . . . . .

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

o.o8

//

0,06

'\

!,'

0,04

\

..

,',

o;

0.02 0.00 = - 0 . 0 2 -~

/

-0.04

.~ ~""~"~--~ .... - 0 . 1 0 ~-"~"~-'~.N'~J'~...'-~-J'V,,'~

/

-o.,.~ ..... "---"'. . .<>.,,,__:

-

,

/

...

.I

/ .," _ ._/I ," ,]

- o . o 6 -Z

-0.14

i

-

~

.....

s "-4

x..

~

;;

\

:x.x

_ i ,, J

-0.18

\

! ' , ..

\d

,,_._; ...

"-..... <J I

-0.20

/,,,.,,'-\ ....

-..\ ..j' .,.,/,,,i

',,I::!

.....',., (c) ""\

f~ '

(d)

\\

\ " ',.

[~

/

/ ""

:' 'i ,,

, #

/\ .,h,,,-"~..\

,,

'""

1

t"')""'~..., ~'-'''''F " ----v-,, /-,,,~,--,,-~..

'

t/ """

"'f

'

I

-0.~ ~

. . . . ~, .. , ~ , ,

-0.24 ~ 2000

/\ :,,,"

1950

1900

1850

1800

1750

1700

1650

1600

1550

i

1500

Wavenumberl

1450 1 4 0 0 (ore-1)

1350

1300

1250

1200

~,7

f

,-.'' 1150

\1

1100

10.50

Fig. 4. Ft-IR spectra of the adsorbed species arising from interaction of 1-butene (a,b) and butadiene (c,d) on MgCr2Oa+x at r.t. (a,c) and at 373 K (b,d). Accordingly, the interaction of butadiene with (VO)2P207, V2Os-TiO2, and MoO3-TiO2 is strong and reactive at r.t. (giving rise to furan-like species) and, at higher temperatures, gives rise to the typical bands of cyclic anhydrides in the region 1900-1700 cm l [29,30]. This reactivity is lacking on some very reactive catalysts like MgCr204+x where both butene and butadiene give rise to molecular adsorbed species at r.t. (Fig. 4, a and c), while at slightly higher temperatures methyl vinyl ketone is formed from butenes (typical doublet at 1670, 1640 cm q, vC=O and vC=C, and band at 1180 cm l, Fig. 4,b) and a mixture of carbonyl compounds and carboxylates from butadiene (Fig. 4,d). As a conclusion of our experiments the following generalized pathway is proposed for n-C4 oxidations:

491

CH3COOH

O

CO, CO 2

O

HOOC

o,

0

,o

O

>

o

o--~~-o 0

OO

-I~

~~I~OOH

o---~~-o 0

Scheme II. Reaction pathways for butane oxidation. 4. CONCLUSIONS The results of the experiments we carded out and the comparison with the data arising from homogeneous oxidation chemistry allow us to propose and generalize the following conclusions that in part contrast previous literature findings: i) activation of hydrocarbons occurs over transition metal oxides by abstraction of two electrons by the oxidized cationic center. ii) nucleophilic oxygen species (lattice oxygen) contribute to the C-H activation step by interacting with the resulting carbenium ion and proton species, giving rise to an alkoxide and to an hydroxide, respectively. iii) the resulting alkoxide has a partial carbocationic nature, the more important the higher is the charge delocalization (i.e. for benzyl and allyl species). iv) so, the C-O bond of the alkoxide intermediate is not irreversibly formed, and its reversible rupture justifies the symmetry of the aUyl species in propene oxidation. v) the alkoxides so formed can be further oxidized to carbonylic compounds. vi) if the surface and the carbonyl compounds are both sufficiently weakly reactive, desorption can occur and partial oxidation takes place.

492 vii) if the surface is very reactive and the carbonyl compounds has hydrogen in the ~position, enolization occurs that finally gives rise to C-C bond breaking and either to combustion or to the production of acetic acid. viii) total combustion pathways should involve nucleophilic (lattice) oxygen species and a sequence of partial oxidation reactions that finally give overoxidation of the products to COx. The authors acknowledge MURST (Rome, Italy) and NATO for financial support. REFERENCES

1. .

3. 4. .

6. .

.

9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28 29. 30.

C.N. Satterfield, Heterogeneous Catalysis in Industrial Practice, 2nd ed., McGraw Hill, New York, 1991. A. Bielanski and J. Haber, Oxygen in Catalysis, Dekker, New York, 1991. V.D. Sokolovskii, Catl. Rev. Sci. Eng. 32 (1990) 1. M.F.M. Zwinkels, S.G. Jaras, P.G. Menon and T.A. Griffin, Catal. Rev. Sci. Eng. 35 (1993) 319 P. Mars and D.W. van Krevelen, Chem. Eng. Sci. 3 (special supplement) (1954) 41. J.J. Spivey, Ind. Eng. Chem. Res., 26 (1987) 2165; J.J. Spivey, in "Catalysis", vol 8, The Royal Society of Chemistry, Cambridge, 1989, pag.158. J.E. Germain and L. Laugier, Bull. Soc. Chim. France, (1972) 541 and 2910;. J.E. Germain and R. Perez, Bull. Soc. Chim. France, (1972) 2042 and 4683. E. Garbowski, M. Guenin, M.-C. Marion and M. Primet, Appl. Catal. 1990, 64, 209 L. Ya. Margolis, Advan. Catal. Relat. Subj. 14 (1963) 493 E. Finocchio, G. Busca, V. Lorenzelli, R.J. Willey, J. Chem. Soc. Faraday trans., 90 (1994) 3347. E. Finocchio, G. Busca, V. Lorenzelli and R.J. Willey, J. Catal. 151 (1995) 204. E. Finocchio, G. Busca and V. Lorenzelli, J. Chem. Soc. Faraday trans., submitted G. Busca, M. Daturi, E. Kotur, G. Oliveri and R.J. Willey, in "Preparation of Catalysts VI", G. Poncelet et al. eds., Elsevier, Amsterdam, 1995, p. 667 K. Hadjiivanov and G. Busca, Langmuir, 10 (1994) 4534 G. Busca, R. Guidetti and V. Lorenzelli, J. Chem. Soc. Faraday Trans. 86 (1990) 989. G. Tulyev and S. Angelov, Appl. Surface Sci. 32 (1988) 381; B. Marcus-Saubat, J.P. Beaufils and Y. Barbaux, J. Chim. Phys. 83 (1986) 317. G. Busca, G. Ramis and V. Lorenzelli, J. Mol. Catal. 50 (1989) 231 G. Busca, Catalysis today, in press A.H. Haines, Methods for the Oxidation of Organic Compounds, Academic Press, 1988. G.A. Olah and A. Molnar, Hydrocarbon Chemistry, Wiley, New York, 1995. G. Busca, E. Finocchio, G. Ricchiardi and G. Ramis, submitted paper G.K. Boreskov, in Catalysis Science and Technology, J.R. Anderson and M. Boudart eds., Vol. 3, Springer Verlag, New York, 1982, p. 39. K.B. Wiberg and G. Foster, J. Am. Chem. Soc. 83 (1961) 423. E. Finocchio, G. Busca, V. LorenzeUi and R.J. Willey, submitted paper T.P. Snyder and C.G. Hill, Jr., Catal. Rev. Sci. Eng. 31 (1989)43 D.J. Driscoll, K.D. Campbell and J.H. Lunsford, Advan. Catal. 35 (1987) 139. G. Busca, J. Chem. Soc. Faraday trans. , 89 (1993) 753; G. Busca, in "Catalytic Selective Oxidation", S.T. Oyama and J.Hightower eds., ACS, Washington, 1993,p. 168 H.C. Kung and M. Kung, Advan. Catal. Relat. Subj. 40 (1994) G. Busca, G. Ramis, V. Lorenzelli and G. Oliveri, in "New Developments in Selective Oxidation II", V. Cortes Corberan and V. Bellon ed., Elsevier, 1994, p. 253, G. Ramis and G. Busca and V. Lorenzelli, J.Mol. Catal. 55 (1989)

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

493

Biomimetic oxidation on Fe complexes in zeolites Gennady I. Panov, Vladimir I. Sobolev, Konstantin A. Dubkov and Alexander S. Kharitonov Boreskov Institute of Catalysis, Novosibirsk, 630090 Russia

One-step hydroxylation of aromatic nucleus with nitrous oxide (N20) is among recently discovered organic reactions. A high efficiency of FeZSM-5 zeolites in this reaction relates to a pronounced biomimetic-type activity of iron complexes stabilized in ZSM-5 matrix. N20 decomposition on these complexes produces particular atomic oxygen form (a-oxygen), whose chemistry is similar to that performed by the active oxygen of enzyme monooxygenases. Room temperature oxidation reactions of a-oxygen as well as the data on the kinetic isotope effect and Moessbauer spectroscopy show FeZSM-5 zeolite to be a successful biomimetic model.

1. INTRODUCTION A perfect operation of enzyme monooxygenases provides a wide variety of fine oxidation reactions occurring in the living nature at ambient conditions. This remarkable phenomenon relates to a unique ability of monooxygenases (MO) to activate dioxygen and to produce very reactive oxygen species. These species coordinated on the iron active sites can selectively add to organic molecules to yield oxygenated products [ 1]. Man-made catalysts are not so perfect as enzymes. One often needs to perform several stages under harsh conditions to accomplish a simple chemical reaction. As an example we may mention two oxidation reactions, which will be discussed further, i.e. benzene oxidation to phenol via the three stage cumin process, and methane oxidation to methanol via the intermediate CO and 1-12production. MOs elegantly do these reactions by a single move just adding one oxygen atom to a starting molecule. Researchers take strong efforts to understand the mechanism of MO oxidation and to apply biomimetic strategy in the practice. Progress in this field is strongly related to the development of successful models capable to mimic MO activity. In this paper we for the first time consider biomimetic aspects of a recent approach in the catalytic oxidation dealing with the iron-containing zeolites as catalysts and nitrous oxide (N20) as a source of oxygen supply [2]. The paper will refer rather frequently to the previous results obtained in this laboratory, since during several years our research group was being intensively involved in the N20 oxidation study. Further we shall show that analyzing the published data and new results concerning the room temperature methane and ethane oxidation we have found Fe complexes in ZSM-5 matrix to be a promising biomimetic model.

Acknowledgment. Research described in this paper was made possible in part by Grant No JDB 100 from International Science Foundation and Russian Government.

494 2. AROMATICS

HYDROXYLATION WITH NITROUS OXIDE

2.1. Discovering zeolite catalysts Partial oxidation reactions are usually carried out over transition metal oxides capable of changing their valent state during their interaction with reacting molecules. Naturally, zeolites with their alumina-silicate composition did not prove themselves as good oxidation catalysts. They failed also to serve as efficient catalyst supporters, since transition metals being introduced into the zeolite matrix lose their ability to activate dioxygen [3,4]. In the 1980's zeolites attracted a renewed attention. They were shown to be rather promising catalysts if, instead of 02, a chemically pre-modified oxygen entering the oxygen-containing molecules is used. The most known example is an excellent catalytic performance of titanium silicalites in the liquid phase oxidations with H202 [5]. A gas phase oxidation with nitrous oxide is another approach in this field being intensively developed in the last years [2]. Nitrous oxide as an efficient oxygen donor was noticed when used in such a delicate reaction as the direct oxidation of benzene to phenol: C6H6 + N20

=

C6HsOH + N2

(1)

This reaction was first demonstrated over V, Mo and W oxides [6]. At 823 K vanadium oxide provided phenol selectivity up to 71%, which was much higher than it had been ever achieved with 02. This result stimulated further efforts in searching for more efficient catalytic systems. As a result, in 1988 three groups of researchers [7-9] have independently discovered ZSM-5 zeolites to be the most efficient catalysts. They allowed the reaction to proceed at much lower temperature (573-623 K) with nearly a 100% selectivity. Later, more complex aromatic compounds were also hydroxylated in this way [2]. These results being quite untypical for zeolites give rise to a number of fundamental questions: i) what makes the zeolite to function as an active catalyst; ii) what makes N20 to function as a selective oxidant; iii) what is the reaction mechanism. We shall shortly discuss the situation with these issues because of their importance for our further consideration.

2.2. Role of iron A strong zeolite acidity had been first suggested to explain the catalytic activity origin [7]. Later, it was not supported experimentally. The activity was shown to relate to Fe admixture always presented in zeolites in a small amount. The role of iron is clearly seen in Fig 1. The later presents the data obtained with two ZSM-5 zeolite systems of Fe-Si [ 10] and Fe-AI-Si [ 11] composition with a widely varied Fe concentration. In fact, the rate of benzene oxidation to phenol over the purest samples (0.003 and 0.004 wt % Fe2Oa) is very low and strongly increases with the increasing iron content. The catalytic activity of Fe proved to be very high (especially in the AI-Si matrix) so that an admixture of 0.01 wt % Fe2Oa may cause a noticeable benzene conversion. This discovery was quite unexpected, since iron oxide has been never reported as an active catalyst in either partial or full oxidation. The studies of two simplest reactions, i.e. 02 isotopic exchange and N20 decomposition, revealed a dramatic change of Fe properties in the ZSM-5 matrix compared to Fe203 [4]. Fe atoms lose their ability to activate 02 but gain remarkably in their ability to activate N20. It gives rise to a great effect of the oxidant nature in the reaction of benzene oxidation over the FeZSM-5 zeolite (Table 1). Thus, in the presence of N20 benzene conversion is 27% at 623 K, while in the presence of 02 it is only 0.3% at 773 K. And what is more, there is a perfect change of the reaction route. Instead of selective phenol formation with

495 N20 (S=98%), their are only full oxidation products with 02. Over Fe203 none of the oxidants produces phenol.

~ 30 o

623 K -

O I

30[

O

--

=

20~

o

20

~

623K

"

o 573 K

O

o 10 [ Fe-Si ] 0

I

I

I

I

2

3

4

5

0

Fe concentration ( wt %Fe203)

0.2

0.4

0.6

Fe concentration ( wt % Fe203)

Figure 1. Effect of iron concentration on FeZSM-5 catalytic activity. Beside iron, the catalytic properties of many other transition metals (V, Mo, Cr, Mn, Co, Ni, Cu, Ti, Zn, Pd, Pt) have been tested. These metals exhibited no activity in phenol production [7,11 ]. This means that Fe might be the one of few particular elements or even the only one, which can effeciently catalyze this reaction. Table 1 Effect of oxidant nature on benzene oxidation (X-benzene conversion, S-selectivity to phenol) Sample

Oxidant N20

Fe-AI-Si (0.08 wt % FEE03) Fe203

Oxidant 02

T (K)

X (%)

S (%)

T (K)

X (%)

623

27

98

773

0.3

0.0

623

24.5

0.0

623

5.5

0.0

S (%)

2.3. Particular features of nitrous oxide In order to understand the reason for such a beneficial N20 oxidizing effect, a detailed mechanism of its decomposition as a stage supplying oxygen to the surface has been studied [4,12]. This study revealed a special type of iron active sites in ZSM-5 matrix (called a-sites), which decompose N20 producing a new oxygen form (a-form): N20 + ( )or

-(O)ct

+N2

(2)

At temperatures lower than 573 K a-oxygen is termally stable and reaction (2) selectively occurs with no oxygen evolution into the gas phase (Fig 2). Note that this phenomenon is not the result of oxygen consumption for the surface reoxidation, since our experimental conditions completely exclude a reduction of the sample alter its treatment in O2 [ 12]. When all ix-sites are occupied, the reaction terminates. By measuring the amount of N2 produced (or that one of

496

12~

N20 + ( )~, --) N2 + (O)~'

90

I

N~ o

60

Figure 2. Kinetics of N20 decomposition at 523 K followed by (x-oxygen loading on FeZSM-5 zeolite surface.

30

0

4

8 12 Time (min)

16

N20 consumed) one can determine an amount of (x-oxygen loaded and the density of (x-sites. To verify these results, the reaction of isotope exchange can be additionally used to measure (xoxygen amount [ 12]"

~SO2 + (~60)~

=

~602 + (~SO)a

(3)

(x-Site density can be regulated by the iron content and procedure of FeZSM-5 activation. High density of (x-sites provides a good experimental opportunity to study properties of (xoxygen, what has been done in a number of papers [4,12-15]. A very low energy of bonding with the surface and a very high reactivity are the most remarkable features of (x-oxygen. At room temperature, it participates in the isotopic 02 exchange as well as in oxidation of various organic molecules. (x-Oxygen has been probably first observed when studying the mechanism of N20 decomposition over FeM zeolite [16], though its concentration was too low for its clear manifestation. 2.4. Reaction mechanism

After discovering (x-oxygen formation as a particular feature of N20 decomposition, an important question concerning the reaction mechanism arises: does (x-oxygen participate or not participate in the oxidation of benzene to phenol? According to a generally accepted view [ 17], the surface oxygen providing partial oxidation must not have a low energy of bonding to the surface, and must not exhibit a high reactivity, which is in a conspicuous contrast to (x-oxygen properties. Therefore, the idea to relate (x-oxygen to phenol formation needs a strong experimental support. Attempts to use for this purpose some spectroscopic techniques 0IL NMR) have not provided reliable data because of their low sensitivities. Another approach was found to be successful. If the interaction of benzene with (x-oxygen actually produces phenol, and if this phenol can be extracted from the surface, its amount will be sufficient for reliable chromatographic analysis. Based on this idea, experiments were carried out according to the following three stage scheme [18]:

497

1. N20 + ( )or 2. Cd-I6 + ( O )or 3. (Cd-I~OH)ct

520K

)

(O)tx + N2

.~

(Cd-IsOH)ot

;

Cd-IsOH

298 K 298 K

+

()~

The scheme includes a-oxygen loading to the surface (stage 1), its interaction with benzene at room temperature (stage 2), and product extraction from the surface (stage 3). Results obtained with the Fe-AI-Si sample are given in Table 2. Phenol was found to be the only reaction product, and its average amount corresponded to 90% of the a-oxygen loaded. A somewhat underestimated yield is probably related to the incomplete phenol extraction. Similar results were obtained with ZSM-5 sample of Fe-Si composition [ 18]. Table 2 Room temperature benzene oxidation with a-oxyBen Sample

FeZSM-5 (0.07 wt % Fe203)

Run

(O)a (~mol/g)

C6I-~OH

Phenol yield

(p.mol/g)

(%) 93 85 93

1

6.0

5.6

2 3

5.5 5.5 0.0

4.7 5.1 0.0

4 (blank run)

-

These experiments clearly showed that it is a-oxygen participation that provides FeZSM-5 zeolites with such a remarkable catalytic performance in the reaction of benzene to phenol oxidation. Equations (1-3) written above are the main stages of the reaction mechanism.

3. BIOMIMETIC FEATURES OF FeZSM-5 ZEOLITE 3.1. cz-Oxygen m e t h a n e oxidation Results on the room temperature oxidation of benzene to phenol are rather suggestive with respect to the monooxygenase mimicking, since the hydroxylation of aromatics is one of typical reactions catalyzed by these enzymes [19]. Understanding the mechanism of MO oxygen activation is a difficult problem and is a subject of many studies and debates. Using the complexes of iron and other metals, including those encapsulated into zeolites [20-22], some successful models have been developed to simulate the MO function of cytochrome P-450. But all attempts to mimic methane monooxygenase (MMO) have failed because of not enough powerful oxygen activation. The MMO dinuclear iron sites are capable to produce oxygen species of much superior activity compared to other MOs. Therefore, in addition to aromatics and various other compounds, MMO exhibits a unique ability to hydroxylate methane producing methanol [23]. That is why, it is of special interest to test the a-oxygen oxidizing potential with respect to methane.

498

J

120 I00

Temperature ( K )

Temperature ( K )

300 400 500 600 700 800

300 400 500 600 700 800

i

AB ~ ,

.

|

J

I

-

i

i

1

i

120 ~

,

~

CH4

100

a)

80

80

~

60

60

o

40

40

~

co

20

,

|

~_

2'0

4'0

6'0

T i m e ( m i n)

8'0

i

.

i

,

1

CH4

.

|

.

b)

CO

20 0

I~

.

AB -]

,

0

20

40

.

,

.

60

,

.

gO

T i m e (m i n)

Figure3. Temperature-programmed reaction of methane with FeZSM-5 surface before a-oxygen loading (a) and after a-oxygen loading (b). A - time moment of opening the microreactor; B - time moment of switching on the programmed heating (6 K/s). For this purpose we studied a temperature-programmed interaction of CI-L with a-oxygen. Experiments were carried out in a static setup with FeZSM-5 zeolite catalyst containing 0.80 wt % Fe203. The setup was equipped with an on-line mass-spectrometer and a microreactor which can be easily isolated from the rest part of the reaction volume. The sample pretreatment procedure was as follows. After heating in dioxygen at 823 K FeZSM-5 cooled down to 523 K. At this temperature, N20 decomposition was performed at 108 Pa to provide the a-oxygen deposition on the surface. After evacuation, the reactor was cooled down to the room temperature, and CI-I4 was fed into the reaction volume at 108 Pa. Fig 3 shows the results of two temperature-programmed experiments. In the first (blank) experiment CH4 reacts with a "bare" FeZSM-5 zeolite, while in the second one it reacts with the zeolite after a-oxygen loading on its surface. Obviously, the bare surface is quite inert towards methane (Fig 3a): after reactor opening a weak CI-I4 adsorption occurs at room temperature. A slight heating results in a complete recovery of the CI-L pressure. There is quite a different picture in the a-oxygen presence (Fig 3b). Just after reactor opening a large and irreversible CH4 consumption occurs now, which evidences its chemical reaction with a-oxygen to take place at room temperature. The reaction product is strongly bound to the surface, and can not desorb into the gas phase under heating without destruction, which is accompanied by the CO evolution. In order to identify the product, we used a procedure of its extraction from the surface similar to that used in the case of a-oxygen benzene oxidation [18]. For this purpose, a number of single-turn-over runs in the room temperature methane oxidation were carried out according to the following scheme: 1) a-Oxygen loading to FeZSM-5 surface by N20 decomposition at 523 K; 2) Its interaction with CH4 (108 Pa) at room temperature; 3) Sample unloading and product extraction by an acetonitrile-water mixture with its further NMR and chromatographic analysis. Results are shown in Table 3. In all runs methanol has been detected as a sole product and its amount within the experiment error correlates with that of CH4 consumed. It is of interest that methane reaction with a-oxygen:

499 CH4 + (O)(x

=

(CH3OH)a

(4)

occurs surprisingly fast. We failed to measure its rate not only at room temperature but even after sample cooling down to 243 K. Similar experiments on a-oxygen oxidation were also carried out with ethane resulting in a selective formation of ethanol. Table 3 Room temperature ,methan,e. oxidation with a-oxygen Run CI-h reacted CH3OH formed No (gmol/g) (pmol/g) 20 23 18 20

1 2 3 4

CH3OH yield

(%)

19 21 18 18

95 91 100 90

3.2. Kinetic isotope effect A rate-determining step of MO alkane oxidation involves the cleavage of C-H bonds, which brings about high values of kinetic isotope effects (KIE = ka/kD) [24]. We measured KIE for methane oxidation according to reaction (4) using the intramolecular competition of C-H and C-D bonds of CH2D2 molecules in their reaction with a-oxygen. Product methanol was extracted with a deuterated mixture of acetonitrile and water, and its isotope composition was analyzed by the M R spectroscopy. Unlike monooxygenases, FeZSM-5 zeolite is a robust system and may be studied under wide conditions. In a temperature range of 223 - 373 K KIE value for methane oxidation varies from 5.5 to 1.9 (Table 4). It corresponds to an "activation energy" of 5.0 kJ/mol, which is in a good agreement with the difference of the zero point energies of C-H and C-D bonds. Thus, similar to MMO with its KIE value of 5 [25], the rate determining step of CI-L oxidation by a-oxygen also involves the cleavage of C-H bond.

Table 4 NMR analysis of product methanol isotops and KIE values of CH2D2 methane oxidation with a-oxygen (8 is a chemical shift, IH and ID are signal intensities) CHD2OD

CH2DOD

Run

T

KIE = 2IH/ID

No

(K)

8 (ppm)

IH (rel.unit)

8 (ppm)

ID (rel.unit)

1 2 3

223 293 373

3.30 3.30 3.30

63 68 56

3.32 3.32 3.32

23 42 58

5.5 3.2 1.9

3.3. Active state of iron The distribution of iron and other metals in zeolites has been studied by many authors. In general, Fe can occupy three positions in ZSM-5 matrix [26]" (1) as isolated ions in the tetrahedral lattice positions; (2) as isolated ions or small complexes outside the lattice but inside

500 the intracrystalline micropore space; (3) as clusters and finely dispersed oxide particles on the outer surface of zeolite crystals. Which of these positions occupy a-sites is an important question for our consideration. The first position can be safely excluded since a high temperature calcination, causing the removal of Fe atoms from the lattice, remarkably increases the a-site concentration [27]. Besides, a-sites can be prepared via the impregnation of a ready zeolite matrix [28], when the probability for Fe atoms to incorporate into the lattice is very low. a-Sites do not occupy also the 3rd type position: deactivation of the outer zeolite surface by its coveting with an inert SiO2 layer affects neither catalytic activity no a-site concentration [29]. Thus, we may deduce that the active iron occupies the second type position in ZSM-5 matrix and is either isolated Fe ions or small complexes inside the mieropore zeolite space. According to the following evidences, a-sites are most probably di-iron complexes similar to the di-iron active sites of methane monooxygenase: a) The ratio between the number of a-sites and that of Fe atoms in the FeZSM-5 zeolites sometimes achieves but never exceeds 2. b) The active iron is invisible in the ESR spectra [12], which is consistent with the formation of dinuclear complexes. c) Below 110 K, the magnetic susceptibility of FeZSM-5 decreases with the decreasing temperature, which is typical for the antiferromagneticaUy bound iron complexes [30]. d) Experimentally observed a-oxygen features are well interpreted within a dinuclear quantum-chemical model of a-sites [31 ]. Recent results obtained with Moessbauer spectroscopy [ 13] provided additional arguments in favor of MMO similarity. According to [32] iron atoms in MMO, depending on conditions, exist m both oxidized and reduced states. Two different quadrupol doublets correspond to each of these states with main parameters given in Table 5. The Moessbauer spectra of FeZSM-5 also reveal two states of iron, which are represented by two quadrupol doublets ffig 4). The narrow doublet parameters correspond to the Fe(III) state, the broad doublet parameters correspond to the Fe0I) state. One can see an excellent agreement between the spectral characteristics of both reduced and oxid~ed iron complexes in MMO and in the ZSM-5 matrix. An agreement of this quality is a rather difficult situation to achieve. Beside FeZSM-5, Table 5 includes also Moessbauer data for the FeY zeolite and for some model compounds specially I

L i

Fe-'+ ,.."

9

"."

,

I

I

/"

9

.~.~,-~

Figure 4. Moessbauer spectrum of 57Fe enriched FeZSM-5 recorded at 90 K Solid line is the best fit obtained with the two quadrupole doublets. The sample was calcined at 820 K, excess of inactive iron was removed by oxalate extraction.

i l

.

-2

0

i

4 .................

Velocity (mm/s) prepared to simulate dinuclear iron centers in biology. One can see that their spectral parameters do not correlate so well with MMO, as FeZSM-5 parameters do

501 Table 5 Moessbauer spectra parameters of Fe-containing systems Sample .

.

.

.

.

.

.

Fe state . . .

Isomer shif~ (mm/s)

Quadrupol splitting AE(mm/s)

Reference

MMO

Fe(II)-Fe~) Fe(III)-Fe(III)

1.30 0.50

3.014 1.05

32

FeZSM-5

F~ Fe(I~

1.34 0.47

3.09 0.99

this work

[Fe2(OH)(OAc)2(Mes TACN)] CIO4

Fe(I~Fe(ID

1.16

2.83

33

[Fe2(BPMP)(OPr)2](BPh4)

Fe(II)-Fe(II)

1.20

2.72

34

[Fe20(OAc)2I-IB(pz)s)2]

Fe(RI)-Fe(III)

0.52

1.60

35

[Fe20(O2CH)4(BiPhMe)2]

Fe(III)-Fe(IlI)

0.54

1.81

36

FeY

Fe(II) Fe(III)

1.17 0.34

2.14 1.22

37

4. CONCLUSION Results discussed above show in several lines a distinct biomimetic-type activity of iron complexes stabilized in the ZSM-5 matrix. The most important feature is their unique ability to coordinate a very reactive c~-oxygen form which is similar to the active oxygen species of M/dO. At room temperature or-oxygen provides various oxidation reactions including selective hydroxylation of methane to methanol. Like in biological oxidation, the rate determining step of this reaction involves the cleavage of C-H bond. These data allow to consider FeZSM-5 zeolite as a new successful monooxygenase model. It is worth noticing the important role of the zeolite structure for constructing such an inorganic model, though in this paper we did not have room to discuss this topic. A remarkable advantage of this robust model is an opportunity to study the monooxygenase-like oxygen as well as the active state of iron under easily controlled and reproducible conditions. It opens new possibilities in getting more knowledge on the mechanism of enzymatic oxidation. REFERENCES

1. 2. 3. 4. 5. 6. 7.

A.E. Shilov, in Activation and Functionalization of Alkanes (ed. C.L. Hill), Wiley, New York, 1989, p. 13. G.I. Panov, A.S. Kharitonov and V.I. Sobolev, Appl. Catal., 98 (1993) 1. D.B. Tagiyev and Kh.M. Minachev, Stud. Surf. Sci. Catal., 28 (1986) 981. G.I. Panov, V.I. Sobolev and A.S. Kharitonov, J. Mol. Catal., 61 (1990) 85. G. Belussi, A. Carati, M. Clerici, and R. Millini, J. Catal., 133 (1992) 220. M. Ivamoto, K. Matsukami and S. Kagawa, J. Phys. Chem., 87 (1983) 903. E. Suzuki, K. Nakashiro and Y. Ono, Chem. Lett., (1988) 953.

502 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. 34. 35. 36. 37.

M. Gubelmann and Ph. Tirel, Preparation of Phenol by Direct Hydroxylation of Benzene, Fr. Patent No.2 630 735 (1989). A. Kharitonov, L. Vostrikova, K. Ione and G. Panov, Phenol Preparation, Rus.Patent No. 1805127 (1989). A.S. Kharitonov, G.A. Sheveleva, G.I. Panov, Ye.A. Paukshtis and V.N.Romannikov, Appl. Catal. A, 98 (1993) 33. G.I. Panov, G.A. Sheveleva, A.S. Kharitonov, V.N. Romannikov and L.A.Vostrikova, Appl. Catal. A, 82 (1992) 31. V.I. Sobolev, G.I. Panov, A.S. Kharitonov, V.N. Romannikov, A.M. Volodin and K.G. Ione,. J. Catal., 139 (1993) 435. G.I. Panov, V.I. Sobolev, K.A. Dubkov, V.N. Parmon, N. Ovanesyan, A.Ye. Shilov and A.A. Shteinman, J. Amer. Chem. Soc., submitted for publication. V.I. Sobolev, O.N. Kovalenko, A.S. Kharitonov, Yu.D. Pankrat'ev and G.I. Panov Mendeleev Commun., 1 (1991) 29. V. Zholobenko, L. Kustov and V. Kazansky, in R.Ballmoos and M. Treacy (Eds.), Proc. 9th Intern. Zeolite Conf., Butterworth-Heinemann, Boston 1992, vol. 2, p.299. J. Leglise, J.O. Petunnchi and W.K. Hall, J.Catal., 86 (1984) 392. G.K. Boreskov, Catalysis: Science and Technology, 3 (1982) 40. V.Sobolev, A.Kharitonov, E.Paukshtis and G. Panov, J. Mol. Catal., 84 (1993) 117. Y. Moro--oka, Stud. Surf. Sci. Catal., 54 (1990) 53. B.V. Romanovsky, Micromol. Symp., 80 (1994) 185. C.A. Tolman, J.D. Druliner, M.J. Nippa and N. Herron, ref. 2, Chapter 10. R.E. Parton, I.F.J. Vankelecom, M.J.A. Casselman, C.P. Bezoukhanova, J.B. Uytterhoeven and P.A. Jacobs, Nature, 370 (1994) 541. H. Dalton, Catal. Today, 13 (1992) 455. J.T. Groves and G.A. McClusky, Biochim. Biophys. Res. Commun., 81 (1978) 154. A.M. Khenkin and A.E. Shilov, New J. Chem., 13 (1989) 659. P. Kamasamy and R. Cumar, Catal. Today, 9 (1991) 328. V.I. Sobolev, K.A. Dubkov, Ye.A. Paukshtis, L.V. Pirutko, M.A. Rodkin, A.S. Kharitonov and G.I. Panov, Appl. Catal., submitted for publication. L. Pirutko, A. Kharitonov and V. Buchtiyarov, Kinet. Katal., to be published. L.Pirutko, O.Parenago, E.Lunina, A.Kharitonov, L.Okkel and G.Panov, React. Kinet. Katal. Lett., 52 (1994) 275. E. Smimov, L. Makarshin, V. Sobolev, V. Parmon and G. Panov, in preparation. M. Filatov, A. Pelmenschikov and G. Zhidomirov, J. Mol. Catal., 80 (1993) 243. J.G. Dewitt, J.G. Bentsen, A.C. Rosenzweig, B. Hedman, J. Green, S. Pilkington, G.C. Papaefthymion, H. Dalton, K.O. Hodgson and S.J. Lippard, J. Am. Chem. Soc., 113 (1991)9219. J. Hartman, R. Rardin, P. Chaudhuri, K. Pohl, K. Wieghardt, B. Nuber, J. Weiss G.Papaefthymion, R.Frankel and S.Lippard, J. Am. Chem. Soc., 109 (1987) 7387. A.S. Borovik and L.Jr. Que, J. Am. Chem. Soc. 110 (1988) 2345. W.H. Armstrong and S.J. Lippard, J. Am. Chem. Soc., 106 (1984) 4632. W. Tolman, S. Lin, J. Bentsen and S. Lippard, J. Am. Chem. Soc., 113 (1991) 152. Aparicio. J.A. Dumesic, S.M. Fang, M.A. Long, M.A. Ulla, W.S. Millman and W.K. Hall, J. Catal., 104 (1987) 381.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) I I th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

503

A Molecular Approach to Synergy Generation in Co-Mo Binary Sulfide Catalysts for Hydrodesulfurization Yasuaki Okamoto and Hiromoto Katsuyama Department of Chemical Engineering, Faculty of Engineering Science, Osaka University, Toyonaka, Osaka 560, Japan The mechanism of synergy generation between Co and Mo sulfides for Hydrodesulfurization was studied for highly dispersed Co-Mo binary sulfide clusters encaged in a NaY zeolite. It was shown by EXAFS, XRD, and HREM that highly dispersed Mo sulfides were prepared in the zeolite cages by sulfiding adsorbed Mo(CO)6 . Highly dispersed Co-Mo binary mixed sulfides were synthesized in the zeolite cavities by using Mo(CO)6 and Co(CO)3NO as precursors. It is concluded that thiophene hydrodesulfurization and butadiene hydrogenation take place on the Co sites of the Co-Mo binary sulfide clusters and that chemicai bondings of Co-S-Mo are required for the synergy generation. Coordinatively unsaturated sites were detected by NO adsorption only on the Co sites of the Co-Mo mixed sulfides. It is proposed that the catalytic activities of the Co sites are promoted by electronic modifications induced by the Mo sulfides. 1. INTRODUCTION Hydrotreatmgs of petroleum feedstocks have recently become more and more crucial not only for protecting environments but also for efficient utilization of limited natural resources. Developments of highly active and selective hydrotreating catalysts, in particular hydrodesulfurization (HDS) catalysts, are one of the most urgent problems in petroleum industries. Sulfided Co-Mo or Ni-Mo based catalysts have been used in industry for HDS reactions. It is well known that strong catalytic synergies generate between Co(Ni) and Mo sulfides [ 1-4]. The mechanism of synergy generation in Co-Mo binary sulfide catalysts has been extensively studied for HDS but still remains controversial at present [1-4]. Two main synergy models are under discussion now: a CoMoS model [1,5] and a contact synergy model [3]. In the CoMoS model, the formation of atomically dispersed Co sulfide species anchored on the edge sites of MoS 2 is claimed to be the origin of the synergy generation. The local structures of Co and Mo atoms for supported Co-Mo catalysts have been proposed on the basis of EXAFS results [6-9]. In the contact synergy model, the promotional effects of Co are explained in terms of a remote control mechanism. It is believed that the HDS activity of Mo sulfides is enhanced by spillover hydrogen originally generated on highly dispersed Co sulfides in contact with or in the proximity of the Mo sulfides. Preparations of supported Co-Mo binary sulfide clusters well defined in the structure on a molecular level are considered to be a promising approach to clarify the mechanism of synergy generation. Supported and unsupported metal sulfide clusters have been examined for the purpose up to now [ 10,11 ]. In the present study, we synthesized in zeolite cavities Co-Mo binary sulfide clusters by using Co and Mo carbonyls and characterized the clusters by extended X-ray absorption fine structure (EXAFS), X-ray photoelectron spectroscopy (XPS), Fourier transform infrared spectroscopy (F'FIR), and high resolution electron microscopy (HREM). The mechanism of catalytic synergy generation in HDS is discussed.

504

2. EXPERIMENTAL 2.1. Catalyst Preparation A NaY zeolite (A1/Si atomic ratio; 0.41) was supplied by Shokubai Kasei Kogyo Ltd. After an evacuation at 673 K for 1 h (lx 10 .3 Pa), the zeolite powder was exposed to a vapor of Mo(CO) 6 or Co(CO)3NO at room temperature, followed by an evacuation at room temperature for 10 rain to remove physisorbed metal carbonyl molecules on the external surface of the zeolite. Mo(CO)6/NaY or Co(CO)~NO/NaY was sulfided in a stream of an atmospheric pressure of 10% H2S/H 2 (0.2 dm 3 rain1). The sulfidation temperature was increased from room temperature to 373 K at a rate of 2 K rain ~ and kept at the temperature for 1 h. Subsequently, the temperature was increased up to 673 K at a rate of 5 K min ~ and kept at 673 K for 1.5 h. After the sulfidation, the sample was cooled in the H2S/H 2 stream to room temperature. The Mo and Co sulfide catalysts thus prepared are denoted MoSx/NaY and CoSx/NaY, respectively. Mo sulfide catalysts, MoS2/NaY , were also prepared by a conventional impregnation method by using ammonium heptamolybdate, for comparison. A series of CoSx-MoSx/NaY catalysts was synthesized by intoducing Co(CO)~NO into MoSx/NaY evacuated at 673 K for 1 h, followed by second programmed sulfidation procedures. MoSx-CoSx/NaY catalysts were prepared in the reversed order of the metal sulfide accommodations into the zeolite cavities. When Co2(CO)8 was used as the Co precursor, MoSx/NaY was impregnated with Co2(CO)8 dispersed in n-hexane, followed by evacuation at room temperature to remove the solvent. Co2(CO)s/MoSx/NaY was subsequently sulfided at 673 K to give CoSx/MoSx/NaY. The catalyst composition was determined by AAS and ICP.

2.2. Reaction Procedures The sulfided catalysts were evacuated for 1 h before catalytic reactions. The reactions were carried out under mild conditions by using a circulation system (0.2 dm 3) made of glass. The HDS of thiophene was conducted at 623 K and an initial pressure of 20 kPa (H2/CnH4S = 36). The thiophene pressure was kept constant (0.54 kPa) during the reaction by holding a small amount of liquid thiophene kept at 273 K in the bottom of a U-tube in the reaction system. The products were analyzed by gas chromatography. The HDS activity was calculated from the amount of H2S produced during the reaction. The hydrogenation (HYD) of butadiene was subsequently conducted at 473 K over the catalyst, that had been used for the HDS of thiophene for 1 h, after evacuation at 673 K for 1 h. The initial pressure was 14 kPa (H2/C4H6 = 2). The HYD products were butenes with a small amount of butane. The catalytic activity was calculated on the basis of the total amount of the reaction products.

2.3. Catalyst Characterization The Mo K-edge EXAFS spectra for the catalysts and reference compounds (1k/loS2 and ]'qa2MoO4) were measured on the BL-10B instruments of the Photon Factory at the National Laboratory for High Energy Physics by using a synchrotron radiation. The EXAFS spectra were obtained at room temperature without exposing the sample to air by using an in situ EXAFS cell with Kapton windows [12]. Data analysis was carried out assuming a plane wave approximation. The XPS spectra of the freshly sulfided Co-Mo/NaY catalysts were measured on an XPS-7000 photoelectron spectrometer (Rigaku, A1 anode; 1486.6 eV). The sample mounted on a holder was transferred from a glove bag into a pretreatment chamber attached to the spectrometer as possible as carefully not to be contacted with air. The binding energies (BE) were referenced to the Si2p band at 103.0 eV for the NaY zeolite, which had been determined by the Cls reference level at 285.0 eV due to adventitious carbon. The FTIR spectra of the NO molecules adsorbed on the catalyst were measured on a JEOL 3IR-100 spectrophotometer in a diffuse reflectance mode. After a back ground spectrum was measured in situ for a freshly sulfided catalyst, NO was introduced to the catalyst as 10 % NO/He pulses (5.1 cm3). The FTIR spectra were recorded after an introduction of 5 pulses.

505

3. RESULTS and DISCUSSION 3.1. Mo and Co Sulfide Catalysts The catalytic activities of MoSx/NaY, CoSx/NaY, and MoS2/NaY are shown as a function of the metal atoms per supercage (SC) for the HDS of thiophene and for the HYD of butadiene in Figs. 1 and 2, respectively. MoSx/NaY catalysts exhibited higher catalytic activities by a factor of 3-4 than the corresponding impregnation catalysts, MoSJNaY, as previously reported [13]. The HDS activity of MoSx/NaY increased linearly with the Mo content up to 2Mo atoms/SC, suggesting a formation of uniform Mo sulfide species in this concentration range. The activity leveled off, however, at a higher Mo content, indicating an agglomeration of the Mo sulfide species. Similar loading effects were observed for the HYD of butadiene over MoSx/NaY and MoS 2/NAY as shown in Fig.2. It is noteworthy that CoSx/NaY showed a considerably high HDS activity, being comparable with that of MoSx/NaY. In contrast to relatively low HDS activities of the Co sulfide catalysts supported on A1203, the Co sulfide species supported on activated carbon have been reported to show even higher HDS activities than Mo sulfide catalysts [ 14,15]. This is attributed to an extremely high dispersion of the Co sulfide species on activated carbon. The high HDS activity of CoSx/NaY suggests a high dispersion of the Co sulfide species. With the HYD of butadiene, CoSx/NaY showed a much lower activity than MoSx/NaY. The local structure and dispersion of the Mo sulfide species were examined by EXAFS techniques. The Fourier transforms of the Mo K-edge EXAFS for MoSx/NaY showed two peaks ascribed to Mo-S and Mo-Mo bondings (vide infra). The coordination number of the Mo-Mo bondings for MoSx/NaY was calculated to be close to unity, as summarized in Table 1, indicating the formation of highly dispersed Mo sulfide species, possibly Mo dimer species, in the supercage of the zeolite. High NO adsorption capacities of MoSx/NaY [ 13] are consistent with the high dispersion of the Mo sulfide species. Mo sulfide catalysts were also synthesized by sulfiding Mo oxide dimer species encaged in the NaY zeolite, (MoO3)2/NaY, which had been prepared by a mild oxidation of Mo(CO)6/NaY by using molecular oxygen [16]. The structural parameters derived from EXAFS analysis were identical with those for MoSx/NaY in Table 1.

6ot

80 120 ,.o0 100 x

=-->.- 8O z

o

60

40

"'o-

._Z. 40

20

..~>_ f,O

<

,'"" "~0

2o

A,

( 0

1

2

.

Metal-atoms [ supercage

;

6

Figure 1. Catalytic activities of MoSx/NaY(Q), CoSx/NaY(O), and MoS2/NaY (&) for the HDS of thiophene as a function of the metal content (metal atoms/SC),

Metal-atoms / supercage

Figure 2. Catalytic activities of MoSx/NaY (O), CoSx/NaY (O), and MoSz/NaY (&) for the HYD of butadiene as a function of the metal content (metal atoms/SC).

506 Table l Structural Parametersa)as Derived from the Mo K-Edge EXAFS for Mo and Co-Mo Sulfide Catalysts Encaged in a NaY Zeolite catalyst

loadings/SC b) Mo

MoSx/NaY c)

2.1

MoSx/NaY d)

2.1

MoS2/NaY e)

0.8

CoSx-MoSx/NaY

2.1

bondings

R/A

CN

E 0/eV

Mo-S Mo-Mo Mo-S Mo-Mo Mo-O Mo-S Mo-Mo Mo-S Mo-Mo Mo-Co 0

2.40 3.15 2.41 3.18 1.78 2.42 3.16 2.41 3.18 2.83

4.7 1.1 4.5 1.2 0.8 1.1 1.2 5.0 0.7 -

-2.3 -2.0 -2.1 2.3 12.7 -5.0 -3.8 -5.8 0.2 _

A

or2~A 2

Co

2.1

a) R" bond distance, CN; coordination number, E0; inner potential factor b) Metal atoms/supercage c) Sulfided at 673 K for 1.5 h, d) for 20 h, e) for 5 h f) Bond distance was estimated by using theoretical parameters

0.0048 0.0089 0.0034 0.0075 0.0017 0.0016 0.0110 0.0020 0.0079 _

cr " Debye-Waller-like

It was shown by EXAFS that the structure and dispersion of the Mo sulfide species in MoSx/NaY were unaltered by a prolonged sulfidation of 20-h at 673 K in a stream of I~S/t-~. This fact indicates that highly dispersed Mo sulfide species are thermally stable. On the other hand, with MoSx/AI,O3 prepared by using Mo(CO)6 [12], a considerable agglomeration of highly dispersed Mo sulfide species was observed at a shorter treatment time. It is considered that highly dispersed Mo sulfide clusters are thermally stabilized in zeolite cavities. In contrast to MoSx/NaY, the Fourier transform for MoSJNaY clearly showed Mo-O bondings as well as Mo-S and Mo-Mo bondings as summarized in Table 1. It is evident that Mo oxide species in calcined MoO3/NaY are only partially sulfided. XPS results corroborated the EXAFS results. The incomplete stdfidation of the Mo species in MoS z/NaY may explain in part the relatively low HDS and HYD activities of the catalysts in Figs. 1 and 2. The location or distribution of the Mo sulfide species, that is, inside or outside the zeolite cavities, was examined by HREM, XRD [17], and pore volume measurements by using benzene as adsorbate [ 18]. HREM observations for MoSx/NaY possessing 2Mo/SC obviously demonstrated that no Mo sulfide species were formed on the outside of the zeolite and that the framework structure of the zeolite was not destroyed at all on the accommodation of Mo sulfide species. The XRD and pore volume measurements confirmed the HREM observations. It is concluded that highly dispersed intrazeolite Mo sulfide species are produced by using Mo(CO) 6.

3.2. Preparation and Catalysis of Co-Mo/NaY The HDS activity of CoSx-MoSx/NaY is shown in Fig.3 as a function of the Co/Mo atomic ratio. The Mo content in the catalyst was kept constant at the saturation value for a Mo(CO)6 adsorption in NaY, 2.1Mo/SC. The HDS activity increased as the Co/Mo ratio increased up to about unity, followed by an activity decrease at a further addition of Co. The simple sum of the activities of the corresponding CoSx/NaY and MoSx/NaY is presented in Fig.3 for comparison. It is evident that the activity of the Co-Mo binary sulfide catalyst is considerably higher than the simple sum of the activities of the composite sulfides, indicating

507 the generation of a catalytic synergy for the HDS between Co and Mo sulfides in the zeolite system. Taking into consideration the preparation procedures, the Mo content of 2.1Mo/SC, and the Co/Mo atomic ratio of ca. unity at the maximum HDS activity, highly dispersed Co-Mo binary sulfide clusters, possibly Co2Mo2Sx, in the supercage of the NaY zeolite are suggested for catalytically active species. The HDS activity of the CoSx-MoSx/NaY was not changed even after a 20-h treatment at 673 K in a stream of H2S/H 2 (Fig.3), demonstrating a high thermal stability of the active species. MoSx-CoSx/NaY catalysts, which were prepared by introducing Mo(CO)6 into CoSx/NaY (1.1Co/SC), showed the identical HDS activities with those of CoSx-MoSx/NaY at the same compositions, as illustrated in Fig.4. Figure-4 suggests that the dispersions of Mo and Co sulfides are not mutually affected by the presence of the other sulfide species or that the formation of catalytically active species, e.g. Co-Mo mixed sulfide species, is independent of the accommodation order. As shown below, FTIR of NO adsorption, EXAFS, and XPS results supported the latter possibility. The activity of CoSx-MoSx/NaY (2.1Mo/SC) is shown in Fig.5 for the HYD of butadiene as a function of the Co/Mo atomic ratio. The HYD activity decreased slightly on the addition of Co up to Co/Mo - ca. 1, followed by a steep decrease at a further incorporation of Co. The HYD/HDS activity ratio decreased with increasing Co content and reached the ratio for CoSx/NaY at the Co/Mo atomic ratio of the maximum HDS activity (Fig.3). The product selectivity in the HYD of butadiene shifted from t-2-butene rich distribution to 1-butene rich one on the addition of Co, as presented in Fig.6. It is worthy of noting that at the Co/Mo ratio of the maximum HDS activity, the butene distribution is close to that for CoSx/NaY. It should be noted, however, that these product distributions are not the initial distributions of the HYD over the catalyst but the distributions modified by successive isomerization reactions. It was found that MoSx/NaY showed high isomerization activities of butenes even in the

200

"- 160

x 120 ~o

g

g 80

&128

&

>,

-

0

f

>,

'

0.5

1 1.5 Co/Mo atomic ratio

:~

Figure 3. HDS activities of CoSx-MoSx/NaY (2.1Mo/SC) (O) as a function of the Co/Mo atomic ratio. F--l;after a 20-h treatment in HzS /I-I2 at 673 K. The sum activities of CoSx/NaY and MoSx/NaY are shown for comparison (O).

40

0

'

1 2 Mo atoms/SC

'

Figure 4. HDS activities of MoSx-CoSx/ NaY (1.1Co/SC) ( O ) and CoSx-MoSx/NaY (1.1Co/SC) ([--3) are compared as a function of Mo atoms/SC.

508 presence of butadiene and that CoSx/NaY exhibited high isomerization activities in the absence of butadiene, whereas very small activities in the presence of butadiene. CoSx/NaY and CoSx-MoSx/NaY (Co/Mo = c a . 1) provided almost the original 1-butene rich selectivity in the HYD of butadiene, demonstrating that the isomerization sites and, consequently, the hydrogenation sites on the Mo sulfides are completely masked by the presence of the Co sulfide species. It is well established [19] that the hydrogenation needs active sites possessing a higher coordinative unsaturation (triply cus) than the isomenzation does (doubly cus). Summarizing the results in Figs.5 and 6, it is concluded that the HYD of butadiene takes place only on the Co sites of the Co-Mo binary sulfide catalyst at Co/Mo = c a . 1 and that the Co sites play important roles in the HDS reaction, too. As illustrated in Fig.2, the HYD activity of the Co sulfide species in CoSx/NaY is considerably lower than that of the Mo sulfide species. The relatively high HYD activity of CoSx-MoSx/NaY apparently indicates that the HYD activity of the Co sulfide is greatly enhanced by the presence of the Mo sulfide species. 3.3. Characterization of Co-Mo/NaY

The major contributions of the Co sites in CoSx-MoSx/NaY to the HYD and HDS reactions were corroborated by a FTIR study of NO adsorption. Figure 7 shows the IR spectra of NO adsorbed on CoSx/NaY (2.1Co/SC) and CoSx-MoSx/NaY (2.1Co + 2.1Mo/SC). Nitric oxide molecules adsorb on Co or Mo sulfides forming dinitrosyl complexes [20,21 ]. With CoSx/NaY, two peaks were observed at 1890 and 1811 cm 1 with a weak shoulder band at 1865 cm", suggesting the presence of two kinds of NO adsorption sites. Topsoe and Topsoe [21 ] reported two NO adsorption bands at 1842 and 1778 cm" for Co/AI203 (2 % Co) sulfided at 653 K. The differences in the wavenumbers between the zeolite and A1203 systems may be due to the differences in the structure and dispersion of the Co sulfide species. MoSx/NaY showed two absorption bands at 1795 and 1660 cm" (not shown) in conformity with those reported for

1 x

60

100

..

80

-r- .~. .--

05 a

6( 4

.>'~

oo

40

-

(

2q o

[

,

0.5 9 1 C o / M o a t o m i c ratio

o oo

Figure 5. HYD activity (O) and HYD/HDS activity ratio (IS]) of CoSx-MoSx/NaY (2.1 Mo/SC) as a function of the Co/Mo atomic ratio.

,

0.5

,

,

1

....

,~,,

,

oo

C o / M o a t o m i c ratio

Figure 6. Butene distribution of the HYD of butadiene over CoSx-MoSx/NaY (2.1Mo/ SC) as a function of the Co/Mo atomic ratio. O ; t-2-butene, ( ~ cis-2-butene, O ; 1-butene.

509

sulfided Mo/A1203 (1780 and 1685 cm -~) [20,21]. CoSx-MoSx/NaY exhibited doublet bands at 1867 and 1807 cm ~, accompanying a weak shoulder peak at c a . 1880 cm ~ . These signals are apparently assigned to those of NO molecules adsorbed on Co sulfides. No peaks ascribable to the NO adsorption on Mo sulfide sites were detected at all. What is important in Fig.7 is that in CoSx-MoSx/NaY, coordinative unsaturation sites are present only on the Co sites in spite of the coexistence of the same amount of Mo sulfide species in the zeolite cavities. These results clearly support that the Co sites in CoSx-MoSx/NaY play major roles in the HYD and HDS reactions. The IR spectra in Fig.7 indicate the preferential adsorption of NO on the Co sites. It may be conjectured that the Mo sulfide species are physically covered by the Co sulfide species or that Co-Mo mixed sulfide species are formed and the chemical natures of the Co and Mo sulfides are mutually modified. The Mo K-edge EXAFS spectra were measured to examine the formation of mixed sulfide species between Co and Mo sulfides. The Fourier transforms are presented in Fig.8 for MoSx/NaY and CoSx-MoSx/NaY. The structural parameters derived from EXAFS analysis are summarized in Table 1. The structure and dispersion of the Mo sulfides in MoSx/NaY are discussed above. With the Co-Mo binary sulfide catalyst, the Mo-Co bondings are clearly observed at 0.283 nm in addition to the Mo-S and Mo-Mo bondings. The Mo-Co distance is close to that reported by Bouwens et al. [7] for a CoMoS phase supported on activated carbon. Detailed analysis of the EXAFS results for CoSxMoSx/NaY will be presented elsewhere. It is concluded that the Co-Mo mixed sulfides possessing Co-S-Mo chemical bondings are formed in CoSx-MoSx/NaY.

o (.3

CoSx-MoSx/NaY

6

0

CoSx-MoSx/NaY

f

x/NaY 21oo

1

1

19oo

1"/oo

v

15oo

Wavenumber (cm "I)

Figure 7. FI'IR spectra of NO adsorption on CoSx/NaY (2.1Co/SC) and CoSx-MoSx/NaY (2.1Mo + 2.1Co/SC).

0

1

2

x.,," - - - - ~

3

4

5

Distance / A Figure 8. Fourier transforms of ~-weighted EXAFS modulations of the Mo K-edge for MoSx/NaY and CoSx-MoSx/NaY.

510 Table 2 XPS Results for Co-Mo Binary Sulfide Catalysts catalyst a)

binding energy b/ev Co2P3/2

CoSx-MoSx/NaY(2.2)

$2P3/2

Mo3d5i 2

Co2p/Si2p

161.9

228.9

1.21

778.4

161.9

228.8

778.4

162.1 d)

779.1

CoSx/MoSx/NaY(1.2)

intensity ratio Mo3d/Si2p 0.77

2.62

0.85

o

Co9S8 c)

.

.

.

.

.

.

.

.

.

a) Mo content; 2.1-2.2 atoms/SC. The number in parentheses denotes Co atolns/SC. b) Referenced to the Si2p level at 103.0 eV for the NaY zeolite. c) Alstrup et al., J. Catal., 77 (1982) 397. d) $2Pl/2 + S2P3/2 The formation of Co-Mo mixed sulfide species was also supported by the Co2p XP spectra. The XPS results are summarized in Table 2. The binding energy (BE) of the Co2p band for CoSx-MoSx/NaY is higher by 0.7 eV than that of Co9S8 [22], indicating the chemical reactions between Co and Mo sulfide species [22]. It is concluded that Co-Mo mixed sulfide clusters, possibly CozMozSx, are formed in the zeolite cavities by using metal carbonyl techniques and that the formation of the mixed sulfide clusters preferentially makes the Co species coordinatively unsaturated and, thereby, actively involved in the HDS and HYD reactions, while it renders the Mo sites coordinatively saturated and less active. The formations of mixed sulfides might be due to high reactivities of the preexisting Mo sulfide clusters with the Co carbonyl [23] and/or with highly dispersed Co sulfide clusters and v i c e v e r s a . The exact composition and structure of the catalytically active Co-Mo binary sulfide clusters are under investigation at present. 3.4. Mechanism of Synergy Generation As briefly reviewed in the introduction, two major models, proposed to explain catalytic synergies between Co and Mo, are the CoMoS model [ 1,5] and the contact synergy model [3]. The Co-Mo mixed sulfide clusters encaged in the NaY zeolite are considered to represent the CoMoS model where Co-S-Mo chemical bondings are formed. On the other hand, if we could prepare Co sulfides at the outside of the zeolite with Mo sulfides at the inside of the zeolite, this catalyst system would represent the contact synergy model, in which spillover hydrogen is expected to play major roles in HDS reactions. Co2(CO)8 was employed for the purpose instead of Co(CO)~NO, since the rate of the migration of Coz(CO)sinto the supercage of the zeolite was much slower than the monomeric species (1% Co adsorption after 16 h vs. 8 % Co within 1 min.). The Co-Mo binary catalyst prepared by using Coz(CO) 8 is denoted CoSx/MoSx/NaY. The XPS results are presented for CoSx/MoSx/NaY in Table 2. The BE of the Co2p band is close to that for Co9Ss, indicating no formation of Co-Mo mixed sulfide species, at least in the catalyst surface. In addition, the Co2p/Si2p XPS intensity ratio was much higher than that for CoSx-MoSx/NaY in spite of a lower Co content. These results demonstrate a preferential segregation of Co sulfides at the outside of MoSx/NaY. The HDS activities of the Co-Mo binary sulfide catalysts are presented in Fig.9 and compared with the sum activities of the corresponding composite catalysts. CoSx-MoSx/NaY showed a higher HDS activity than the sum activity as shown in Fig.3, while CoSx/MoSx/NaY exhibited almost the same activity with the sum activity. Figure 9 suggests no strong synergy generation in CoSx/MoSx/NaY. It is, therefore, concluded that under the present reaction conditions, the formation of Co-Mo mixed sulfides are necessary for the synergy generation.

511

Figure 9. Activities of CoSx-MoSx/NaY and CoSx/MoSx/NaY for the HDS of thiophene. The corresponding sum activities of the composite catalysts are also shown for comparison.

The present results support the CoMoS model for the mechanism of synergy generation. According to Riaz et al. [24], a bimetallic cluster Cp'2Mo2Co2S3(CO)4(Cp': methylpentadiene) (1) desulfurizes thiols and thiophene to give a cubane cluster CP'2Mo2Co2S4(CO)2 (2). This reaction is not catalytic but just stoichiometric, since the cubane cluster (2) is not converted to (1) at the low reaction temperature, 373-423 K, adopted by them. The initiation process of the desulfurization reaction is shown to be a nucleophilic attack of the sulfur compounds on the Co sites of the bimetallic cluster (1) but not on the Mo sites. These results by Riaz et al. are consistent with the present findings that the Co sites of the Co-Mo mixed sulfide clusters in the zeolite play important roles in the HDS. An analogous nucleophilic attack or adsorption of thiophene on the Co sites is also considered to be of great importance in the HDS of thiophene over the Co-Mo mixed sulfide catalysts. The preferential formation of coordinatively unsaturated Co sites, which is induced by chemical bondings between the Co and Mo sulfides, apparently promotes the initiation process of HDS. The increased BE of the Co2p band for the Co-Mo mixed sulfides suggests an electron transfer from the Co to Mo atoms, in conformity with the quantum chemical calculations by Hams and Chianelli [25] for a mixed Co-Mo model compound. The increased charge on the Co sites is expected to facilitate the nucleophilic attack of thiophene, thereby, further promoting the HDS activity of the Co sulfide species. We are inclined to propose that electronic modifications of the Co sites by the Mo sulfides, that is, the preferential formation of coordinative unsaturation and the increase in the positive charge on the Co sites, are one of the most important factors for the synergy generation. In addition, increases in the Co sulfide dispersion are brought about by chemical bondings with highly dispersed Mo sulfide species. The latter effects are considered to be particularly important for alumina supported Co-Mo sulfide catalysts. In the zeolite systems, individual Co and Mo sulfide species are highly dispersed and, therefore, the electronic effects play the major roles in the synergy generation.

512 4. CONCLUSIONS It was demonstrated in this study that highly dispersed Co-Mo binary sulfide clusters were synthesized in the cavities of a NaY zeolite by using the metal carbonyls as precursors. The catalytically active species are proposed to be Co-Mo mixed sulfide clusters, possibly C02Mo2Sx. The Co sites were found to constitute the active sites for the thiophene hydrodesulfurization and butadiene hydrogenation. It is considered that chemical bondings of Co-S-Mo are necessary for the synergy generation. The catalytic activities of the Co sites are suggested to be promoted by electronic effects of the chemically bonding Mo sulfides. ACKNOWLEDGMENTS

We express our thanks to Prof. M.Yamada (Tohoku University) for the PTIR measurements and Mr. M.Matsuo (Rigaku Ltd.) for the XPS measurements. We are also grateful to Prof. M.Nomura and staff of the Photon Factory, National Laboratory for High Energy Physics, for assistance in measuring EXAFS spectra (Proposals: 89138 and 93G 163). REFERENCES

1. H.Topsoe and B.S.Clausen, Catal. Rev. -Sci. Eng., 26 (1984) 395. 2. R.Prins, V.H.J. de Beer, and G.A.Somorjai, Catal. Rev. -Sci. Eng., 31 (1989) 1. 3. B.Delmon, Catal. Lett., 22 (1993) 1. 4. R.R.Chianelli, M.Daage, and M.J.Ledoux, Adv. Catal., 40 (1994) 177. 5. H.Topsoe, B.S.Clausen, N.Topsoe, and E.Pedersen, Ind. Eng. Chem. Fund., 25 (1986) 25. 6. S.M.A.M.Bouwens, R.Prins, V.H.J. de Beer, and D.C.Koningsberger, J. Phys. Chem., 94 (1990) 3711. 7. S.M.A.M.Bouwens, J.A.R. van Veen, D.C.Koningsberger, V.H.J. de Beer, and R.Pnns, J. Phys. Chem., 95 ( 1991) 123. 8. W.Niewman, B.S.Clausen, and H.Topsoe, Catal. Lett., 4 (1990) 355. 9. S.M.A.M.Bouwens, F.S.M. van Zon, M.P. van Dijik, A.M. van der Kraan, V.H.J. de Beer, J.A.R. van Veen, and D.C.Koningsberger, J. Catal., 146 (1994) 375. 10. W.Eltzner, M.Breysse, M.Lacroix, and M.Vnnat, Polyhedron, 5 (1986) 203. 11. E.I.Stiefel, T.R.Halbert, C.L.Coyle, L.Wei, W.-H.Pan, T.C.Ho, R.R.Chianelli, and M.Daage, Polyhedron, 8 (1989) 1625. 12. Y.Okamoto, M.Odawara, H.Onimatsu, and T.Imanaka, Ind. Eng. Chem. Res., 34 (1995) 3703. 13. Y.Okamoto, A.Maezawa, H.Kane, and T.Imanaka, J. Molec. Catal., 52 (1989) 337. 14. J.C.Duchet, E.M. van Oers, V.H.J. de Beer, and R.Prins, J. Catal., 80 (1983) 286. 15. J.P.R.Vissers, V.H.J. de Beer, and R.Prins, J. Chem. Soc. Faraday Tram.l, 83 (1987) 2145. 16. Y. Okamoto, Y.Kobayashi, and T.Imanaka, Catal. Lett., 20 (1993) 49. 17. Y.Sakamoto, N.Togashi, O.Terasaki, T.Ohsuna, Y.Okamoto, and K.Hiraga, Materials Science and Engineering A, Elsevier, in press. 18. Y. Okamoto, H.Katsuyama, K.Nakai, and O.Terasaki, in preparation. 19. K.Tanaka and T.Okuhara, Catal. Rev. -Sci. Eng., 15 (1977) 249. 20. Y.Okamoto, Y.Katoh, Y.Mori, T.Imanaka, and S.Teranishi, J. Catal., 70 (1981)445. 21. N.Topsoe and H.Topsoe, J. Catal., 84 (1983) 386. 22. I.Alstrup, I.Chorkendorff, R.Candia, B.S.Clausen, and H.Topsoe, J. Catal., 77 (1982) 397. 23. T.R.Halbert, T.C.Ho, E.I.Stiefel, R.R.Chianelli, and M.Daage, J. Catal., 130 (1991) 116. 24. U.Riaz, O.J.Curnow, and M.D.Curtis, J. Am. Chem. Soc., 116 (1994) 4357. 25. S.Harris and R.R.Chianelli, J. Catal., 86 (1984) 400.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 1 l th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

513

Electrochemical Promotion of NO Reduction by CO and by Propene

Alejandra Palermo~, Mintcho S. Tikhov, Neil C. Filkin and Richard M. Lambert* Department of Chemistry, University of Cambridge, Lensfield Road, Cambridge, CB2 1EW, U.K. Ioannis V.Yentekakis and Constantinos G. Vayenas Department of Chemical Engineering, Patras, GR-26500, Greece

Electrochemical promotion (EP) provides an efficacious means of catalyst promotion. The effects are reversible and the phenomenon provides a uniquely effective and controllable means for in s i t u tuning of the working catalytic system. EP studies of the catalytic chemistry of NO reduction by CO and by propene over Pt films supported on ~"-alumina (a sodium ion conductor) demonstrate that major enhancements in activity are possible when Na is electrochemically pumped to the catalyst surface. Both reactions exhibit strong electrochemical promotion under appropiate conditions of temperature, gas composition and catalyst potential. The data indicate that Na increases the strength of NO chemisorption relative to CO or propene, a process that is accompanied by weakening of the N-O bond, thus facilitating NO dissociation, thought to be the reaction initiating step. The overall kinetic behaviour and the selectivity towards N2 formation on catalyst potential are in agreement with this hypothesis. XP spectroscopy data confirm that the mode of operation of the electrochemically promoted Pt film does indeed involve reversible pumping of Na to or from the solid electrolyte. 1. INTRODUCTION Electrochemical promotion (EP) of catalytic activity has been described for many metal-catalysed reactions [ 1], including reactions ocurring at the metal/solution interface [2]. Back-spillover ions pumped from the electrolyte modify the metal work function with concomitant changes in adsorption enthalpies and hence reaction activation energies. These EP systems are characterised by a very large value of the ratio between the change in turnover rate and the rate of charge transport through the electrolyte (up to 105). Reported gains in reaction rate with respect to the unpromoted rate lie in the range 2-100 [3]. ~tpermanentaddress: IN'IEMA-CONICET,Universidad.Nac. de Mar del Plata, 7600 Mar del Plata, Argentina

514 Here we describe an EP study of catalytic NO reduction by CO and by propene over Pt/[~" alumina. The latter process is especially significant for the removal of NO in oxidising environments, a key problem that must be overcome for pollution abatement from lean-burn gasoline engines. Despite the fact that the NO+CO reaction over Pt has been intensively studied over a range of conditions and with a variety of catalysts ranging from single crystals [4, 5, 6, 7, 8] to practical dispersed materials [9, 10, 11, 12] there have been no previous studies of the effects of alkali promotion. Similarly, despite the extensive recent literature on the catalytic chemistry of NO+ hydrocarbon reactions [13, 14, 15, 16, 17, 18] very few studies on the effects of alkali promotion have been carried out. Electrochemical promotion of a Pt film using 13" alumina support as ion conducting provides an effective way of examining the effects of Na promotion in a controllable and reversible manner. Large changes in catalyst activity and selectivity can be induced in both reaction systems and our results strongly suggest that NO dissociation is the key reaction initiating step. 2. E X P E R I M E N T A L Three different specimens were used for the kinetic and spectroscopic experiments. The catalyst samples (working electrodes) consisted of porous but continuous thin Pt films deposited on the 16 mm diameter discs of Pt/I3"-AI203 as described in detail elsewhere [ 19]. Au reference and counter-electrodes were attached to the other face of the electrolyte wafer by vacuum evaporation producing good adhesion and extremely low resistance (
N O + C O ---> C O 2 + 2

(1)

2 N O + C O ---> C O 2 + N 2 0

For the NO+propene reaction the rN2 values were obtained directly from experiment by gas chromatographic measurement of the N2 production.

515 Catalytic rate measurements under potentiostatic or galvanostatic conditions were carried out using a galvanostat-potentiostat (Amel type 553). The reactant gas mixture was delivered at total flowrates of 1-2 x 10 -4 mol sec -1, with partial pressures PNO, Pco, Ppropene varied between 0 - 6.5 k Pa, 0 - 1.5 k Pa, 0 - 0.4 k Pa, respectively with PHe, bringing the total pressure to 1 atmosphere in every case. Conversion of the reactants was typically =15%. Control experiments confirmed that the Au reference and counter electrodes were catalytically inert under all conditions. The sample used for the XPS measurements was first tested in the EP reactor to ensure that it exhibited the same catalytic behaviour as that of the samples used to acquire the reactor data. XPS measurements were carried out in a VG ADES 400 UHV spectrometer system. The sample was mounted on a molybdenum block resistively heated by imbedded, electrically insulated tungsten filaments. XP spectra were acquired with Mg Ka radiation with the working (Pt) electrode always at ground potential; appropriate electrochemical potentials (Vwc) were applied between the working (Pt) electrode and the Au counter electrode by applying voltage bias to the latter. The potential of the working electrode with respect to the Au reference electrode (Vwr) was also measured. Quoted binding energies are referenced to the Au 4t"7/2 emission at 83.8 eV; the Au reference spectra were provided by a grounded Au film deposited on the outer face of the quartz sample holder, as illustrated in the inset to Figure 6. 3. RESULTS 3. 1. NO Reduction by CO Steady state measurements of NO decomposition in the absence of CO under potentiostatic conditions gave the expected result, namely rapid self-poisoning of the system by chemisorbed oxygen: addition of CO resulted immediately in a finite reaction rate which varied reversibly and reproducibly with changes in catalyst potential (VwR) and reactant partial pressures. Figure 1 shows steady state (potentiostatic) rate data for CO2, N2 and N20 production as a function of VWR at 621 K for a constant inlet pressures (P~ P~ of NO and CO of 0.75 k Pa. Also shown is the VWR dependence of N2 selectivity where the latter quantity is defined as SN ~ =

rN; rN2 + rN:O

(2)

Note that there is a sharp increase in activity as VWR is reduced below --0.2 V (Na pumping towards the Pt catalyst) with a concomitant threefold enhancement in the selectivity to N2. As it can be seen from Figure 1, the highest selectivities to nitrogen production always occur in the presence of the highest Na loading (most negative potential). The N20 rate actually goes through a maximum at VWR=-0.2 V. Control experiments were carried out in which the total flowrate was varied by a factor of 2 in order to check that the observed step change in

516 activity is due to a true increase in catalyst activity and is not influenced by mass transfer limitations or by multiple steady states in the reactor. As in previous EP studies with 13" alumina [20,25] for any fixed gaseous composition the promotional effect of Na is fully controlled by VWR and the resulting Na coverage. Upon current interruption the promoted rates remain practically constant over periods of many min. Potentiostatic imposition of the initial potential is necessary to restore the initial (unpromoted) rate [20]. 10

1

P NO= Pco = 0.75 k P~z, T = 621 K

'7r I,,.

r9 8

CO

0 E

'o

2

Io

~e (D

0

>"

oo

.~

selectivity

Z

V wr

0

0 -200mV

9 +200mV

.e..,

Z

04

9 +lO00mV

(clean Pt)

04

0

N2

T=621 K, PNO=0.52 kPa &

V wR / volts

-2

-'1

zx 4 ~

zx

()

zx

zx

1

Figure 1- NO+CO. Effect of VWR on CO s, N 2, N O rates and N 2 selectivity.

0

0.5

P co / kPa

1

1.5

Figure 2: NO+CO. Effect of Pco on N=, rate as a function of catalyst potential.

Figure 2 depicts the dependence of N2 rate on Pco at fixed PNO= 0.52 k Pa for three different values of the catalyst potential. VWR=+1000 mV corresponds to the clean Pt surface (unpromoted rate) and VWR=- 200 mV corresponds to a sodium promoted surface. Both CO2 and N2 rates exhibit Langmuir-Hinshelwood behaviour and as can be seen from Figure 2 for N2 rate, increased levels of Na result in a systematic increase in the CO partial pressure (P'co) necessary for inhibition. The N20 rate also exhibits Langmuir-Hinshelwood kinetics, but the effect of increased Na is somewhat different: in particular, high levels of Na tend to suppress the N20 rate and there is no systematic shift in P'co. 3. 2. NO Reduction by Propene In the case of NO reduction by propene, the only detectable reaction products were CO2, N2, N20 and H20. The overall mass balance was found to close within 5% as observed by a combination of GC and mass spectroscopic analyses. Figure 3 shows the effect of varying the catalyst potential on the rate of production of CO2, N2, N20 and on the selectivity towards nitrogen formation, SN2. As can be seen from this figure, both the CO2 and N2

517 reaction rates exhibit volcano-type behaviour: the rate of production is low at both high positive and high negative catalyst potentials, corresponding to zero and high coverage of sodium species, respectively. The region of strong electrochemical promotion corresponds to intermediate values of VWR, where a sharp asymmetric peak in rate is observed. In this promotion regime (+ 100 to -350 mV), the activity shows an exponential dependence on VWR. The behaviour of N20 production is similar, though the peak is not so sharp, and the rate at high negative potentials is higher than at high positive ones by approximately a factor of two. The selectivity towards N2 formation v e r s u s N20 production is thus slightly lower at high negative Vw-Rthan at high positive VWR, with a maximum (0.8) at -0.3 V. 12 10

P No=1.3 kPa, P =0.6 kPa t propene T--648 K &

":

"...

8

i i

_r ~

4

T=648 K, PNO=1.4 kPa

._>

10

-0.5

300mV

0

9

9

O

zx

0.5

Figure 3: NO+propene. Effect of V WR CO 2,N2' NO rates and N2 selectivity.

9

8

3<"..... ~(.

0 V WR / Volts

~

o

Z

".

N2 0 0 -"

9 -200mV

|

z 2

-3(X)mV

O 0mY

%

(9 u) 04

!

9 o

.'t=

IN2

V wr 0

[3 m

""':A

o

6

20

15

selectivity 9......,.

"&....----A. . . .

1

O ZX

0.1

A 0.2

0.3

0.4

P(propene) / kPa OR

Figure 4: NO+propene. Effect of P

on

N rate as a function of catalyst potential. 2

Figure 4 shows the effect on the N2 rate of varying the propene partial pressure for fLxed NO partial pressure and for several values of catalyst potential. The observed rate exhibits Langmuir-Hinshelwood behaviour with a characteristic rate maximum reflecting competitive adsorption of the two reactants. As the catalyst potential is changed to more negative values, corresponding to supply of sodium to the catalyst, the rate maximum shifts to higher propene partial pressure. The CO2 and N20 rates showed essentially similar behaviour as Pprot~ne was varied. Figure 5 summarises results for the CO2, N2 and N20 formation rates for the dependence of apparent activation energies on catalyst potential. Although there is a notable increase in activation energy with increased Na coverage in each case, the variation is not as abrupt at that characteristic of EP CO oxidation [24] and NO+CO reactions.

518

3. 3. XPS Study of 13"alumina Figure 6 shows Na ls spectra acquired at 500 K under electrochemical bias conditions that replicate those in the reactor. Spectrum A was obtained with Vwr ~- +600 mV, corresponding to the electrochemically cleaned surface. Spectrum B was obtained with Vwr--400 mV, corresponding to the Na-promoted surface. The high binding energy feature at 1072.8 eV is assigned to Na on the Pt surface and the smaller feature at lower binding energy (1071 eV) is ascribed to Na at the I]"-alumina surface, visible through cracks in the porous polycrystalline film. Most importantly, the spectral behaviour was reproducible and reversible: it was possible to cycle back and forth between the two types of spectra by repeatedly reversing the sign of Vwr The same behaviour was observed when Na was vacuum-deposited on the working electrode under open circuit conditions: application of positive bias (VwR > 0) resulted in complete disappearance of vacuum-deposited Na from within the XPS samping depth. This behaviour clearly shows that the electropumped Na and the vacuum-deposited Na behave identically on the surface of the Pt electrode. Further studies to quantify the transport and morphological properties of Na on the catalyst surface are in progress.

Figure 5: NO+propene. Dependence of apparent activation energy for CO 2, N 2, NO rates on VwR.

Figure 6: XPS of Pt/beta alumina acquired under electrochemical bias. A) cleaned surface, B) Na-promoted surface.

4. DISCUSSION It is important to point out that in the discussion that follows the term "Na coverage" is used; this does not imply that the promoter is thought to be present in the form of chemisorbed metallic sodium as it would be in vacuum. The reactive gas atmosphere is expected to lead to the formation of surface compounds of Na, and single crystal data indicate

519 that stable Na-CO complexes [23] or Na carbonates [24] can be formed, depending on the composition of the ambient gas. Adsorbed polar alkali compounds lead to large decreases in work function, of the same order as those produced by the alkali metal itself, so the general theory of electrochemical promotion [1 ] is nevertheless applicable. The NO+CO reaction exhibits strong electrochemical promotion under Na pumping to the catalyst when the partial pressures of both reactants are similar. This and the EP behaviour as a function of CO pressure may be rationahsed as follows. At low partial pressures of CO the rate is low due to the restricted availability of CO~a~; CO+O is the rate limiting step and so the EP effect is small. At intermediate PCO the rate determining step becomes NO dissociation, as the coverages of NO and CO are similar. Thus a strong electrochemical promotion effect is observed: supply of sodium to the catalyst causes enhanced NO dissociation and hence a dramatic rate increase (Figure 1). At high CO partial pressures the EP effect is again attenuated as a result of limited NO coverage and site blocking by CO-Na surface complexes [20]. The selectivity towards N2 formation versus N20 production increases with increasing Na coverage; this follows from increased availability of N(a ) + O(a ) versus NO(a ), again strongly supporting the idea of Na-enhanced NO dissociation. The maximum p values obtained are 13 and 1.5 for N2 and N20 respectively, at a gas composition of P0t~o = P0co = 0.75 k Pa. For P0co = 0.75 k Pa, Na pumping to the catalyst leads to an increase in SN2 by up to a factor of three. Sharp changes in activity which occur as a function of Pt~o at fixed catalyst potential resemble the behaviour observed for the CO+O2 reaction over Pt/I]"alumina [20] and, as in that case, are ascribed to a surface phase transition. At low PNo the surface is dominated by islands of CO; reaction occurs only at the peripheries of these islands, resulting in a low rate. At sufficiently high PNo the CO islands are disrupted by NO chemisorption and the rate rises sharply as intermixing of the reactants occurs. This model is strongly supported by the observed effects of Na promotion. The higher the Na coverage (more negative catalyst potentials) the lower the value of PNO at which the phase transition occurs, reflecting the increased strength of NO chemisorption relative to that of CO. In the case of NO reduction by propene, very large changes in activity also occur under the influence of Na pumping to the Pt catalyst. The overall activity towards formation of carbon dioxide is strongly enhanced when Na is pumped to the Pt catalyst (Figure 3). Activity towards formation of the other products shows a similar dependence on catalyst potential. The maximum gain in rate over the clean surface rate is of the order of a factor of 10 in the case of nitrogen and CO2 production. In the region of strong electrochemical promotion (+100 to - 350 mV), the activity towards formation of all products shows an exponential dependence on catalyst potential, as predicted by theory [25]. It is evident from Figure 3 that there is a precipitous fall in rates as the coverage of Na species increases beyond a critical valued, i.e. the regime of electrochemical promotion (~ 0 -- -350 mV) is followed by a regime of strong poisoning. Results obtained with related systems, including single crystal/electron spectroscopy data obtained with model planar

520 catalysts, indicate that the poisoning behaviour is associated with the formation of stable surface compounds of Na which serve to block active sites at high Na loadings. In this regard, there is a striking difference between the two reactions at very high negative catalyst potentials, NO+propene exhibits a much greater sensitivity to poisoning while the reaction of NO+CO is relativily resistant to poisoning. To put it another way, promotion of NO+propene occurs over a much narrower range of Na coverage. Logically, one should consider this in terms of possible differences in the surface chemical behaviour of CO and propene on a Nadoped Pt surface in the presence of chemisorbed oxygen, in addition to other adsorbed species. Two factors may contribute to this, although further spectroscopic data are needed in order to provide a definitive answer. Factor 1" single crystal studies show that the surface compounds formed in the NO+CO system under reaction conditions undergo significant agglomeration to form 3-dimensional crystallites. In other words, pumping the equivalent of one monolayer of Na to the Pt does n o t lead to the formation of a monolayer of Na compound which would completely block the surface and presumably completely poison the system. Propene has a much greater tendency to deposit carbon on Pt than does CO [26,27,28]. The presence of strongly adsorbed carbon atoms could inhibit agglomeration of Na surface compounds, thereby maximising the number of Pt sites that are affected by a given amount of Na. Factor 2: the chemisorption of both CO and NO is strengthened by coadsorbed Na [29]; the chemisorption bond of propene should be weakened by Na. Therefore at high levels of Na the coverage of propene should be strongly attenuated, with a corresponding large decrease in reaction rate. The observed dependence of the CO2 rate on propene partial pressure for fixed NO partial pressure and for a range of different catalyst potentials (Figure 4) demonstrates that the system exhibits classical Langmuir-Hinshelwood behaviour - a characteristic rate maximum reflecting competitive adsorption of the two reactants. Under all conditions of partial pressure, there was an overall increase in activity with increased pumping of Na to the catalyst: we again associate this with dissociation of NO induced by the Na promoter. The rate maxima shift systematically to higher propene partial pressures as the sodium coverage is increased, reflecting the increase in the binding of NO relative to propene with increasing Na coverage. This kind of behaviour is exactly what one would expect in the case of an electropositive promoter: the chemisorption strength of electron donors (propene) should be decreased whereas the chemisorption of electron acceptors (NO and its dissociation products) should be enhanced. For both reactions studied, NO+CO and NO+propene, the effect of electrochemically pumped Na in increasing the extent of NO dissociation is large and significant. This is because unpromoted low index planes of Pt, Pt(111), are relatively inert towards NO dissociation and we adscribe the NO dissociation as the key reaction-initiating step. Such dissociation of diatomic molecules in the field of coadsorbed cations has been discussed in detail by Lang et al [29]. The rates of production of CO2, N2 and N20 all depend on

521 dissociation of NO for their formation, as it can be analised from the following proposed elementary steps: CO(a ) 4" O(a ) "~ CO 2 N( a ) + N( a ) ===)N 2

(3)

N(a) + NO(a ) --=) N 2 0

The observed increase in the selectivity towards N2 is a consequence of NO dissociation, i.e. a decreased amount of molecular NO, and increased amount of atomic N on the surface, both factors favours the second of the above reactions over the last one. This dissociative mechanism is the generally accepted pathway under ultra high vacuum conditions [7, 8, 30]. However, a recent study by Klein et al [9] has questioned the validity of the dissociative mechanism under atmospheric pressure conditions in favour of a non-dissociative mechanism. A particular difficulty with the non-dissociative mechanism is that it cannot readily account for the lack of reactivity of low index planes of Pt. Our EP results strongly suggest that the dissociative mechanism holds, even in the high pressure regime. The catalyst film consist of large polycrystalline Pt particles whose surfaces are dominated by low index planes that are inactive for NO dissociation. The low rates observed at high positive catalyst potentials (Na-free system) may be ascribed to defects and high index planes that are inevitably present at crystallite edges. Both N2 and N20 are produced in this region as there is a mixture of molecular NO plus atomic N and O. Na supplied to the Pt surface strongly enhances the overall activity by inducing NO dissociation on the otherwise ineffective low index planes in accord with both theory and experiment. For the reduction of NO with propene, the catalyst potential dependence of the apparent activation energies does not show a step change and is much less pronounced than it is for the CO+O2 and NO+CO systems. There is persuasive evidence [20] that the step change is associated with a surface phase transition - the formation or disruption of islands of CO. It is reasonable to assume that this phenomenon cannot occur in the NO+propene case, since there is no reason to expect that large amounts of chemisorbed CO can be present under a n y conditions. That there should be a difference in this respect between CO+O2/CO+NO on the one hand, and NO+propene on the other hand, is therefore understandable; however, the chemical complexity of the adsorbed layer in the NO+propene precludes any detailed analysis of the Ea(VwR) effect. The central assumption underlying all of the preceding discussion is that under EP conditions, reversible changes in VWR correspond to the reversible pumping of Na to/from the Pt from/to the solid electrolyte. Our XPS data clearly show that such reversible transport of Na between 15"-alumina and the surface of the Pt film does indeed occur under the conditions of voltage and temperature that were used for the reactor studies; increasingly negative potentials corresponding to increasing amounts of Na. Furthermore, we have demonstrated the equivalence of vacuum-deposited and electrochemically pumped Na on the catalyst

522 surface. These are important observations that serve to underpin the fundamental theory developed by Vayenas et al [1]. ACKNOWLEDGEMENTS Support under grant GR/J00632 from the UK EPSRC is gratefully acknowledged. MST acknowledges support from British Gas plc. AP and RML acknowledge additional support under a grant from the British Council and Fundaci6n Antorchas. REFERENCES

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 29 30

C.G. Vayenas, S. Bebelis, I.V. Yentekakis, and Lintz, H. G., Catalysis Today, Vol. 11, No 3, p. 303. Elsevier, Amsterdam, 1992. S.G. Neophytides, D. Tsiplakides, P. Stonehart, M.M. Jaksic and C.G. Vayenas, Nature 370 (1994) 45. C. Pliangos, I.V. Yentekakis, X. Verykios and C.G. Vayenas, J. Catal. 154 (1995) 124. J. A. Rodriguez and D. W. Goodman, Surf. Sci. Reports 14 (1991) 1. Y. O. Park, W. F. Banholzer and R. I. Masel, Surf. Sci. 155 (1985) 341. W. F. Banholzer, R. E. Parise and R. I. Masel, Surf. Sci. 155 (1985) 653. D'Arcy Lorimer and A. T. Bell, J. CataL 59 (1979) 223. B. A. Banse, D. T. Wickham and B. E. Koel, J. Catal. 119 (1989) 238. R. L. Klein, S. Schwartz and L. D. Schimdt, J. Phys. Chem. 89 (1985) 4908. D. N. Belton and S. J. Scmieg, J. Catal. 138 (1992) 70. D. N. Belton and S. J. Scmieg, J. Catal. 144 (1993) 9. S. E. Oh., G. B. Fischer, J. E. Carpenter and D. W. Goodman, J. Catal. 100 (1986) 360. A. Obuchi, A. Ohi, M. Nakamura, A. Ogata, K. Mizuno and H. Ohuchi, Appl. Catal. B: Environmental 2 (1993) 71. J.R. Hardee and J.W. Hightower, J. Catal. 86 (1984) 137. S. Naito and M. Tanimoto, Chem. Lett. 1993 1935. R. Burch, P.J. Millington, A.P. Walker, Appl. Catal. B: Environmental 4 (1994) 65. T. Miyadera and K. Yoshida, Chem. Len. 1993 p. 1483. H. Hamada, Y. Kintaichi, M.Sasaki, T. Ito and M. Tabata, Appl. Catal. 75 (1991) L 1. C. G. Vayenas, S. Bebelis, S. Neophytides and I.V. Yentekakis, Appl. Phys. A 49 (1989) 95. I.V.Yentekakis, G.D. Moggridge, C.G.Vayenas and R.M.Lambert, J.Catal. 146 (1994) 292. I.V. Yentekakis, S. Neophytides and C.G. Vayenas, J. Catal. 111 (1988) 152. I.V. Yentekakis and S. Bebelis, J. Catal. 137 (1992) 278. J.C. Bertolini, P. Delichere and J. Massardier, Surf. Sci. 160 (1985) 531. I.R. Harkness and R.M. Lambert, J. CataL 152 1995 211. C.G.Vayenas, S. Bebelis and S Ladas, Nature 343 (1990) 625. N.R.Avery, N.S.Sheppard, Proc.Roy.Soc.Lond. A 405 (1986) 1 R.J.Koestner, J.C.Frost, P.C.Stair, M.A.Van-Hove, G.A.Somorjai, Su~Sci. 116 (1982) 85. M.Salmeron, G.A.Somorjai, J.Phys.Chem. 86 (1982) 341. N.D. Lang, S. Holloway and J.K. Norskov, Surface Science 150, 24 (1985). G. Pirug and H.P. Bonzel, J. Catal. 50 (1977) 64.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

523

P r o m o t i o n of m o l e c u l a r h y d r o g e n on solid acid c r a c k i n g a c t i v i t y T. Shishido, T. Nagase, K. Higo, J. Tsuji and H. Hattori Center for Advanced Research of Energy Technology, Hokkaido University, Sapporo 060, Japan IR spectroscopy of adsorbed pyridine, and temperature-progrRmmed desorption (TPD) of hydrogen and deuterium were studied for several solid acid catalysts to examine whether the formation of protonic acid sites from hydrogen molecules occurs and promotes acid-catalyzed reactions. The catalysts examined were Co.Mo/SiO2-A1203, Co.Mo/A1203, SIO2-A1203, H-ZSM-5, physical mixture of Pt/SiO2 and H-ZSM-5, and Pt/SO42"-ZrO2, and the reaction chosen was cumene cracking. For all catalysts, the formation of protonic acid site was observed by heating in the presence of hydrogen, and the amount of hydrogen (or deuterium) adsorbed increased as the adsorption temperature was raised. The promotion effects of hydrogen on the catalytic activity for cumene cracking were observed for the catalysts containing metallic components. The concept "molecular hydrogen-originated protonic acid site" is proposed. 1. I N T R O D U C T I O N It is established that the active sites of solid acid catalysts for most of the acid-catalyzed reactions are protons located on the surfaces. Although a proton is formed by heterolytic dissociation of hydrogen molecule, promotion effects of hydrogen on the acid-catalyzed reactions have scarcely been reported. Most of the papers studying hydrogen effects on the acid-catalyzed reactions reported suppression effects of hydrogen; the reactions were retarded in the presence of hydrogen[ 1-5]. In recent years, several papers have reported promotion effects of hydrogen. Baba et al. reported that the promotion effects of hydrogen on ethylbenzene disproportionation over Ag-Y zeolite are caused by the formation of protons from hydrogen molecules accompanied by reduction of Ag § ions to Ag o metals in the presence of hydrogen[6-8]. The protons thus formed are eliminated by the reverse reaction of Ag o to Ag + if the gaseous hydrogen is removed from the reaction mixture. Although the formation of a proton by reduction of metal ion with hydrogen is generally observable, the unique feature with Ag-Y is the occurrence of the reverse reaction as the gas phase hydrogen is removed. Prominent effect of hydrogen was reported by Sachtler et al. with Pd-Y zeolite for isomerization of methylcyclopentane in the presence of hydrogen[9-11]. They proposed "electron-deficient"Pd-H + adducts as the active sites in the presence of hydrogen. Hosoi et al. reported that Pt/SO42--ZrO2 persists a high activity for a long period in alkane skeletal isomerization when the reaction is carried out in the

524 presence of hydrogen[12]. They ascribed a high and stable activity in the presence of hydrogen to a removal of carbonaceous residues deposited on the catalyst by hydrogenation. We have studied the hydrogen effects on the catalytic acyivity of Pt/SO42-ZrO2 for alkane skeletal isomerization, and concluded that the promotion effect of hydrogen is caused by the formation of protons from molecular hydrogen[1317]. Hydrogen molecule is dissociated on the platinum species to form hydrogen atoms which spiUover onto the support. The spiltover hydrogen atom migrates on the support to reach Lewis acid site where hydrogen atom loses an electron to form a proton. The proton is stabilized on the oxygen atom nearby the Lewis acid site. The electron trapped at Lewis acid site reacts with a second hydrogen atom to form a bond of Lewis acid-H-. We proposed the protons originating from molecular hydrogen act as active sites on Pt/SO42--ZrO2 . Recently, the formation of proton from molecular hydrogen was proposed for the physical mixture of NiS and USY zeolite[18]. Skeletal isomerization of alkanes possibly proceeds by two mechanisms; metal-acid bifunctional mechanism and acid-catalyzed monofunctional mechanism. Since Pt/SO42"-ZrO2 contains metallic component, there still remains a possibility that the promotion effect of hydrogen on skeletal isomerization is caused by some interaction of hydrogen with the surface, but not by the formation of protons. It would be more clearly demonstrated that the protons formed from hydrogen molecules act as active sites for acid-catalyzedreactions if the reaction catalyzed only by acid sites are promoted in the presence of hydrogen over Pt/SO42"-ZrO2. In addition, it would be important to examine whether the concept "molecular hydrogen-originated protonic acid site" can be extended to the catalysts other than Pt/SO42--ZrO2 . The present paper aims to clarify the following points. 1 Does the generation of protonic acid sites originating from molecular hydrogen occur for the catalysts other than Pt/SO42"-ZrO2? 2 Are the promotion effects of hydrogen on the typycal acid-catalyzed reaction observable for Pt/SO42"-ZrO2 and the other catalysts? As a pure acid-catalyzed reaction, cumene cracking was chosen, and hydrogen effects were examined for Co.Mo/SiO2-AI203, Co.Mo/A1203, SIO2Al203, H-ZSM-5, a physical mixture of Pt/SiO2 and H-ZSM-5, and Pt/SO42-ZrO2. 2. E X P E R I M E N T A L

2.1 Catalyst preparation The Pt/SO42--ZrO2 was prepared as follows. The sulfated ion treated Zr(OH)4 was prepared by impregnation of Zr(OH)4 with 1N H2SO4 aq. solution followed by filtration and drying at 383K. The Zr(OH)4 was obtained by the hydrolysis of ZrOC128H20 with aqueous ammonia. The obtained gel was washed with deionized water until no C1- ions could be detected. The Pt/SO42--ZrO2 sample (0.5 wt%Pt) was prepared by impregnation of SO42--ZrO2 with 1% H2PtC16 aq. solution followed by drying at 383K and calcination at 873K in air. The amount of S remained in the resulting catalyst was 1.5 wt% determined by XRF. The catalyst was pretreated in a hydrogen stream at 623K for 2 h before

525 use for reaction. For IR and TPD measurements, the catalyst was treated further in a vacuum at 623K. Pt/SiO2 was prepared by impregnation of silica supplied from the Catalysis Society, Japan (JRC-SIO-1) with H2PtC16 aq. solution followed by drying, and calcining at 773K for 5 h. The content of Pt was adjusted to be 2 wt%. H-ZSM-5 was prepared by ten times ion exchange of Na-ZSM-5 supplied from TOSOH with 1N NH4C1 aq. solution. The SIO2/A1203 molar ratio was 23.3. The catalyst was finally calcined at 803K for 3 h. The physical mixture of the Pt/SiO2 and HZSM-5 was prepared by mixing the two components in the ratio 1:10 (Pt/SiO2 : H-ZSM-5) in an agate mortar. Co.Mo/SiO2-A1203 was prepared as follows. Into a solution of cobalt nitrate, aqueous ammonia was added to form aqueous solution of cobalt ammoniun complex. To the solution, aqueous molybdenum ammoniun was addes to make a mixed solution containing cobalt and molybdenum. Silica-alumina supplied from the Catalysis Society, Japan(JRC-SAL-2) was impregnated with the mixed solution followed by drying at 353K for 36 h, crushing into a powder, and calcining at 773K for 16h. The contents of Co and Mo were 4 and 8 wt% as CoO and MOO3, respectively. The Co.Mo/A1203 catalyst used as a reference catalyst was supplied from Cyanamid Co. Ltd. (HDS-20).

2.2 IR study of adsorbed p y r i d i n e A self-supported wafer of the sample was placed in an in situ IR cell with CaF2 windows, and pretreated. The pretreatment procedures varied with the catalyst samples. For Co.Mo/SiO2-A1203 and Co.Mo/A1203, the calcined sample was pretreated in a hydrogen flow at 673K for 2h followed by outgassing at 773K. For presulfided sample, the calcined sample was treated in a 3%H2S in H2 at 673K for 2h followed by outgassing at 773K for 17 h. For the physical mixture of Pt/SiO2 and H-ZSM-5, the calcined sample was outgassed at 773K and cooled to 673K at which the sample was treated in a hydrogen flow for 2 h followed by outgassing at 773K for 17h. For H-ZSM-5, the calcined sample was outgassed at 773K for 17h. After the pretreatment, pyridine was introduced into the cell at a pressure of ca. 2 Torr for 15 rain at 423K, and then outgassed normally at 673K for 20 rain. In the case of Co.Mo/SiO2-AI203, the final outgassing was done at 423K. To examine the effects of hydrogen on the adsorbed pyridine, the pyridine-covered sample was exposed to 500Torr of hydrogen at room temperature, and heated stepwise from room temperature to 573K or 673K by 50K increments. Following the processes of heating in the presence of hydrogen, the sample was outgassed stepwise from 373K to 673K by 50K increments. All spectra were recorded on an FT/IR-5300 infrared spectrometer(JASCO) at room temperature. For determination of the number of protonic sites and Lewis acid sites on the surface, the integrated absorbances of the bands at 1450 cm -1 (due to pyridine chemisorbed on Lewis acid sites, L-Py) and 1490 cm "1 (due to both the L-Py and pyridine chemisorbed on protonic acid sites, B-Py) were used with the tangent background for all samples. The values obtained were normalized to the weight of the sample wafer. To obtain the apparent absorption coefficients of the bands, a known amount of pyridine was adsorbed on the sample, and the absorbance of each band was measured. Then, a small quantity of water which is sufficient to convert all Lewis acid sites into protonic acid sites was introduced into the IR-

526 cell. The apparent absorption coefficients were calculated from the changes of the absorbances at 1450 and 1490 cm "1 on exposure to water vapor. As the apparent absorption coefficient for a certain band is different for a different wafer, the calibration of the apparent absorption coefficient was performed for each wafer.

2.3 T e m p e r a t u r e - P r o g r a m m e d Desorption of h y d r o g e n and/or deuterium The p r e t r e a t m e n t conditions were the same as those for IR study of adsorbed pyridine. The sample pretreated in a vacuum was exposed to ca. 300Torr D2 at different t e m p e r a t u r e s for 1 h. After cooling to room temperature, the sample was outgassed for 10 rain prior to TPD run. TPD was run at a heating r a t e of 10K rain "1, and the desorbed gases were analyzed by mass spectrometry. 2.4 R e a c t i o n p r o c e d u r e s for c u m e n e c r a c k i n g A pressurized flow reactor and an atmospheric pulse reactor were employed for carrying out cumene cracking. For the reaction over Co.Mo/SiO2-AI203 and SIO2-A1203, the flow reactor was used with carrier flow rate of 100ml min "1 and cumene feed rate of 14ml min "1 under the pressure of 30 arm. For the reaction over Pt/SO42"-ZrO2, the physical mixture of Pt/SiO2 and HZSM-5, and H-ZSM-5, the pulse reactor was used. A dose of cumene, 0.5ml (0.036mmol), was passed over the catalyst in a carrier flowing at 50ml min "1, and the products were trapped at 77K before being flash-evaporated into a gas chromatographic columns(PEG-20M and VZ-7). 3. R E S U L T S

3.1 IR study of the effect of h y d r o g e n on the acid site The acid-site types were examined by IR spectroscopy of adsorbed pyridine. For Co.Mo/SiO2-A1203, no pyridine remained on the surface after outgassing at 673K. There are no acid sites strong enough to retain pyridine against outgassing at 673K. Therefore, the catalyst was exposed to pyridine and outgassed at 423K, and then the catalyst was heated at 523K or 623K in the presence of 300 Torr of hydrogen. The variation of the percentages of protonic acid sites and Lewis acid sites as a function of the temperature at which the Co.Mo/SiO2-AI203 was heated in Fig. 1 Change of the acid sites on Co.Mo/SiO2A1203 with hydrogen t r e a t m e n t and the the presense of hydrogen is following outgassing. shown in Fig. 1.

527

Although most of the acidic sites on the surface are Lewis acid, the fraction of protonic acid sites increased as the catalyst was heated in the presence of hydrogen. The increase in the fraction of protonic acid sites was extended as the heating temperature was raised. This indicates that protonic acid sites were generated by heating in the presence of hydrogen. The protonic acid sites thus formed were eliminated by outgassing gas phase hydrogen at 673 K. The protonic sites are generated and eliminated in response to heating in the presence and absence of hydrogen. These observations are essentially the same as those observed for Pt/SO42"-ZrO2 catalyst, though the protonic acid sites formed and eliminated are much stronger for Pt/SO42--ZrO2 than for Co.Mo/SiO2-A1203. For presulfided Co.Mo/SiO2-Al203, the formation and elimination of the acid sites were almost the same as those observed for non-sulfided Co.Mo/SiO2-AI203. Presufidation with hydrogen sulfide did not affect much the conversion of the acid sitescaused by gas phase hydrogen. For Co.Mo/A1203 catalyst, no IR bands ascribed to pyridine were appreciable after outgassing at 423K. The acid sites on Co.Mo/A1203 are much weaker than those on Co.Mo/SiO2-A]203. For the physical mixture of Pt/SiO2 and H-ZSM-5, pyridine was retained on the surface after outgassing at 673K. On t~e surface of the mixed catalyst, the acid sites as strong as those on H-ZSM-5 exist as expected. The IR spectrum of adsorbed pyridine changed markedly when the sample was heated in the presence of hydrogen. On raising the temperature, the IR band ascribed to protonic acid sites developed with concomitant decrease in the intensity of the IR band ascribed to Lewis acid sites. The variations of the amounts of protonic acid sites and Lewis acid sites as a function of the heating temperature in the presence of hydrogen are shown in Fig. 2. "7

-~

H2 exposure

Evacuation

9 8 i

,~6 a

o4

i2 o

~0

o

<

273 373 473 273 373 473 Temperature / K

573

673

Fig. 2 Change of the acid sites on the physical mixture of Pt/SiO2 and H-ZSM-5 with hydrogen treatment and the following outgassing.

528 On raising the temperature to 373K, the Lewis acid sites disappeared with an increase in the protonic acid sites. On outgassing the gas phase hydrogen, the Lewis acid sites increased and the protonic acid sites decreased. These results are essentially the same as those observed for PtYSO42--ZrO2 and HZSM-5. However, the decrease in the Lewis acid sites occurred at much lower temperature for the physical mixture of Pt/SiO2 and H-ZSM-5 as compared to HZSM-5. For H-ZSM-5, it requires 473K to convert Lewis acid sites into protonic acid sites. Addition of Pt/SiO2 to H-ZSM-5 enable to convert Lewis acid sites into protonic acid sites at a lower temperature. The recovery of the Lewis acid sites and the decrease in the protonic acid sites to the original values by outgassing gas phase hydrogen are rather slow and require a high temperature. The addition of Pt/SiO2 did not affect much the restoration of the Lewis acid sites of H-ZSM-5 by outgassing gas phase hydrogen. Temperature-Programmed D e s o r p t i o n ( T P D ) of h y d r o g e n (or deuterium) The amount and strength of hydrogen adsorption changed with the adsorption temperature for all catalysts examined. As the adsorption temperature was raised, the amount and strength of adsorption increased. TPD plots for D2 adsorbed on the physical mixture of Pt/SiO2 and H-ZSM-5 at different temperatures are shown in Fig. 3. Desorbed gases were composed mainly of D2, only small quantities of liD and H2 being detected though they are not shown in Fig. 3. The peak area increased with an increase in the adsorption temperature. The TPD plots for the physical mixture of Pt/SiO2 and H-ZSM-5 are similar to those for H-ZSM-5. For both catalysts, it appears that a step requiring a high energy exists in the adsorption of D2. Essentially the same phenomena were observed for Co.Mo/SiO2-AI203 and Pt/SO42--ZrO2 in the sense that the desorption peaks became larger and appeared at higher temperature as the adsorption temperature was raised. The TPD plots of D2 for presulfided Co.Mo/SiO2-A]203 are shown in Fig. 4. 3.2

d

3

~2

e~

~

1

e~

~

0

273

373

473

573

673

773

Temperature / K Fig. 3 TPD plots of deuterium adsorbed on the physical mixture of PffSiO2 and H-ZSM-5 at the temperatures indicated.

529

10

8 ~6 e~

= 4 9~- 2

~ 0 273

373

473

573

673

773

Temperature / K Fig. 4

TPD plots of deuterium adsorbed on Co.Mo/SiO2-A1203 at the temperatures indicated.

3.3 E f f e c t o f h y d r o g e n

on cumene

cracking

For Co.Mo/SiO2-A1203 catalyst, the promoting effect of hydrogen on the catalytic ativity was observable in cumene cracking under the hydrogen pressure of 30 atm as shown in Fig. 5. The conversion was less than 10 % in the nitrogen stream, while the conversion was about 30-35 % in the hydrogen stream. The catalytic activity was enhanced more than three times in the presence of hydrogen. Since the active sites for cumene cracking are protonic acid sites, it is suggested that the protonic acid sites generated in the presence of hydrogen act as catalytically active sites in the reaction. The product distributions in the presence of hydrogen and in the absence of hydrogen were the same except the formation of propane in the presence of hydrogen instead of propene which was formed in the absence of hydrogen. For presulfided Co.Mo/SiO2-A1203, the promoting effect of hydrogen was also observed. The conversions were about 5 % in the absence of hydrogen and about 20 % in the presence of hydrogen. 40 0

0

0

9

ffi 30

9 9

9

0

0

9 9

DO

N

~" 0 20 0

~ 0

E

O O

10 _0 o

0

0

0

0

O

O

|

0

Fig. 5

50

I

100 T i m e / min

l

150

200

Cumene cracking over Co.Mo/SiO2-Al203 at 673K and 30atm.

530

40

= 30

~o 20

- 9e ~ ~

o

e ~

o

o

E ~ o

I

0 0

Fig. 6

50

I

100 Time / rain

I

150

200

Cumene cracking over SiO2- A1203 at 673K and 30atm.

For SIO2-A1203, the conversions are plotted against the time on stream in Fig. 6 for both reactions in the hydrogen stream and nitrogen stream. The conversions were the same for both cases. Although the formation of protonic acid sites in the presence of hydrogen was observed by IR study of adsorbed pyridine, the hydrogen effect was not appreciable in the reaction. The conversions of cumene cracking over H-ZSM-5 and the physical mixture of Pt/SiO2 and H-ZSM-5 carried out in the pulse reactor are plotted against the pulse number in Fig. 7. With H-ZSM-5, the conversion rapidly decreased with the pulse number even in the presence of hydrogen. The conversion was not affected by the presence of hydrogen for H-ZSM-5. In contrast, the conversion was affected much by the presence of hydrogen for the physical mixture of Pt/SiO2 and H-ZSM-5. The conversion kept close to 100% for the initial 5 pulses where hygrogen was used as a carrier. As hydrogen was switched into helium at the 6th pulse, the conversion rapidly decreased with the following pulses. The conversion recovered to 70% when the carrier was switched back to hydrogen at the l l t h pulse. Although the conversion of cumene over the physical mixture in the presence of hydrogen included the conversion to isopropylcyclohexane to some extent, promotion effect of hydrogen was clearly demonstrated for the cracking activity of the physical mixture of Pt/SiO2 and H-ZSM-5. For Pt/SO42--ZrO2, essentially the same hydrogen effect as that of the physical mixture of Pt/SiO2 and H-ZSM-5 was observed. The conversion of cumene at 423K was constant (55%) with the pulse number in the hydrogen carrier. In the helium carrier, the conversion decreased to 2% at the 5th pulse. By switching the helium carrier back to hydrogen, the conversion gradually increased to 50% following the 5th pulse. The products were composed mostly of benzene and propene in the reaction under a helium carrier, and benzene and propane with a small amount of propene in a hydrogen carrier. Only propene underwent hydrogenation in the hydrogen carrier, but benzene and cumene did not. No deactivation in a hydrogen may be partly caused by a rapid hydrogenation of propene that otherwise polymerizes to form carbonaceous residue.

531

100

A

A

~OOO H2

He

H2

80 o

60

-

40

-

9

O

9

9

9

OQ

omm

O O

o

20

O

~

1

2

4

9 I

I,

6

8

o

.

Ooo9 o

10 12 14

pulse number

Fig. 7

9 Physical mixture of Pt/SiO2and H-ZSM5 9 H-ZSM5 Cumene cracking over H-ZSM5 and the physical mixture of Pt/SiO2 and H-ZSM-5 at 423K in a pulse reactor.

4. D I S C U S S I O N All the results can be explained in terms of the formation of protonic acid sites from molecular hydrogen. In the IR study of adsorbed pyridine, conversion of L-Py to B-Py was observed for all catalysts regardless of the existence of a metallic component. However, the formation of the protons seems to proceed by different mechanisms for solid acids containing metallic components and those without metallic components. Without metallic component, hydrogen molecules will not dissociate to hydrogen atoms. Hydrogen molecules are likely to be dissociated heterolytically on the Lewis acid sites to form protons and hydrides. On the solid acids with metallic component, hydrogen molecules are dissociated on the metallic component to form hydrogen atoms which spillover onto the support and reach Lewis acid sites. In TPD of adsorbed D2, it was observed for all catalysts that the desorption peaks appeared at a higher temperature and in a larger amount as the adsorption temperature increased. The increase in the amount of adsorbed D2 with the adsorption temperature suggests that a step requiring a high energy is involved in the adsorption process, though it is not definite which step needs a high energy. For the solid acids with metallic components, perhaps, the migration of hydrogen atoms on the support involves repetition of adsorption and desorption at Lewis acid sites, and the desorption step seems to require a high energy. For the solid acids without metallic component, hydrogen molecule should be directly adsorbed on the Lewis acid site to dissociate heterolyticaUy to form a proton and hydride; the proton forms acidic OH group and the hydride is adsorbed on the lewis acid site. The heterolytic dissociation of hydrogen molecule may be the step requiring a high energy. Without metallic components, the promotion effect of hydrogen on the

532 cracking activity was not appreciable, though the formation of protonic acid sites was observed by IR. During the reaction, hydrogen molecules compete with cumene molecules for adsorption on the seine sites. The adsorption strength of hydrogen molecule should be much weaker than that of c11mene molecule, and therefore, hydrogen cannot form a proton in the presence of cumene. With metallic component, atomic hydrogen migrates to reach Lewis acid site. Unlike molecular hydrogen, atomic hydrogen can compete with cumene molecule for the adsorption site, and forms a proton. It has been revealed that the formation of protonic acid sites from molecular hydrogen is observable for the catalysts other than Pt/SO42"-ZrO2, and the protonic acid sites thus formed act as catalytically active sites for acid-catalyzed reaction. W e propose the concept "molecular hydrogen-originated protonic acid site" as a widely applicable active sites for solid acid catalysts. The concept should be useful when designing solid acid catalysts used in the presence of hydrogen. Hydrocracking catalyst m a y be a typical example to be considered. The "cracking part" in hydrocracking is supposed to proceed on acid sites. It became clear that the formation of protonic acid sites from hydrogen molecules should be taken into account for the elucidation of hydrocracking mechanisms. The concept also suggests the possibility that the concentration of protonic acid sites on the surface is controlled by adjusting the hydrogen pressure and temperature. REFERENCES

1. S. Gnep, M. L. Martin de Armando and M. Guisnet, "Spillover of Adsorbed Species", p. 309, 1983, Elsevier. 2. M. Guisnet, J. Catal., 88 (1984) 249. 3. H.G. Karge, Z. Sarbak, K. Hatada, J. Weitkamp and P. & Jacobs, J. Catal., 82 (1983) 236. 4. H.G. Karge, Y. Wada, J. Weitkamp, P. & Jacobs, J. Catal., 88 (1984) 251. 5. G. Shults-Ekloff, N. I. Jaeger, C, Vladov and L. Petrov, Appl. Catal., 33 (1987) 73. 6. T. Baba and Y. Ono, Zeolites, 7 (1987) 292. 7. T. Baba and Y. Ono, Appl. Catal., 55 (1989) 301. 8. T. Baba, M. Nomura, Y. Ono and Y. Ohno, J. Phys. Chem., 97 (1993) 1288. 9. X. Bai and W. M. H. Sachtler, J. Catal., 129 (1991) 121. 10. S.T. Homeyer, Z. Karpinski and W. M. H. Sachtler, J. Catal., 123 (1990) 60. 11. T.T.T. Wong and W. M. H. Sachtler, J. Catal., 141 (1993) 407. 12. T. Hosoi, T, Shimadzu, S. Ito, S. Baba, H. Takaoka, T. Imai and N. Yokoyama, Prep. Syrup. Div. Petro. Chem. Am. Chem. Soc, 1988, p562. 13. K. Ebitani, J. Konishi and H. Hattori, J. Catal., 130 (1991) 257. 14. K. Ebitani, J. Tsuji, H. Hattori and H. Kita, J. Catal., 135 (1992) 609. 15. K. Ebitani, H. Konno, T. Tanaka and H. Hattori, J. Catal., 135 (1992) 60. 16. K. Ebitani, H. Konno, T. Tanaka and H. Hattori, J. Catal., 143 (1993) 322. 17. K. Ebitani, T. Tanaka and H. Hattori, Appl. Catal. & 102 (1993) 79. 18. M.-G. Yang, I. Nakamura and K. Fujimoto, Appl. Catal. A, 127(1995) 115.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 1 l th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

533

Selective I s o m e r i z a t i o n of A l k a n e s on S u p p o r t e d T u n g s t e n O x i d e A c i d s Enrique Iglesia, David G. Barton, Smart L. Soled
E.

Department of Chemical Engineering, University of California at Berkeley, Berkeley, CA 94720 USA; (t)Exxon Research and Engineering Co., Route 22 East, Annandale, NJ 08801 USA; (2)Departamento de Ingenieria de Procesos, Universidad Autonoma MetropolitanaIztapalapa, A.P. 55-534, 09340 Mexico, D.F., Mexico; (3)Edge Analytical, P.O. Box 2365, Stanford, CA 94309 USA. ABSTRACT Tungsten oxide species form strong acid sites on ZrO2 supports. After calcination at 1000-1100 K and promotion with Pt, these solids catalyze C7+ alkane isomerization at 400-500 K with much higher selectivity than sulfated oxides or zeolitic acids at similar turnover rates. Alkane isomerization proceeds via bimolecular reactions involving hydrogen transfer from alkanes or H2, which cause the desorption of isomeric carbocations before 13-scission occurs. On Pt/SOx-ZrO2, carbocation desorption is slow, leading to long surface residence times and extensive cracking. On Pt/WOx-ZrO2, carbocation desorption is rapid and surface isomerization steps limit n-heptane isomerization turnover rates. Saturation coverage by WOx surface species inhibits ZrO2 sintering and its tetragonal to monoclinic structural transformation. High isomerization turnover rates appear to require the presence of WOx clusters on ZrO2 surfaces. X-ray absorption at the W-LI and W-Lm edges suggests the predominant presence of distorted octahedral species, even after dehydration at 673 K, in all WOx-ZrO2 samples calcined at 1073 K. Tetrahedral species, which lead to a strong pre-edge feature in the W-LI absorption edge, are not detectable in these samples. UV-visible spectra suggest an increase in WOx domain size with increasing loading. These distorted octahedral WOx domains on ZrO2 differ markedly in structure, reduction rates, and alkane isomerization turnover rates and selectivities from tetrahedral WOx species on A1203. 1. INTRODUCTION Acid catalysis is a critical requirement in octane enhancement processes such as isomerization and alkylation, which form highly branched isoparaffins with high octane. Isomerization and alkylation processes in current practice rely on liquid acids (e.g., H2SO4, HF) and on metal oxides promoted by halogens (HCI-A1203). These materials pose significant corrosion, containment, and environmental challenges that can be eliminated by the use of stable metal oxides. As halogen use in chemical processes becomes restricted and ultimately eliminated, solid acids based on metal oxides will provide the only catalytic option to replace current technologies. New solid acids must also improve process efficiency by increasing isoparaffin selectivity and catalytic stability in alkylation and isomerization practice. This work has been in part funded by a grant from the National Science Foundation, #CTS-9510575.

534 Recent work on strong solid acids for isomerization and alkylation has focused on sulfate-promoted metal oxides [ 1]. Catalyst deactivation, sulfur leaching during reaction and regeneration, and high cracking selectivity and low multibranched isomer selectivity in C7+ alkane isomerization reactions currently prevent the industrial use of these materials. Tungsten oxide species on the surface of ZrO2 and TiO2 [2, 3] also form strong and permanent acid sites. The addition of small amounts of a metal component to WOx-ZrO2 leads to n-alkane isomerization [4-6] and ring opening [7] catalysts with unprecedented selectivity and stability. This study examines the structural and mechanistic basis for this unique catalytic behavior. 2. E X P E R I M E N T A L M E T H O D S

Acid sites are formed by impregnation of high surface area zirconium oxyhydroxide, prepared by precipitation methods, with a solution of ammonium metatungstate. ZrOx(OH)4.2x was prepared by hydrolysis of a 0.5 M ZrOC12 solution (Aldrich, 98% ZrOC12"8H20) using dropwise addition of NH4OH (Baker, 28% NH3) until a final pH value of 10. Previous studies have shown that high surface area and pore volume require hydrolysis at high pH [8]. Dried ZrOx(OH)4.Ex was impregnated with a solution of (NH4)6H2W12040 (Strem Chemicals, 99.9%) and calcined in flowing dry air by increasing the temperature at 0.16 K s ~ and holding at the final temperature for 3 h. Final calcination temperatures of 1000-1100 K are required for high isomerization activity at low temperatures (400-500 K) after Pt impregnation. Pt (0.3% wt.) was introduced by incipient wetness impregnation of calcined WOx-ZrO2 samples with a solution of tetraammine Pt hydroxide (Strem Chemicals, 99%), followed by decomposition in dry air at 723 K for 3 h. Pt/SOx-ZrO2 samples were prepared using previously reported procedures [9]. Bulk WO3 powder (Aldrich, 99.995%) was used as received. Surface areas were obtained from N2 physisorption isotherms using standard BET methods. X-ray diffraction spectra were obtained using Cu-Ka radiation (Siemens D5000 diffractometer, 0.002 deg s1 scans). Monoclinic and tetragonal volume fractions were calculated from the integrated intensity of the monoclinic (11 1) and (111) lines and the tetragonal (111) line using a previously reported non-linear function [ 10]. Crystallite diameters were calculated from the corrected integral width of these lines using the Scherrer equation [11]. Temperature-programmed reduction studies were carried out in a flow apparatus using 20% H2/Ar and a post-reactor 13X sieve trap to remove water formed during reduction. Reduction rates were obtained by thermal conductivity measurements of the H2 concentration in the exit stream as the sample temperature was raised from room temperature to 1343 K at 0.167 K s~ and held for 0.5 h. Catalytic studies were carried out using both fixed-bed and recirculating gradientless batch reactors. Catalytic tests were run at 0.1-2.5 MPa total pressure and 443-513 K after catalyst samples (0.025-0.050 mm diameter particles) were treated at 673 K in flowing He for 3 h and then reduced in HE at 473 K for 1 h. HE and D2 (Matheson, UHP) were passed through a catalytic purifier and a 13X sieve trap before use. n-Heptane (Fluka, 99.7%) and n-heptane/adamantane (Aldrich, >99%) mixtures were introduced into a heated zone held at 523 K using a liquid micropump in the fixed-bed studies and by vaporizing from a bulb in the batch reactor tests. Reactants and products were analyzed by capillary chromatography using flame ionization and mass spectrometric detection methods. Reaction rates are reported as turnover rates based on the total number of W-atoms in the catalyst sample. This choice reflects the mechanistic evidence for rate-determining acid-catalyzed steps during alkane isomerization on these catalysts, the structural evidence for dispersed WO• species, and the

535 difficulty in determining accurate acid site densities using conventional site titration techniques. Product selectivities are reported on a carbon basis as the percentage of the carbon atoms in converted n-heptane that appears as a given product. Weight hourly space velocity (WHSV) is reported as (g n-heptane/g catalyst-h). Ultraviolet-visible (UV-vis) diffuse reflectance spectra of supported WOx samples and standard W compounds were obtained with a Varian (Cary 5E) spectrophotometer using polytetrafluoroethylene as a reference. The Kubelka-Munk function was used to convert reflectance measurements into equivalent absorption spectra [12]. Spectral features of surface WOx species were isolated by subtracting from the WOx-ZrO2 spectra that of pure ZrO2 with equivalent tetragonal content. All samples were equilibrated with atmospheric humidity before UV-vis measurements. X-ray absorption spectra were obtained using beamlines IV-3 at Stanford Synchrotron Radiation Laboratory (SSRL) and X-10C at the National Synchrotron Light Source. Si(220) monochromators were detuned to 33% of entering intensity and operated at a resolution of 2 eV at the W-L1 edge. Data were obtained at 0.25 eV increments near the LI edge (12100 eV) and 0.04 A j in the fine structure (EXAFS) region. EXAFS data for the W-LIH were limited to the 10207-10739 eV because of overlap with the LIH- edge of a Hf impurity in ZrO2 (0.5 % wt. Hf). Spectra were obtained using an in-situ cell equipped with Kapton windows held by flanges with Cu gaskets and a sample wafer diluted with rI-Al203 to give 2.4% wt. W. Data were analyzed using previously reported procedures [13] and difference spectra were calculated by subtracting the bulk WO3 powder W-LI spectrum from those of each sample [ 14]. 3. RESULTS AND DISCUSSION 3.1 Alkane Isomerization Rates, Selectivity, and Kinetics The addition of a metal function (0.3% wt. Pt) to WOx-ZrO2 leads to very high isomerization turnover rates and selectivities at very low temperatures (473 K, Figures 1 and 2). These isomerization selectivities are much higher than on Pt/SO• and other acid catalysts. At 473 K, isomerization selectivities at about 50% n-heptane conversion are 85% on Pt/WOx-ZrO2, but only 35% on Pt/SOx-ZrO2 (Figure 1). n-Heptane isomerization turnover rates are similar on Pt/SO• and Pt/WOx-ZrO2 at low temperatures (Figure 2). At higher reaction temperatures, Pt/WOx-ZrO2 catalysts are significantly more active because of their greater stability against deactivation by sublimation or decomposition of ZrO2-supported acid sites. Pt/WOx-ZrO2 catalysts also retain high isomerization selectivities at higher temperatures (71% at 523 K and 36% conversion). A previous study has shown that n-heptane isomerization on Pt/SO• catalysts proceeds by chain transfer pathways, in which carbenium ions propagate by hydrogen transfer from a neutral molecule to a carbocation (Scheme l a) [9]. This process controls the rate at which sites turnover by removing an isomerized carbocation from the surface and replacing it with a new carbocation. Hydrogen transfer rates determine the desorption rate, the surface residence time of isomerized carbocations, and thus the probability that ]3-scission reactions occur before desorption. On Pt/SO• catalysts, carbocation desorption is slow and limited by the rate of hydrogen transfer from H2, alkanes, or hydrogen donors such as adamantane [9]. As a result, the surface lifetime of isomerized carbocations and the probability that they undergo 13-scission is very high. This mechanism (Scheme l a) is consistent with the observed linear dependence of isomerization turnover rate and selectivity on H2 and adamantane concentrations on Pt/SOx-ZrO2 (Table 1, Figure 2) [9]. It suggests that

536 Table 1. Reaction orders in n-heptane and dihydrogen on Pt/WOx-ZrO2 (0.3% wt. Pt, 12.7% wt. W, 1073 K calcination) and Pt/SOx-ZrO2 (0.4% wt. Pt, 4.5% wt. S, 873 K calcination) [473 K, 0.033-0.2 MPa n-heptane, 0.8-2.9 MPa H2]. Catalyst

>1I0

-0.5 (0.2) 1.0 (0.1)

i"

"

t"

8O

-.-

-,.i...

i...,.,

...

Pt/VVOx-ZrO 2

uJ 6o

._I

l.U t4)

Dihydrogen n-Heptane Reaction Reaction Order Order

PffWOx-ZrO2 Pt/SOx-ZrO2

~oo

z 0 I-

40

_~ 2o

0.9 (0.1) 0.2 (0.1)

"

~A

Pt/SOx_ZrO2

n," u,I

o

Note. <15% n-heptane conversion; values in parenthesis correspond to 95% confidence intervals.

o

0

I

I

I

20

40

60

80

n-HEPTANE CONVERSION (%)

Figure 1. Isomerization selectivity on Pt/WOx-ZrO2 (0.3% wt. Pt, 12.7% wt. W, 1073 K calcination) and Pt/SOx-ZrO2 (0.4% wt. Pt, 4.5% wt. S, 723 K calcination) [473 K, 650 kPa H2, 100 kPa n-heptane].

H-atoms formed from H2 or adamantane on Pt sites are involved in carbocation hydrogen transfer steps. The surface isomerization step is quasi-equilibrated, but carbocation formation and desorption by hydrogen transfer steps are irreversible steps. On Pt/SOx-ZrO2 catalysts, the surface is almost saturated with carbocations at steady-state. This leads to the observed weak dependence of isomerization rates on n-heptane concentration (Table 1). I Scheme 1. Isomerization reaction pathways on Pt/SOx-ZrO2 and Pt/WO,-ZrO2 catalysts. ] (a)

r

I Pt/sO*'zrO2 ] H+

(b)

[ PtNVO~-ZrO2]

-

11, -

.,. A A , k,

,),,,+ A

\

---,k+^

On Pt/WOx-ZrO2, n-heptane isomerization turnover rates depend weakly on adamantane (Figure 2) and dihydrogen (Table 1) concentrations. Turnover rates on Pt/WOx-ZrO2 become negative order (-0.5) in HE concentration, in contrast with the linear dependence observed on Pt/SOx-ZrO2 catalysts. Isomerization selectivities are also less sensitive to adamantane concentration on Pt/WO• These data are consistent with faster hydrogen transfer rates, which lead to short surface lifetimes, to low steady-state carbocation coverage, and to desorption of carbocations before [3-scission. On Pt/WO• n-heptane isomerization occurs on a surface sparsely populated with carbocations and surface isomerization is the rate-determining step (Scheme l b). In effect, carbocation formation/desorption by hydrogen transfer steps is quasi-equilibrated at reaction conditions and the carbocation equilibrium coverage is proportional to n-heptane pressure, leading to the observed first order on

537 n-heptane. The negative H2 pressure order (-0.5) arises from competitive occupation of active sites by H § and carbocations, which causes the coverage of the latter to decrease with increasing H2 pressure. Pt/WOx-ZrO2 catalyzes the activation and storage of H-atoms much more efficiently than Pt/SOx-ZrO2 and provides rapid carbocation neutralization and desorption pathways that prevent cracking of adsorbed carbocations. This mechanistic conclusion is confirmed by the rapid isotopic exchange between n-C7H16 and DE observed on Pt/WOx-ZrO2 during n-heptane isomerization (Figure 3). H-D equilibrium between n-C7HI6 and DE is reached very rapidly on Pt/WOx-ZrO2. These data show that n-heptane adsorbs, exchanges with a surface hydrogen pool initially consisting of D-atoms, and desorbs frequently before isomerization occurs. In contrast, adsorption of n-heptane on PIfSOx-ZrO2 is irreversible and several surface isomerization events occur during its residence at the surface (Figure 3). These kinetic measurements and the low n-heptane reaction orders on Pt/SOx-ZrO2 (Table 1) confirm that acid sites on these materials are saturated with carbocations that undergo C-C scission frequently before desorption. 12

80

10 t~

uJ Ir O z tr F---

PtJVVOx-ZrO 2

8

,A

6 4 2

< • .c:_

,at "O--

--0--

--

--

--0"

-"

Pt/WOx_Zr02 I

0.1 ADAMANTANE

--

-0

o

_O--

--

--

-O -

-

-

-

-

t

I

I-Q.

Pt/SOx-ZrO 2

4t

.4.

6O I

4o

I I

2o Pt/SOx-ZrO 2

I 0.2 PRESSURE (kPa)

Figure 2. The effect of adamantane addition on n-heptane isomerization turnover rates [473 K, 650 kPa H2, 100 kPa n-heptane].

L,_,_

0

-

0.17

-

-5-

0.33

-'

0.50

T I M E (h)

Figure 3. Deuterium contents in "unreacted" n-heptane [gradientless batch reactor, 443 K, 100 kPa D2, 4 kPa n-heptane].

Dihydrogen dissociation is required before H2 can act as a source of hydrogen atoms. Pt clusters in SOx-ZrO2 catalysts do not chemisorb H2 at room temperature [9] and appear to be present in the form of oxidized or sulfur-poisoned Pt crystallites, which are less effective in H2 dissociation and H-transfer than Pt metal during alkane isomerization. Reduced Pt sites on Pt/WO• dissociate and store H-atoms required for hydrogen transfer and carbocation desorption. These species are not available on Pt/SOx-ZrO2 and may account for the high isomerization selectivity observed on Pt/WO• catalysts. As a result, Pt clusters on Pt/WOx-ZrO2 catalysts chemisorb significant amounts of H2 at room temperature and catalyze exchange of H-atoms between D2 and "unreacted" n-heptane much more efficiently than Pt clusters on Pt/SO• (Figure 3). 3.2 Surface Areas and Crystal Structures The synthesis of WOx-ZrO2 solids with high isomerization activity requires oxyhydroxide precursors, calcination at 1000-1100 K, and W loadings of 10-12% wt., as also reported

538 earlier by Hino and Arata for metal-free - 100 50 WOx-ZrO2 used in related reactions [1, 2]. / X-ray diffraction and surface area -8o 4o / measurements suggest that these W-atom N / -._I '" / surface densities correspond to saturation -60 z < 30 0 / coverages, which markedly inhibit zirconia (.9 "' sintering and tetragonal to monoclinic -40 ~-LU transformations at high temperatures. ZrO2 I-i0 surface areas after 1073 K calcination are ra 1 0 -20 o~ 4 mEg1 and increase to an asymptotic value 4 5 6 7 W nm 2 I I ! I of 51 m2g"1 for W surface densities above 0 0 5-6 W-atoms nm 2 (Figure 4). Similarly, 4 0 2 4 6 8 10 12 %wt.W WOx-ZrO2 samples with surface densities above 5-6 W-atoms nm 2 contain only the metastable tetragonal phase (Figure 4), but Figure 4. The effects of W loading on BET pure ZrO2 contains only the thermosurface area and crystal modification of the dynamically stable monoclinic phase. This zirconia support (1073 K calcination). apparent saturation coverage is higher than monolayer capacities reported for WO3/A1203 (3.7 W-atoms nm 2) [15]. Particle diameters obtained from X-ray line broadening are similar for monoclinic and tetragonal ZrO2 crystallites; these diameters are consistent with the surface areas measured using BET methods. These data suggest that high isomerization rates occur on samples with near saturation W surface densities. This saturation coverage appears to be assembled during calcination by the sintering of the underlying ZrO2 support. Three-dimensional growth of bulk WOa-like species begins to occur when this saturation coverage is exceeded. W-loadings above 10 W-atoms nm 2 lead to detectable WO3 lines in X-ray diffraction spectra after calcination at 1073 K.

?

f

J

// /

I

I

I

I

I

3.3 Temperature-Programmed Reduction of WOx/ZrO2 All WO• samples reduce in three steps (Figure 5), as previously reported [16]. These steps appear to correspond to reduction to WO29, WO2, and W with increasing temperature. The temperature required for the second and third reduction steps decreases with increasing W loading for WOx-ZrO2 samples (Figure 5a) and approaches the reduction behavior of bulk WO3. A direct comparison with bulk WO3, however, is impaired by marked differences in crystal size between bulk and supported samples. The integrated intensity of the reduction peaks corresponds to full reduction of WO3 to W metal at temperatures below 1350 K for all WOx-ZrO2 samples. WOx-ZrO2 reduces at significantly lower temperatures than WOx/AI203 samples of similar W surface density (Figure 5c). Isolated tetrahedral WOx species in WOx/AI203 have been detected by X-ray absorption [17] and account for their stability against reduction [18]. The reduction of W species in WOx-ZrO2 occurs at temperatures typical of those required for the reduction of hetero-polyoxoanion WOx clusters [ 19]. The introduction of a Pt function influences weakly the behavior of the two hightemperature reduction peaks, but markedly decreases the temperature of the minor lowtemperature reduction step from 700 K to 350 K (Figure 5b). These data suggest that some reduction of WOx-ZrO2 species can occur during n-alkane isomerization reactions (440-500 K). These reducible W species may act as redox sites required for the conversion of H-atoms to H + species on WOx-based solid acids.

539

A possible mechanism for H-atoms from Pt to form acid protons on WOx-ZrO2 is proposed in Eqn. 1, where the net negative charge is distributed over the tungsten oxide cluster in species (II). The accommodation of a proton by electron transfer and charge delocalization across an extended W-O network results in an electronic structure similar to tungsten bronze. The ability to delocalize the net negative charge on the conjugate base may be linked with the strong acidity appearing near the saturation coverage of WOx groups, for which the domain sizes become large enough to distribute the charge over several W atoms. This mechanism is similarto that responsible for acidity generation in heteropolytungstates (H+)8.x(M+XO4)(WO3)12 [19]. This simple mechanism may also account for the ability of H-atoms to desorb carbocations and regenerate neutral WO3 clusters (Eqn. 2).

8 C -5. 4 oE m, ~ -r"5 E v tu

H

+ (W+603)n

(I)

-

2 0 10-

.Ii(.."\'\",, '\ \.

//, _~,....._~. i i

--t

-

. '~,. v ~

J

(b)

8 6

D

~" n

0

~ 10 ~ 8 O ~ 6 T 4

I-

1

I

I

Ai

(c)

0

i

<

F

-I

G

l

2C)0 400 600 800 100012001400 TEMPERATURE (K)

Figure 5. Temperature-programmed reduction profiles of WOx-ZrO2 samples (a): 2.2% wt. W (A), 7.9% wt. W (B), and 21.0% wt. W (C); Pt/WOx-ZrO2 samples (b): 11.6% wt. W (D) and 0.3% wt. Pt, 12.7% wt. W (E); and W standards (c): bulk WO3 (F) and 7.9% wt. W/AI203 (G).

; (H +) (Wn+.61)(W +5) O3n + Pt, (II)

(C7Hl%) (Wn+61)(W .5) O3 n Jr- H- Ptx

'

,~ 4 rr ~3 2

3.4 Ultraviolet-Visible S p e c t r o s c o p y UV-vis spectra obtained for WOx-ZrO2 samples, (NH4)6H2W12040, and bulk WO3 standards are shown in Figure 6. The absorption spectrum for (NH4)6H2WI2040 exhibits two bands with maxima at 4.9 and 4.0 eV. The band at 4.9 eV has been assigned to electron transfer from O to W in O=W species, based on previous studies for MoOx samples [20]. The band at 4.0 eV probably reflects similar processes in W-O-W species. Individual bands are not apparent in bulk WO3 because of its polymeric nature and wide range of W-O-W distances. Pt x -

(a)

10

;

(WO3) ~ + CTH,6(g ) + Pt•

(1)

(2)

All WOx-ZrO2 samples show a maximum near 4.8-4.9 eV. An unresolved band at lower energies appears at high W loadings ( >15% wt. W), suggesting the incipient formation of WO3 crystallites. For surface densities up to 6 W-atoms nm 2, the spectra can be described by a single absorption band of Gaussian shape. This band cannot be unequivocally assigned to a

540 Table 2. Surface density (0w), direct band gap energy (Egd), and number of next nearest neighbors (Nw) on WOx-ZrO2calcined at 1073 K.

,,,,I,,,,I,,,,~E,,, I

12

o.,,'~ m

W loading (% wt. W)

0w Egd(~ Nw(2) (W-atoms nm"2) (eV)

",

~6-

~.

I!

2.2 4.5 7.9 9.3 11.6 21.0 (NH4)6H2WI204o

3.05 3.93 5.01 6.01 7.51 15.0 -

4.15 4.05 4.05 3.85 3.80 3.75 3.45 2.85

1.7 2.0 2.0 2.7 2.8 3.0 4.0 6.0

bulk WO3 o) calculated from F(R~) = k(hv-Egd)/hv (2) b7 line~.extrapolation from standard compounds

0

~ 2

-I I .... 3 4 ENERGY, eV

l'';'l 5

6

Figure 6. UV-vis spectra of WOx-ZrO2 samples: 2.2% wt. W (A), 11.6% wt. W (B), and 21.0% wt. W (C), and W standards: (NH4)sH2WI2040 (D) and bulk WO3 (E) plotted as the Kubelka-Munk function vs. energy.

specific local WOx symmetry [21 ]. Structural assignments based on standard compounds are difficult because of spectral changes caused by interactions of WOx with ZrO2 [22] and by changes in cluster size [21 ]. The low-energy absorption edge in UV-vis spectra is sensitive to domain size in the 1-10 nm range for many semiconductors [23] including WO3 [24]. In this size range, the large fraction of atoms residing at the surface alters the cluster electronic properties. As a result, the energy gap between the valence and conduction bands decreases with increasing domain size. This energy band gap ranges from 2.7 eV for crystalline WO3 [25] to 6.21 eV for the HOMO-LUMO gap in isolated WO42"(aq.) ions [22]. Optical band gap energies (Eg) for WOx-ZrO2 samples calcined at 1073 K were obtained from UV-vis spectra using procedures based on direct and indirect transitions between valence and conduction bands [26]. Direct band gap energies (Egd)decreased monotonically from 4.15 to 3.75 eV as the W loading increased from 3.05 to 15.0 W-atoms nm -2 (Table 2). Recently, Egd values for supported MoOx samples were found to be proportional to the number of next nearest neighbors (Nw) [27]. Using Egd measurements for WO3 (N~ = 6) and (Nn4)6HEW12040 (Nw = 4), Nw values for WOx-ZrO2 can be obtained by linear extrapolation. The number of next nearest neighbors appears to increase monotonically with increasing W-loading (Table 2). Clearly, these Nw values reflect variations in the orbitals involved in the electronic transitions with changes in loading and they are not to be taken as purely geometrical quantities [27]. Nonetheless, it appears that isolated WOx species are not present in any WOx-ZrO2 samples after 1073 K calcination and that all samples contain significant concentrations of polymeric -W-O-W- groups. The size of WOx domains appears to increase with increasing W surface density.

3.5 X-Ray Absorption Studies W-LI absorption edges for WO3 and for several WOx-ZrO2 samples calcined at 1073 K are shown in Figure 7. The absorption edge reflects the excitation of W 2s electrons by X-ray photons. A pre-edge feature at -5 eV is caused by 2s to 5d dipole-forbidden transitions, which become detectable in non-centrosymmetric compounds because of d-p orbital mixing [28].

541 Therefore, this feature is sensitive to the symmetry of the absorber [29] and it has been used to probe the symmetry of W atoms supported on A1203 and TiO2 [ 17, 30]. This pre-edge feature is not present in perfect octahedral structures, but appears in distorted octahedra and with higher intensity in tetrahedral structures. The similarity among all the spectra in Figure 7, shown also by the difference spectra, suggests similar distorted octahedral symmetry for WOx species in all WO• samples. In-situ dehydration of these samples at 673 K in flowing He for 1 h has no detectable influence on the X-ray absorption spectra. Therefore, the distorted octahedral symmetry of WOx-ZrO2 samples is not caused by hydration of tetrahedral WOx species during brief exposure to ambient air. The radial structure functions corresponding to W-LI edges are very similar in all WO• samples and qualitatively resemble those for bulk WO3 out to distances of about 6/~,. Coordination number calculations for the WO3 system are not reliable because of the low symmetry and multiple bond distances in most stable WO3 phases [31]. These data, however, show that the bonding symmetry around W atoms is not strongly influenced by loading or by the ZrO2 support for samples calcined at 1073 K, even though such differences in loading lead to significant differences in catalytic behavior. Previous studies have concluded that 4-, 5-, and 6-coordinate W species are present on A1203 and TiO2 supports [ 17, 30] depending on surface W density and on hydration state. The present study has detected WO3-1ike distorted octahedral domains at all surface densities and irrespective of hydration on ZrO2. These species catalyze alkane isomerization reactions with much higher turnover rate and selectivity than dispersed WOx moieties on alumina or titania. These results are consistent with active sites consisting of highly distorted octahedral WOx clusters on ZrO2. Acid sites formed by these octahedral WOx surface species are more effective isomerization sites than previously reported tetrahedral WOx species on A1203 [ 17], possibly because of the ability of WOx clusters to form metastable proton-containing complexes during catalytic isomerization reactions.

w 0 z< rn n," 0 if? m < a u.l N ,(

!

w _J I1. <

n,' 0 Z --40

-20

~ 0

20

(E - Eo) / eV

40

60

I'-40

i -20

i 0

i 20

I 40

60

( E - Eo) / eV

Figure 7. Normalized W-Li edge absorbances from WOx-ZrO2 samples" 2.2 % wt. W (A), 11.6% wt. W (B), 0.3% wt. Pt, 12.7% wt. W (C), and 21.0% wt. W (D), and bulk WO3 (E). Difference spectra at right, with equivalent vertical scale, calculated by subtracting the WO3 spectra from sample spectra.

542 References

[1] Arata, K., Adv. Catal. 37 (1990) 165; Tanabe, K., Crit. Rev. Surf. Chem. 1 (1990) 1. [2] Hino, M. and Arata, K., J. Chem. Soc. Chem. Commun. (1987) 1259. [3] Arata, K. and Hino, M., Proc. 9th Intern. Cong. Catal. ("Oxide Catalysts and Catalyst Development", Phillips, M.J. and Teman, M., eds.). The Chemical Institute of Canada, Ontario (1988) 1727. [4] Iglesia, E., Barton, D.G., Soled, S.L., Miseo, S., Baumgartner, J.E. and Gates, W.E., 1995 Spring Meeting American Chemical Society (1995); Proc. 14th North American Meeting Catal. Soc. (1995). [5] Soled, S.L., Miseo, S., Baumgartner, J.E., Gates, W.E., Barton, D.G., and Iglesia, E., Proc. 13th Intern. Conf. Catal. ("New Trends in Solid Superacids and Superbases", Izumi, Y., Anpo, M., and Izumi, Y., eds.). The Taniguchi Foundation (1994) 17. [6] Soled, S.L., Gates, W.E., and Iglesia, E., U.S. Patent 5,422,327 (1995). [7] Chang, C.D., Santiesteban, J.G., and Stem, D.L., U.S. Patent 5,345,026 (1994). [8] Gimblett, F.G.R., Rahman, A.A., and Sing, K.S.W., J. Colloid Interf. Sci. 84 (1981) 337. [9] Iglesia, E., Soled, S.L., and Kramer, G.M., J. Catal. 144 (1993) 238. [ 10] Toraya, H., Yoshimura, Y., and Somiya, S., J. Am. Ceram. Soc. 67 (1984) C 119. [ 11 ] Klug, H.P. and Alexander, L.E., "X-Ray Diffraction Procedures". Wiley and Sons, New York (1974). [ 12] KortOm, G., "Reflectance Spectroscopy- Principles, Methods, and Applications". Springer-Verlag, New-York (1969). [13] Sayers, D.E., Lyttle, F.W., and Stem, E.A., Phys. Rev. Lett. 27 (1971) 1204. [ 14] Meitzner, G.D. and Sinfelt, J.H., Catal. Lett. 30 (1995) 1. [ 15] Salvatti, L., Makovsky, L.E., Stencel, J.M., Brown, F.R., and Hercules, D.M., J. Phys. Chem. 85 (1981) 3700. [16] Vermaire, D.C. and van Berge, P.C., J. Catal. 116 (1989) 309. [17] Horsley, J.A., Wachs, I.E., Brown, J.M., Via G.H., and Hardcastle, F.D., J. Phys. Chem. 91 (1987) 4014. [18] Soled, S.L., McVicker, G.B., Murrel, L.L., Sherman, L.G., Dispenziere, N.C., Hsu, S.L., and Waldman, D., J. Catal. 111 (1988) 286. [ 19] Pope, M.T., "Heteropoly and Isopoly Oxometallates". Springer-Verlag, Berlin (1983). [20] Bartecki, A. and Dembicka, D., J. Inorg. Nucl. Chem. 29 (1967) 2907. [21 ] Foumier, M., Louis, C., Che, M., Chaquin, P., and Masure, D., J. Catal. 119 (1989) 400. [22] Iannibello, A., Marengo, S., Tittarelli, P., Morelli, G., and Zecchina, A., J. Chem. Soc. Faraday Trans. I 80 (1984) 2209. [23] Brus, L., J. Phys. Chem. 90 (1986) 2555. [24] Ozin, G.A. and ()zkar, S., J. Phys. Chem. 94 (1990) 7556. [25] Kopp, L., Harmon, B.N., and Liu, S.H., Solid State Commun. 22 (1977) 677. [26] Butler, M.A., J. Appl. Phys. 48 (1977) 1916. [27] Weber, R.S., J. Catal. 151 (1995) 470. [28] Shadle, S.E., Hedman, H., Hodgson, K.O., and Solomon, E.I., Inorg. Chem. 33 (1994) 4235. [29] Li, P. and Chen, I.W., J. Am. Ceram. Soc. 77 (1994) 118. [30] Hilbrig, F., Gobel, H.E., Kn6zinger, H., Schmelz, H., and Lengeler, B., J. Phys. Chem. 95 (1991)6973. [31] Loopstra, B.O. and Boldrini, P., Acta Cryst. 21 (1966) 158.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

543

Tungsta and Platinum-Tungsta Supported on Zirconia Catalysts for Alkane Isomerization Gustavo Larsen,* Edgar Lotero and Rub6n D. Parra Department of Chemical Engineering, University of Nebraska-Lincoln, Lincoln, NE 685880126 USA Tungsta- and platinum-tungsta supported on zirconia catalysts (WZ and PtWZ) were tested for isomerization of n-butane with hydrogen at 573 K. The catalysts' reduction kinetics and the chemical state of tungsten were studied by X-ray absorption near-edge spectroscopy (XANES) and conventional temperature-programmed (TPR) techniques. Platinum was found to catalyze the reduction of the support and the tungsta co-catalyst. An effective n-butane isomerization catalyst results from the combination of high-temperature calcination and lowtemperature reduction. Pyridine adsorption, as monitored by diffuse reflectance infrared spectroscopy (DRIFTS), indicates that water of reduction generates Br0nsted acid sites. 1. I N T R O D U C T I O N Oxoanion-promoted zirconia catalysts are active systems for acid-catalyzed reactions such as alkane isomerization and alkylation. The long-term performance of such materials is strongly affected by a number of fouling processes, such as coking, loss and reducibility of oxoanion and loss of support surface area. For example, in spite of the fact that noble metals and the presence of hydrogen or good hydride transfer agents are known to improve the stability of sulfated zirconia [ 1], partial loss of sulfate ions is known to occur during catalyst reduction [2]. Tungstated zirconia is also a good candidate [3], and the stability of this system is also improved by the presence of Pt [4]. This paper deals with the characterization of WO3/ZrO2 (WZ) and Pt/WO3/ZrO2 (PtWZ) catalysts for alkane isomerization, as monitored by diffuse reflectance infrared spectroscopy of adsorbed pyridine, X-ray absorption near-edge structure of W L edges, and temperature-programmed techniques. 1.1 Previous work and motivation In two recent papers [8,10], we have initiated studies aimed at understanding the catalytic behavior of WZ and PtWZ. Our observations, which motivated the present study, can be summarized as follows, a) Water of reduction results in the formation of Brcnsted acid sites, as monitored by pulsed addition of pyridine to a DRIFTS chamber at room temperature [8,10]. In this paper, we have complemented those results with similar pyridine adsorption experiments at

* To whom all correspondence should be addressed

544 different temperatures, keeping in mind that these materials decompose the base molecule above 723 K. b) A negative reaction order in hydrogen and a positive one in n-butane were observed for the isomerization pathway at two different metal dispersions, whereas variable and positive reaction orders in both hydrogen and n-butane were measured for the formation of hydrogenolysis products, c) Coke deposition was higher on Pt-free and low dispersion catalysts, but the selectivity toward isomerization also increased with decreasing catalyst dispersion. Additional reaction data is presented in this contribution, d) X-ray diffraction showed that the presence of tungsta does not prevent the tetragonal-to-monoclinic zirconia transition if samples are treated at the high calcination temperatures required to induce strong acidity (- 1096 K) [3]. A separate tungsten oxide phase is present after catalyst calcination, e) An acid catalyzed process, the temperature-programmed reaction of methanol, was used to study the differences between alumina- and zirconia-supported catalysts. The latter yields diethyl-ether with high selectivities over the decomposition to CO+H2 pathway and is also more reactive than the ahmfina-supported material. Given that WZ behaves as a strong solid acid using different probe reactions, we wished to a) determine whether reduction of tungsta occurred in the presence of Pt and hydrogen, in an effort to understand what the chemical state of W is expected to be after catalyst pretreatment and prior to reaction with n-butane and to b) study the effect of temperature on the adsorption of pyridine to gain some insight into the acid-base features of these materials. Reaction measurements, coupled with X-ray absorption near-edge (XANES) and diffuse reflectance infrared (DRIFTS) spectroscopy and temperature-programmed reduction (TPR) are used as chemical and physical probes.

1.2 Experimental

Upon calcination at 1096 K, WZ becomes an active n-butane isomerization catalyst and has BET surface areas 5-10 times higher than tungsta-free zirconia [8,10]. If the Pt function is loaded a posteriori (via Pt(acac)2/ether impregnation) instead of using aqueous co-impregnation techniques, the Pt dispersion can be controlled and catalysts with improved activity for the nbutane/H2 reactions are obtained. However, an increase in Pt dispersion also resulted in an increase in the hydrogenolysis selectivity, as typically observed in other Pt catalysts [ 10]. Low and medium dispersion samples (as monitored by carbon monoxide adsorption) were prepared following the protocols described in reference 10. In brief, a solution of ZrO2C12.xH20 (Aldrich) was used to obtain hydrous zirconia (HZ) upon precipitation with 30 % NH3 (Aldrich) to pH 10. The precipitate was filtered and washed with deionized water until it became Cl--free (AgNO3 test). The solids were dried in an oven at 373 K for 3 hours. The incipient wetness technique was used to obtain a 12 wt% W material, using ammonium metatungstate as tungsten precursor. The impregnated solid was subsequently oven-dried at 373 K for 3 hours. At this point, one third of this material was impregnated with an aqueous solution of PtC16H2.xH20 from Aldrich (38-40 wt% Pt) to yield a 1.2 wt% Pt sample. This Pt- and tungstate-loaded sample was calcined at 1096 K for 1 hour in a quartz U-tube under an air flow of 1 liter/min g. The calcination temperature was achieved by applying a 13.3 K/rain ramp. The sample was labeled as PtWZ (Std) and stored until use. The remaining dried portion of the tungstate-impregnated solid was calcined using exactly the same protocol as in PtWZ (Std). The calcined sample was divided in two equal portions. One of them was labeled as WZ (Pt-free sample) and stored until use. The remaining portion was impregnated several times with an ether solution of platinum acetylacetonate from Strem Chemicals until the 1.2 Pt wt% target loading was achieved. This sample was labeled as PtWZ (acac) and stored in a closed vial. Finally, a 7-A1203 (A) support from Strem Chemicals was subjected to the same aqueous Pt loading and calcination series as in

545

PtWZ (Std) to obtain a reference catalyst. Prior to reaction, all samples were calcined for one hour at the chosen temperature (1096 K for WZ, Pt/A1203 and Pt/WZ (Std), and 773 K for Pt/WZ (acac)) and subsequently reduced for 1 h under a H2 flow of .5 liter/min g at 623 K. The choice of temperature for the second calcination cycle of the Pt/WZ (acac) sample is not in any way arbitrary. The idea is to use a non-aqueous scheme to keep the incorporation of moisture to a minimum, and a calcination temperature low enough to guarantee better metal dispersions. Pyridine was adsorbed in situ in a DRIFrS chamber by means of pulse addition of the base molecule until saturation was achieved, followed by purging of the cell for 0.5 h with ultrahigh purity nitrogen. Given the estimated saturation loadings and the Pt wt%, we expect the DRIFTS spectra of pyridine to be essentially unaffected by the presence of noble metal. Prior to pyridine adsorption, samples were calcined and reduced in situ as described above. The excess hydrogen after reduction was purged with nitrogen at 623 K for 0.5 h. The DRIFTS setup is described elsewhere [2,11 ]. Preliminary results on these catalysts suggested that fragmentation of pyridine starts to occur above 773 K. Thus, we kept our desorption temperatures below that value to minimize the cracking reactions of the base molecule with the surface. We are aware of the fact that some degree of fragmentation may occur at lower temperatures without necessarily having low-molecular weight fragments leave the surface of the catalysts. Nevertheless, the constancy of the number of bands due to pyridine may suggest that this effect may not be of significance. For the TPR experiments, a Residual Gas Analyzer (MKS) quadrupole mass spectrometer was connected to a quartz cell and a stainless steel reactor manifold equipped with mass-flow and temperature controllers. Several masses (m = 18, 2, 32, 4, etc.) were followed during reduction of samples that had been calcined in situ. Calibration of MS signals was done using pulsed addition of standard amounts of H20 and H2 at a temperature were both hydrogen consumption and water evolution had ceased. Typically, 0.25 g of catalyst were crushed and sieved (20-40 mesh fraction), placed in a quartz reactor and subjected to a He/H2 flowing mixture and a temperature ramp of 6.6 K/min. A similar setup coupled with gas chromatographic detection was used to carry out kinetic measurements. For kinetic studies, 0.1 g catalyst was used and the partial pressure of n-butane and hydrogen were adjusted to 51.7 and 16.9 kPa respectively, using nitrogen as balance gas to 101.3 kPa. Conversions were kept below 2 %. We have also conducted in situ XANES experiments at the X18b beamline of the National Synchrotron Light Source (NSLS). The W L edges were investigated using a stainless steel XAS cell similar to that previously described [ 12]. Since our cell did not allow for the use of temperatures above 773 K, we proceeded to calcine our samples at the required temperature in our laboratory, and to reactivate them (mainly to eliminate moisture) at 773 K under an air flow in the XAS cell. XANES data was recorded under air, and once cooled to room temperature, the air was eliminated switching the flow to pure helium. The next step was to ramp the temperature up to 723 K under a hydrogen flow, to check whether the L edges of tungsten were affected by reduction. Apart from multiple scattering, the LII and LIII edges are primarily affected by the density of emtpy d states, since they represent the allowed pl/2 __> d and p3/2 ._> d promotions of core p electrons respectively [12]. Hence, the area of these so-called white lines give a qualitative indication of the reducibility of the WO3 phase under different gas environments. On the other hand, the LI is a dipole-forbidden transition (s --> d) in a strictly octahedral field, and distortions to the latter cause a pre-edge peak or shoulder to appear immediately after the threshold of photoelectron ejection. This has been used as a fingerprint technique in tungsta supported on alumina catalysts [ 13,14].

546 2. R E S U L T S AND D I S C U S S I O N Figure 1 shows the time on stream (TOS) behavior of PtWZ(acac) in the atmospheric-pressure packed bed reactor operated at 573 K. We have reported that once an activity plateau is reached (typically after lh TOS), the reference Pt/AI203 catalyst, which results in a nearly identical CO uptake to that of PtWZ(acac), yields hydrogenolysis products only in comparable quantities to those observed in PtWZ(acac) [10]. Interestingly, the low dispersion PtWZ(Std) yields isobutane only, but at a rate which is about three times lower than that of PtWZ(acac). In turn, the Pt-free sample (WZ) also yields isobutane only at the same reactor operation conditions, but at a rate which is about one order of magnitude lower than that of PtWZ(acac) [ I0]. All these tungsta-zirconia samples had nearly identical BET surface areas. This led us to suggest that, under the current conditions, the hydrogenolysis activity is primarily metal-catalyzed, whereas the isomerization activity is mainly due to the tungsta-zirconia function. The n-butane conversion to hydrogenolysis products decreases much more rapidly than the isobutane production.

0.25

0.2 v~

al

0.15 O t..

= O

"~

0.1

O.O5 ~ +

% I

10

1

20

I

30

T i m e on

1

I

40

50

l

60

I

70

80

Stream (min)

Figure 1. Time on Stream behavior of PtWZ(acac), (e) isobutane, (21) methane, (O) ethane and (+) propane production rates.

We note that in a purely Pt-catalyzed hydrogenolysis-isomerization competition, we expect the hydrogenolysis pathway to be less demanding in terms of surface configuration [7]. This is also in agreement with our idea that the WZ support alone may be responsible for the isomerization reaction, aided by the metal function by both creation of acid sites by hydrogen spillover and coke prevention as generally suggested for Pt/SO4=/ZrO2 . Figures 2.a-c show the pyridine adsorption results. Br0nsted acidity is manifested by the bands at 1440-1445, 1630-1640 and 1530-1550 cm -1. Bands at 1600-1630 cm -! are assigned to pyridine bonded to Lewis acid sites. Certain bands such as the 1440-1460 and 1480-1490 cm-1 can be due to hydrogen-bonded, protonated or Lewis-coordinated pyridine species. Under continuous nitrogen purging, spectra labeled as "A" in Figures 2a-c represent saturation of the surface at room temperature (904_-_25~tmol pyridine/g found in all three tungsta catalysts) and "F" show the baseline due to the dry catalyst. We cannot entirely rule out the possibility of some extent of weakly bound pyridine at room temperature. Nevertheless, the pyridine DRIFTS experiments show the presence of BrCnsted acidity, which is expected to be the result of water of reduction that did not desorb upon purging at the reduction temperature. It is noted that, regardless of the presence of Pt, the intensity of the DRIFTS signals due to pyridine are

547

markedly suppressed upon desorption above 623 K. Experiments with H-Y zeolite (not shown) indicated that much higher temperatures are required to desorb pyridine completely. Sachtler and co-workers [5] have suggested that reasons other than acid strength are reponsible for the high n-butane isornerization activity of sulfated zirconia (SZ). These authors concluded that the acidity of SZ is comparable to that of protonated zeolites. Since the acidity of WZ is expected to be lower than that of SZ, the lower pyridine retention of WZ with respect to that of H-Y should not be surprising. The relative importance of the 1530-1550 cm -1 (Br~nsted) band is more pronounced on

C

F f 1600

1500 Wavenumbers (cm-1)

Pt/WZ (acac).

A

1600

1500 Wavenurnbers (cm-1)

b

1600

1500 Wavenurnbers (cm-1)

c

Figure 2. Pyridine adsorption at A) room temperature, B) 423, C) 523, D) 623 and E) 723 K. F) corresponds to the dry catalyst at room temperature, a) WZ, b) PtWZ(Std) and c) PtWZ(acac) In order to investigate the state of tungsten after reduction, we began by carrying out TPR measurements on different samples with and without Pt or W. Figure 3 shows the TPR profile of PtWZ(Std). Although the reduction process appears to start at fairly low temperatures, as evidenced by the consumption of hydrogen below 400 K (and desorption of a small amount of water at 20-30 K higher temperature), it is clear that further reduction occurs above 800 K. The amount of water produced (or hydrogen consumed) largely exceeded the one we would predict if Pt were the only reducing species. During calcination, we noticed that temperatures above 900 K were required to completely eliminate chlorine (as HCI) from PtWZ(Std). As expected in materials which can retain water at high temperatures such as promoted zirconias, the water signal is slightly retarded with respect to that due to hydrogen consumption (m=2).

548

9

f

9

1

m=2

; -

\

PtWZ ( S t d ) ~ ~ = 1 8

,

,

t

I

,

l

,

i

<--Piiateau:

109fi

K-->

400 600 800 1000 Figure 3. TPR on PtWZ(Std), hydrogen and water signals.

Z(Std) WZ

.,d

z ,

i

,

I

,

l

,

!

<--Plateau:, 10~}6 K-->,

400 600 800 1000 Figure 4. Water evolution during hydrogen TPR on all samples.

549 Figure 4 presents the water evolution signals for all samples investigated using the TPR technique. While the W- and Pt-free material undergoes very little reduction, WZ is reduced above 1000 K in an amount that is roughly consistent with the transformation of all W 6§ to W 4+. If we assume that the low-temperature signal was due to Pt4+ reduction to Pt 2+, reduction peaks above 650 K should be due to complete reduction of Pt and Pt-catalyzed W 6+ and Zr 4+ reduction. Note that the W-free zirconia support also experiences significant reduction in the presence of Pt. Co-calcination, whether with or without W, seems to induce the most dramatic effect in support reducibility. We noticed that the reduction temperature employed prior to reaction was 623 K, and we suspected (see Figure 4, PtWZ(acac) reduction profile) that tungsta reduction, at least in part, could have been catalyzed by the presence of Pt. As an example, in the XAS in situ experiments we found a change in the W LIII white line intensity of PtWZ(acac) which was beyond experimental error, as shown in Figure 5. Even though quantification is difficult because XANES spectra are a convolution of multiple scattering and bound states, it was a priori expected that (also judging from TPR data) this phenomenon could occur. In addition, if we accept the conventional m e c h a n i s m of BrOnsted acid site generation by hydrogen spillover from Pt sites, as proposed for Pt-SO4=/ZrO2, we must recognize that this would be in fact a reduction process (zerovalent hydrogen y i e l d i n g - O H groups). PtWZ(acac) has a higher dispersion, presumably due to the lower calcination (reactivation) temperature employed after the Pt loading step. In agreement with this picture, note that the contribution of protonated pyridine bands after reduction are highest in PtWZ(acac). The three-times higher isomerization activity, and the lower coke deposition rates [10], could well be due to the relatively higher ability of PtWZ(acac) to give spiltover hydrogen. The compromise lies in the a p p e a r e n c e of hydrogenolysis in PtWZ(acac). Given the modest change in the W L III white line, we had reason to believe that the bulk properties of the WO3 were hardly affected. In fact, the monitoring of the W LI edges confirmed this idea. The two extreme cases, the fully oxidized WZ and PtWZ(acac) after reduction at 723 K, are compared to three different reference materials.

4 ~0 qJ I

qJ N

olm,I m

0

Z

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

r~ I

10200 10220 Energy (eV) Figure 5. PtWZ(acac) under hydrogen at 337 (dashed) and 692 K (solid).

550 The series of deviation from a strictly W 6+ octahedral crystal field follows the sequence ammonium tungstate > tungstic acid > WO3 [14]. The s-->d transition becomes allowed as distortion of the octahedral environment increases. Viewed as a fingerprint technique, Figure 6 shows that the bulk of W present in both WZ (calcined) and PtWZ(acac, reduced) behaves as WO3, as also shown by our XRD data [10], sugesting that only a small extent of W reduction is expected to have taken place during catalyst conditioning prior to reaction. At this point, we should mention that we were unable to carry out Pt XAS experiments because of the low wt% content of this element with respect to the optimum wafer thickness dictated by the strong Zr and W background absorbers and the strong overlapping of W L edges EXAFS signals with the Pt L absorption lines. In addition, it should be understood that it is extremely difficult to perform a two-shell EXAFS modeling of data that i n v o l v e s contributions from backseatters so close in the periodic table such as W and Pt. We have come across this difficulty in the past when trying to model the EXAFS signals of Re and Pt in PtRe bimetallics [15]. In the future, we will employ surface-sensitive techniques such as SIMS to determine the extent of mixed oxide formation in these catalysts.

1.2 O!

e~o

e' Oat

"~ 1.0 r,,rd

I

4

Ot !

!

Ot

0.8 Od

N 0.6 otllm t=m,,=l

E o 0.4 Z

~J

0.2 1

,

!

,

!

1209O 12110 Energy (eV) Figure 6. LI edges of W, (O) (NH4)2WO3. (11) WO3, (O) H2WO3. (+) PtWZ(acac)reduced and (A) WZ-oxidized.

The present characterization studies have motivated us to investigate in the future the optimum calcination and reduction temperatures to maximize the isomerization pathway, while keeping coke deposition and hydrogenolysis to a minimum. Our results suggest that compromises are expected to be made to achieve those goals.

551 A CKNO WLEDGEMENTS This research has been funded by the Donors of the Petroleum Research Fund (ACS, grant # 28336-G5) and by the National Science Foundation (CTS-940618). Support from the Layman Fund and the Research Council at the University of Nebraska is also acknowledged. We thank the National Synchrotron Light Source and the Stanford Synchrotron Radiation Laboratory for X-ray beamtime.

REFERENCES 1. E. Iglesia, S. L. Soled and G. M. Kramer, J. Catal., 144 (1993) 238. 2. G. Larsen, E. Lotero, R.D. Parra, L.M. Petkovic, H.S. Silva and S. Raghavan, Appl. Catal. A: Gral, 130 (1995) 213. 3. M. Hino and K. Arata, J. Chem. Soc. Chem. Commun., (1987) 1259. 4. S.L. Soled, AIChE 1994 Annual Meeting, paper 85f, San Francisco, Nov. 1994. 5. V. Adeeva, J. W. de Haan, J. J/mchen, G. D. Lei, V. Schiinemann, L. J. M. van de Ven, W. M. H. Sachtler and R. A. van Santen, J. Catal., 151 (1995) 364. 6. C. Tsz-Keung, J.L. d'Itri and B.C. Gates, J. Catal., 153 (1995) 344. 7. G. Larsen and G.L. Hailer, Catal. Today, 15 (1992) 431. 8. G. Larsen, S. N. Raghavan, M. M~quez and E. Lotero, Catal. Lett., in press. 9. F. Hilbrig, H.E. G6bel, H. Kn6zinger, H. Schmeltz, H., and B. Lengeler, J. Phys. Chem. 95 ( 1991) 6973. 10. G. Larsen, E. Lotero, S. N. Raghavan, Rub6n D. Parra and C. A. Querini, Appl. Catal. A" Gral, in press. 11. G. Larsen, Edgar Lotero, Manuel M~quez and Hugo S. Silva, J. Catal., in press. 12. B. Qi, I. Perez, P. H. Ansari, F. Lu and M. Croft, Phys. Rev. B, 36 (1987) 2972. 13. F. Hilbrig, H. E. G6bel, H. Kn6zinger, H. Schmeltz and B. J. Lengeler, J. Phys. Chem., 95 (1991) 6973. 14. J. A. Horsley, I. E. Wachs, J. M. Brown, G. H. Via and F. D. Hardcastle, J. Phys. Chem., 91 (1987) 4014. 15. C. G. Michel, W. E. Bambrick, R. H. Ebel, G. Larsen and G. L. Haller, J. Catal., 154 (1995) 222.

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j.w. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

553

n-Butane isomerization on Ni-prornoted sulfated zirconia catalysts W. E. Alvarez a, H. Liua, E. A. Garcia b, E. H. Rueda b, A. J. Rouco b, and D. E. Resasco ~ "School of Chemical Engineering, University of Oklahoma, 100 E. Boyd St., Norman, OK 73019 b Universidad Nacional del Sur, (8000) Bahia Blanca, Argentina

1. INTRODUCTION A few years ago, Hsu et al. [ 1,2] found that Fe and Mn oxides were able to promote sulfated zirconia making it active for the isomerization of n-butane at temperatures much lower than those typically needed on unpromoted sulfated zirconia [3]. Even though the interpretation of this promotion is still a matter of discussion, the idea that the presence of the transition metal oxides promotes an inter-molecular reaction path [4] is gaining strong support [5]. This mechanism [6] consists of a series of steps including the formation of a C4* carbenium ion and a butene molecule which undergo oligomerization and form a surface C8* ion. This C8~ ion isomerizes and splits by fl-scission, forming an isobutane molecule and a C4~ carbenium ion. Our own results on Fe, Mn-promoted catalysts (FMSZ) are consistent with this interpretation [7]. In this contribution, we have investigated the promoting effects of adding Ni to sulfated zirconia catalysts. Previous reports on FMSZ catalysts have indicated that, in the absence of added H2, the isomerization activity exhibited a typical pattern when measured as a function of time on stream [8, 9]. In all cases, the initial activity was very low, but as the reaction proceeded, the conversion slowly increased, reached a maximum, and then started to decrease. In a recent paper [7], we described the time evolution in terms of a simple mathematical model that includes induction and deactivation periods. This model predicts the existence of two types of sites with different reactivity and stability. One type of site was responsible for most of the activity observed during the first few minutes on stream, but it rapidly deactivated. For the second type of site, both, the induction and deactivation processes, were significantly slower. We proposed that the observed induction periods were due to the formation and accumulation of reaction intermediates that participate in the inter-molecular step described above. Here, we present new evidence to support this hypothesis for the particular case of Ni-promoted catalysts. We have recently reported that the addition of Ni results in a promotion of the isomerization activity of sulfated zirconia [10] comparable to that obtained by the addition of Fe and Mn. It has been previously observed that the presence of H2 causes a decrease in isomerization activity, a result consistent with the mechanism that involves olefins as reaction intermediates. Here, we

554 analyze the promoting effects of Ni with and without addition of H2, the activity induction, the catalyst deactivation, and its regeneration in air. 2. EXPERIMENTAL Several Ni-promoted sulfated zirconia catalysts were prepared and investigated. The twostep method of preparation was used [ 11], depositing the Ni promoter and the sulfate over a zirconium hydroxide precursor prior to calcination. The zirconium hydroxide was precipitated from 0.4 M zirconium oxychloride solution by dropwise addition of NH4OH until pH = 10 at room temperature. The solid was then filtered, thoroughly washed with dionized water, and dried in air at 110~ The resulting Zr(On)4 had a BET surface area of 278 m2/g, a pore volume of 0.28 cm3/g, and an average pore size of 4 nm. These physical properties were measured in a Micromeritics ASAP 2010 adsorption apparatus. The addition of Ni was done by impregnation with Ni(NO3)2 solution of various concentrations, keeping a liquid-to-solid ratio of 0.65-0.68 ml/g.. The catalysts were then dried overnight at room temperature and for 8 h at 110~ For the double-promoted catalyst, the addition of Ni and Mn was accomplished by coimpregnation with aqueous solutions containing the nitrates of both metals. The addition of sulfate was done by two alternative methods. The first method was the standard incipient wetness impregnation [12], using aqueous solutions of (Nt'h)SO4 with the proper concentration to achieve the desired sulfate loadings. The second method of sulfate addition was the so-called soaking procedure. After drying for 8 h at 110~ the solid was immersed for 2 h in a well stirred 1M (NI-~)2SO4 solution, using a liquid/solid ratio of 7.5 cm3 per gram of Zr(OI-I)4. Following this soaking step, the solid was allowed to settle and separated from the liquid. After a drying step as that described above, the catalysts were finally calcined at 600~ in air for 13 hours and stored in a desiccator. For comparison, a series of silicasupported catalysts was also prepared using SiO2 Davison Grade 62 as a support. These catalysts were prepared by the same sequential impregnation method as that used for zirconia. Nickel and sulfur contents were determined by Atomic Absorption Spectroscopy (AAS) and by a LECO SC-132 respectively. Some samples were sent for analysis to Galbraith Lab Inc. XRay Diffraction measurements (XRD) were made in a Phillips PW 1710 diffractometer with Cu K radiation, scanned at a rate of 0.035 degrees (2 theta) /sec. X-Ray Photoelectron Spectroscopy (XPS) measurements were made in a Perkin Elmer Physical Electronics 548 spectrometer. Binding energies were referenced to the oxygen 2p line. Temperature programmed desorption and oxidation (TPD and TPO) were conducted in a flow system directly connected to a quadrupole MKS mass spectrometer that allowed for continuous monitoring of up to 16 different masses simultaneously. During TPD experiments gaseous samples were extracted at various temperatures and analyzed in a GCMS to identify all the products being evolved. Temperature programmed reduction (TPR) measurements were conducted in a flow system equipped with a thermal conductivity detector to determine H2 consumption. n-Butane isomerization was studied in a flow reactor. In each run, 0 4 g of catalyst was placed in a stainless steel downward-flow packed bed reactor. The reaction system allows for

555 precise control o f flow rates and temperature. The feed was a mixture of n-butane (99.9 % purity from Matheson) and He (99.999 % purity from Sooner Air Gas). Prior to each reaction run, the catalysts were activated m situ by flowing dry air at 600~ for 2 h. They were cooled under air to the reaction temperature, flushed in a He stream, and finally exposed to the feed mixture. The n-C~-I~0 flow rate was kept constant at 5.46 cm3/min for all runs, it was diluted with a 10.6 cm3/min flow of He or H2. The reaction products were analyzed by a HP-5890 gas chromatograph, equipped with a 50 m KCI/AI20 3 capillary column and a flame ionization detector. The high selectivity exhibited by these catalysts (> 92 % isobutane) allowed us to conduct some experiments in which we only measured the concentration of isobutane and unconverted n-butane. In this way, we were able to obtain data points every minute.

3. RESULTS AND DISCUSSION 3.1. Catalyst Characterization The Ni and S contents on the catalyst series were determined after calcination at 600~ As shown in Table 1, sulfate was only retained on the silica support when Ni was present. Infrared studies have previously shown that sulfate groups impregnated on pure silica are thermally unstable [13]. Therefore, the SOdNi molar ratios, close to unity, together with the colors resulting after calcining the silica-supported samples made us conclude that Ni was in the form of NiSO4. On zirconia, the SOdNi ratios were larger than one because the sulfate can be associated with both, Ni and the support. Table 1: Concentrations of Ni and S in the various catalysts investigated Catalyst

SO4 wt % nominal

Ni wt % nominal

SO4 wt % after calcin,

Ni wt % after calcin,

SO4/Ni molar ratio

Color after calcination

S6Si 1NiSi 3NiSi 1Ni6SSi 3Ni6SSi 6SZ 1NiZ 3NiZ 1NilSZ 1Ni3SZ 1Ni6SZ 3Ni6SZ

6 0 0 6 6 6 0 0 1 3 6 6 soak soak soak soak

0 1 3 1 3 0 1 3 1 1 1 3 -

0.35 0 0 1.54 4.16 4.23 0 0 0.75 1.96 4.15 5.0 NA 8.5 (2.31") 10.8 8.76

0 0.85 2.51 0.97 2.85 0 0.84 2.35 0.99 0.95 0.95 2.21 0.0 0.8 1.2 1.6

0 0 0.97 0.89 0 0 0.46 1.26 2.67 1.38 -

white dark gray dark gray yellow yellow white light gray light gray pink pink pink pink white white/pink white/pink white/pink

SZ(s) 1NiSZ(s) 2NiSZ(s) 3NiSZ.(~s)

556 The sulfate concentration of sample 1NiSZ(s) indicated in parenthesis corresponds to a sample that was treated in He at 600~ The catalysts prepared by the soaking procedure resulted in a higher concentration of sulfate than those prepared by sulfate impregnation. A series of samples prepared by impregnation with higher concentrations of sulfate, i.e., 8 and 12 wt. %, showed the same concentration after calcination as those prepared with lower concentrations. It appears that, on the impregnated catalysts, there is a fixed maximum amount of sulfate (about 4-5 wt. %) that the material is able to retain through the high temperature calcination. By contrast, the soaking procedure appears to significantly increase the amount of stable sulfate. At the same time, the soaking procedure results in significant losses of metals by leaching during the treatment with (NI-~)2SO4 solution. X-ray diffraction (XRD) analysis of the freshly calcined catalysts as well as samples used for several hours in the isomerization reaction, only presented the peaks corresponding to the tetragonal form of zirconia. At the same time, for the silica series, XRD confirmed the presence of NiO on the unsulfated catalysts and NiSO4 on the sulfated ones. However, XRD did not show any evidence of any of these species for the zirconia series, probably due to their high state of dispersion. Similarly, the XPS data clearly showed the presence of NiO and NiSO+ on the unsulfated and sulfated silica-supported catalysts, respectively, but they were not conclusive in the case of zirconia series since both sulfate and oxide species were observed. NiSO4/ZrO2 (550C) NiSO4/ZrO2 (250C) 3Ni6SSi 3Ni6SZ

1Ni6SSi

1Ni6SZ 3NiZ

3NiSi

1NiSi 6SSi 6SZ -~

0

200 400 Temperature (*C)

600

0

.

o

200 400 Temperature (~

600

Figure 1. Temperature programmed reduction of silica- and zirconia-supported catalysts containing nickel and sulfate, a) silica-series; b) zirconia-series. The TPR profiles of the zirconia and silica series are shown in Fig. 1. The silica series showed that the sulfated Ni catalysts are more difficult to reduce than the unsulfated ones. The position of the maxima shifted 30-50~ to higher temperatures as sulfate was added. As opposed to the silica series (see sample 6SSi), the zirconia series exhibited Hz consumption

557 peaks even in the absence of Ni (see sample 6SZ), indicating the reaction of H2 with sulfate. When Ni was added to the zirconia series, the H2 consumption started at much lower temperatures suggesting that, while the presence of sulfate retarded the reduction of Ni, the presence of Ni accelerated the reduction of sulfate. The samples containing sulfate and Ni (3Ni6SZ and 1Ni6SZ) exhibited double peaks. The first of these peaks (about 400~ can be ascribed to the reduction of Ni and sulfate very close to it. The second peak (about 470~ is due to the reduction of sulfate that is not in close contact with the Ni. To demonstrate this, we have impregnated NiSO4 on pre-calcined ZrO2. This sample was further treated in air at 250~ or 550~ The one pretreated at 250~ presented a H2 consumption peak at about the same temperature (400~ as that of the first of the double peaks of the NiSZ catalysts, indicating that these peaks are indeed related to the presence of NiSO4 or, at least, sulfate in close proximity to Ni. The sample pretreated at 550~ showed the decrease of the peak at 400~ and the appearance of one at about 450~ This splitting could be due to the decomposition of NiSO4, which would leave sulfate groups on the zirconia support, separated from Ni, and therefore, more difficult to reduce. 35

20

30 15

25 ~, 20

t-

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o

L)

10 5

I ,

0

10

i

L

,'

20

30

40

50

Time (min) Figure 2.

Conversion

- time curves on

several Ni - promoted catalysts under nC,I-I~0/He mixture, n-Butane conversion to isobutane as a function of time on stream over 0.4 g cat., at 150~ n-C,l-I~o flow rate = 5.46 cm3/min, He flow rate = 10.4 cm3/min. 3NiSZ(s) (squares); 2NiSZ(s) (triangles); 1NiSZ(s) (circles).

0

10

20

30

40

Time (rain)

Figure 3. Conversion -time curves on regenerated catalysts at the same conditions as those in Fig. 2. 1NiSZ(s), first run after regeneration (solid triangles), second run (open triangles); NiMnSZ(s) (open squres); SZ(s) (solid circles); 1NiSZ(s) treated in He at 600~ (crosses).

558 3.2. C a t a l y t i c activity

In previous work [7], we reported that the soaking procedure resulted in FMSZ catalysts that were more active than those prepared by adding sulfate using the incipient wetness impregnation method. We report here the activity of catalysts prepared by the soaking procedure. Fig. 2 shows the variation of activity as a function of time on stream for three different NiSZ(s) catalysts.The typical evolution of activity previously reported for FMSZ catalysts was again observed on these Ni-promoted catalysts. The relatively rapid induction was followed by two deactivation regimes, a rapid deactivation during the first 10 min and a slower deactivation thereat~er. The long term activity was directly related to the Ni concentration, while the initial increase in activity showed no correlation with the amount of Ni. At the same time, it varied very much from sample to sample, and even among different runs on the same sample. In all the runs, the selectivity towards isobutane was very high ranging 92-97 %. In agreement with the results previously obtained on the FMSZ catalysts [7], the conversion - time curve obtained on the fresh catalyst was clearly different from those on the subsequently

regenerated samples. On the regenerated catalysts, the initial induction and deactivation (first 10 min.) were significantly slower than on the fresh catalysts. As opposed to the first runs for which the initial increase in activity varied from run to run, the reproducibility of the activity of the regenerated catalysts was excellent. The behavior exhibited by regenerated catalysts is illustrated in Fig. 3, in which the activity pattern of two different runs on a sample of 1NiSZ(s) is compared to those of the unpromoted SZ and a catalyst promoted with Ni and Mn (0.8 wt% Ni, 0.3 wt% Mn). A significantly higher activity was observed for the double-promoted catalyst. This rate enhancement was also evident during the first run, although it was more pronounced atter the regeneration. The activity of the 1NiSZ(s) catalyst after a pretreatment in He at 600~ is also included in Fig. 3. This treatment rendered the catalyst almost totally inactive. In contrast with all the other pretreatments that did not significantly modify the color of the samples, the heating in He, caused the sample to turn black. This color change can be ascribed to the decomposition of the nickel sulfate species, which occurs with significant sulfate losses. As shown in Table 1, the sulfate content on this sample decreased from 8.5 to 2.3 after the treatment in He, but this sulfate loss did not cause any change in the crystal structure of zirconia. X R analysis showed that zirconia was still in the tetragonal form. To determine whether the induction period was caused by the accumulation of adsorbed species or by an irreversible modification of the catalyst structure we conducted the following experiment. As mentioned above, the regenerated catalysts had a much higher reproducibility in the shape of the conversion - time curves than the fresh catalysts. Then, for this experiment, we used a sample that had been previously used in a reaction period of several hours and then regenerated at 600~ We started the flow of reactants under the typical reaction conditions at 150~ over the 2NiSZ(s) catalyst. After reaching the maximum conversion, we stopped the flow of n-Cd-I10 leaving the catalyst under He at the reaction temperature, and then resumed the flow of n-Cd-I~0. We repeated this sequence several times and monitored the variation of conversion. As shown in Fig. 4, every time that we stopped the flow of n-Cd-I~0 and then resumed it, a new induction period was observed. However, each subsequent induction resulted in a lower maximum conversion. The envelope of these maxima is indicated in the figure and it is almost the same curve as that obtained in a regular run. The induction periods significantly

559 lengthened after each subsequent cycle. For example, it took only 1.5 min to go from zero to 6% conversion for the first induction period, but it needed 8 min to reach the same conversion during the second period, and 20 min. during the third. It appears that, as the deactivation proceeds, the induction rate decreases. This result supports the idea that the induction is related to the accumulation of reaction intermediates, produced during the first stages of the reaction. As the catalyst deactivates, the accumulation of intermediates becomes slower. 16

30

14

12

..

=20

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o~j0 C 0

...

:3 ..Q O

..

.m_

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o

..........

O ,4-,, C

t 9

. ....

.9 ~-10 tO

o

c.c

30

60 90 Time (min)

120

150

Figure 4. Conversion-time curves on a regenerated 2NiSZ(s) sample obtained at intermitent reaction periods after which the flow of n-butane was stopped, leaving the catalyst at the reaction temperature under a 100 cm/min flow of He for a few minutes, and then resuming the flow of n-butane. The arrows indicate 15 ktl injections of l-butene.

=.5-_ -

100

200

300

400

Temperature (*C) Figure 5. n-Butane conversion to isobutane as a function of temperature in a temperature programmed reaction experiment conducted over 0.4 g of 1NiSZ(s) catalyst under an n-butane/ hydrogen mixture (n-C4 molar fraction = 0.34) at a constant heating rate of 2C/min.

Fig. 4 also shows the result of injecting 1-butene in the feed during one of the induction periods. We previously demonstrated [7] that on FMSZ catalysts, the addition of olefins significantly modified the form of the induction period. When 1-butene was added to the He stream before starting the feed of n-C,I-I~0, the induction period was very much shortened and the activity rapidly increased. By contrast, when the addition of olefins was done several minutes after contacting the catalyst to the n-CA-I~0feed, there was almost no difference in the shape of the conversion-time curve. Similar results were obtained on the NiSZ(s) catalyst reported here. When 1-butene was added before starting the feed of n-Cd-I~0 (not shown)the induction period was very much shortened. The arrows in Fig. 4 indicate the injection of 1butene, both at the beginning and almost at the end of the fourth induction period. The

560 activity jump caused by the injection was more pronounced at the beginning, when the surface started to accumulate olefins, than at the end, when it was almost at its saturation point. In a separate experiment, we varied the length of time during which the catalyst was left under He before resuming the butane flow. When this time was 10 rain or longer, we observed that upon resuming the n-Cd-I~0 flow the conversion started from zero, as shown in Fig. 4. By contrast, when the catalyst was left in He for 2 rain or less, the conversion started at about the level that was before shutting the n-Cd-I~0 flow off. This result indicates that the surface species can be removed during the 10 min. periods, but remain on the surface when the period without n-CA-It0was short. Compared to the runs with n-Cd-I~0/He mixtures, those conducted n-C4H10/H2 mixtures exhibited a much lower activity and did not show the rapid induction period exhibited by the fresh catalysts under the n-Cd-I10/He mixtures.

3.3. Temperature Programmed Studies In addition to the isothermal runs, we carried out a temperature programmed reaction (TPRx) experiment to study the evolution of n-Cd-I~0 isomerization activity of the NiSZ catalysts as a function of temperature in the presence of H2. In this experiment, we increased the temperature linearly at a heating rate of 2~ and simultaneously monitored the evolution of products, sampling the reactor outlet every 5 min. The variation of conversion of n-C4H~0 to isobutane in the linear temperature ramp is illustrated in Fig. 5. It can be observed that the conversion initially increased with temperature, reached a maximum at about 250~ and then rapidly decreased. It is interesting that at about 250~ when deactivation begun, the concentration of Ct, C3, and C5 products started to increase, but, when the temperature reached 400~ the concentration of all products dropped to zero. There are several possible causes for the rapid deactivation observed above 250~ One of them is the formation of carbonaceous deposits. This form of deactivation, however, should not be substantial in the presence of H2. Other possible causes of catalyst deactivation are the loss of sulfate groups and the conversion of Ni sulfate into Ni sulfide. In our TPR experiments, we detected the evolution of H2S, but this only occurred at temperatures close to 500~ not at the relatively low temperatures at which deactivation started. As described above, we attempt to explain the observed catalytic behavior in terms of the accumulation of reaction intermediates which involve the participation of olefins [4, 7, 10]. To study the interaction of C4 olefins with the Ni-promoted and the unpromoted sulfated zirconia catalysts, we exposed two samples (1NiSZ(s) and SZ(s) previously treated in air at 600~ to a 10 Tort of 1-butene during 10 rain at room temperature. The catalyst was then exposed to pure He and the temperature was increased with time at a constant heating rate of 8~ The desorption products, continuously monitored by MS, started to appear at about 100~ and had a maximum at about 210~ To identify the products desorbed at the begining of the desorption process and at the point of maximum resorption rate, we sent pulses to a GCMS at 150~ and 210~ The total amount of products detected in the second pulse (maximum desorption rate) was more than 20 times greater than in the first pulse. The product distribution significantly varied for the two pulses. For instance, the 150~ exhibited the presence of butenes, and C5 olefins as main products. This distribution indicates that some of the adsorbed butenes undergo oligomerization

561 and cracking, while others may have just desorbed after a simple double-bond isomerization. The 210~ exhibited a much higher alkane/olefin and the presence of isobutane as the main desorption product. The concentration of isopentane and hexane was also significant. The presence of these products would indicate that, in addition to the oligomerization and cracking steps, a H2 transfer process takes place in which the coke deposits provides the H2 for hydrogenating the olefins. The differences observed in the product distribution and desorption temperatures following the adsorption of 1-butene for the promoted and unpromoted catalysts are relatively minor. This suggests that the important effects of Ni are most evident under reaction conditions. Although olefins do play a role [7] the promoting effect of Ni cannot be ascribed to a direct interaction between the promoter and the olefin but perhaps to a concerted effect in which more than one species is participating. To probe the formation of coke we conducted TPO measurements on samples previously used in the butene TPD experiments. The TPO profiles corresponding to catalyst 1NiSZ(s) are shown in Fig. 6. Significant evolution of CO2 was detected, indicating the formation of coke during the adsorption/desorption of 1-butene. The 1NiSZ(s) catalyst exhibited almost twice as much CO2 as the unpromoted SZ sample.

to'J .

m

O')

t.j t~ CO t~ Ill

0

200 400 T e m p e r a t u r e (C)

600

Figure 6. TPO of carbonaceous deposits lett on the surface of I NiSZ(s) after the adsorption/desorption of 1-butene. Evolution of CO2 and SO2 (curves A and B, respectively). Evolution of SO2 during the butene TPD (curve C, added for comparison). The loss of sulfate during the reaction steps or during regeneration may become a critical issue when analyzing the potential of these materials as commercial catalysts. Sulfate losses during the butene TPD, made evident by the evolution of SO2 (rn/e=64), started to occur at about 500~ We have previously demonstrated the evolution of SO2 in the presence of adsorbates such as ammonia, benzene, or pyridine at temperatures much lower than those required to produce SO2 from clean sulfated zirconia [ 14]. For instance, A treatment in He at 600~ causes drastic losses which result in a significant drop in activity (see Fig. 3). It is

562 important then to compare the evolution of SO2 under different conditions. For example, during the TPO experiment, very small quantities of SO2 were evolved contrasting with relatively large amounts evolved during the TPD of the adsorbed olefins. 4. CONCLUSIONS Several conclusions can be drawn from the present work: 9 The promoting effect on the n-C~8-I~0isomerization increases with the amount of Ni, and further increases with the addition of Mn. We have presented evidence that on the Nipromoted sulfated zirconia catalysts, the nickel is, at least partially, in the form of sulfate. 9 The induction period exhibited by these catalysts is related to the accumulation of reaction intermediates on the surface. 9 When hydrogen is present in the feed, the catalytic activity of NiSZ is much lower than when it is not present. Under hydrogen, the activity increases with temperature up to 250~ but it rapidly deactivates at higher temperatures. 9 Sulfate losses are important above 500~ under reducing conditions and in the presence of adsorbed olefins. They are much less significant under oxidizing conditions such as those found during the catalyst regeneration in air. ACKNOWLEDGMENTS We thank Ms. Wei-Chee Tan for the preparation of the catalyst samples. This work was supported by the National Science Foundation (CTS-9403199), the international cooperative program NSF-CONICET (INT-9415590), and the Exxon Education Foundation. We thank the University of Mar del Plata for a fellowship (WEA), as part of the international exchange program sponsored by the University of Oklahoma and the University of Mar del Plata. REFERENCES l.

2.

.

9. 10. 11. 12. 13. 14.

US Pat. 4,918,041 (1990) C. Y. Hsu, C. R. Heimbush, C. T. Armes, and B. C. Gates J. Chem. Soc. Chem. Comm. (1992) 1645 M. Hino and K. Arata, J. Chem. Soc., Chem. Comm. (1980) 573 Adeeva V., Lei G. D., and Sachtler W. M. H., Appl. Catal. A 118, L11 (1994) A. S. Zarkalis, C. -Y. Hsu, and B. C. Gates, Catal. Lett. 29, 23 5 (1994) Guisnet M. R., Acc. Chem. Res. 23,392 (1990) M. A. Coelho, W. E. Alvarez, E. C. Sikabwe, R. L. White, and D. E. Resasco, Catal. Today (in press) A. Jatia, C. Chang, J. D. MacLeod, T. Okabe, and M. E. Davis, Catal. Lett. 35, 21 (1994) J. Tabora and R. J. Davis, J. Chem. Soc. Faraday Trans. 91, 1825 (1995) M. Coelho, D. E. Resasco, E. C. Sikabwe, and R. L. White, Catal. Lett. 32, 253 (1995) F. R. Chen, G. Coudurier, J. F. Joly, and J. C. Vedrine, J. Catal. 143, 616 (1993) R. A. Comelli, C. R. Vera, and J. M. Parera, J. Catal. 151, 96 (1994) B. A. Morrow, R. A. McFarlane, M. Lion, J. C. Lavalley, J. Catal. 107, 232 (1987) E. C. Sikabwe, M. A. Coelho, D. E. Resasco, and R. L. White, Catal. Lett. 34, 23 (1995)

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

563

RARE EARTH MODIFIED SILICA-ALUMINAS AS SUPPORTS FOR BIFUNCTIONAL CATALYSIS Stuart L. Soled*, Gary McVicker, Sal Miseo, William Gates, and Joe Baumgartner, Exxon Research and Engineering Co., Rt. 22 East, Annandale, NJ 08801 USA ABSTRACT We have explored rare earth oxide-modified amorphous silica-aluminas as "permanent" intermediate strength acids used as supports for bifunctional catalysts. The addition of well dispersed weakly basic rare earth oxides "titrates" the stronger acid sites of amorphous silicaalumina and lowers the acid strength to the level shown by halided aluminas. Physical and chemical probes, as well as model olefin and paraffin isomerization reactions show that acid strength can be adjusted close to that of chlorided and fluorided aluminas. Metal activity is inhibited relative to halided alumina catalysts, which limits the direct metal-catalyzed dehydrocyclization reactions during paraffin reforming but does not interfere with hydroisomerization reactions. INTRODUCTION The worldwide interest in solid acid catalysis that has developed over the last few years has focused primarily on oxide-based strong "super" acids as possible replacements for conventional liquid or halide-containing acids (1-3). Additional opportunities exist in processes that use intermediate strength acids. In particular, hydrocarbon conversions that require the simultaneous participation of both a metal hydrogenation-dehydrogenation function and an acid function, such as catalytic reforming and paraffin hydroisomerization, have traditionally used bifunctional catalysts consisting of group VIII metals dispersed on chlorided or fluorided alumina supports. In some applications, halide is stripped off these catalysts (by water generated from the reaction, for example) and periodic halide replacement by treatment with organic halides or HCI is required. Availability and use of halogen-containing treatment gases, howver, are coming under increasing environmental pressure. In addition, variation in halide level often produces unwanted changes in acid strength that result in selectivity changes during reaction. Consequently, there is interest in developing alternative environmentally-compatible halide-free bifunctional catalysts (4). Silica-aluminas contain acid sites stronger than those found in most halide-treated aluminas. We have attempted to tailor the acidity of amorphous silica-alumina by adding varying levels of "permanent" inorganic basic titrants. Such titrants were chosen in order to meet the following four criteria: 1) be weakly basic- in order to gradually titrate the acid sites so that their strength matches that found in halided-aluminas. This precludes the use of alkali and alkaline earth oxides. 2) be easily dispersible- the weakly basic oxides should spread as a monolayer on the SiO 2A120 3 surface. If the oxides agglomerate, the acid sites would not be uniformly titrated. 3) be non-reducible- because of the metal present on bifunctional catalysts, we want to avoid any possible alloy formation on the support surface during reduction. 4) be able to disperse group VIII metals- it has been known that it is difficult to maintain high metal dispersions on pure SIO2-A120 3. To obtain an effective metal/acid balance, enhanced

564 dispersion of the group VIII metal on modified SIO2-A120 3 relative to neat SIO2-A120 3 would be advantageous. We have found that rare earth oxides are mild bases that spread uniformly on SIO2-A120 3 supports and systematically reduce the acid strength into the range of halided aluminas. Of course, solid acidity is a complex parameter that encompasses acid site number, type, strength, hardness etc. so that matching the entire spectrum of acidic properties of a particular support is not likely. In fact, subtle differences in acid and metal properties provide opportunities to exploit differences in catalytic performance. Rare earth oxides do not reduce during catalytic reactions of interest, and we show, surprisingly, that they provide advantages in improving Pt metal dispersion. Consequently, we evaluated Pt on rare-earth oxide-modified silica-aluminas as environmentally compatible bifunctional catalysts. Acidity has to be closely balanced with an active and dispersed group VIII metal, usually Pt, to minimize either metal or acid cracking side reactions (5). Both physical and chemical characterization probes as well as selected model compound reaction tests were used to characterize these new catalysts. Acid site strengths have been determined using an olefin isomerization test, metal dispersions estimated with dihydrogen chemisorption and benzene hydrogenation, and bifunctional properties monitored during heptane reforming and n-C12 isomerization. We have characterized both acid and metal functions and compared them with conventional bifunctional catalysts. EXPERIMENTAL Preparation and Characterization The SiO2-AI20 3 (Si-AI) powder samples were obtained from Davison; both MS25 (75% wt. S i t 2, 25% wt. A1203) and MS13 (87% wt. S i t 2, 13% wt. A1203) were evaluated. The rare earths were added by incipient wetness impregnation of nitrate solutions, dried overnight at 120~ and then air calcined to oxides at 500-600~ We included yttrium oxide in our studies, since it behaves similarly to rare earth oxides. Group VIII metals were added by incipient wetness impregnation. For those samples containing Pt as the group VIII metal, we used a standardized chloroplatinic acid solution as the platinum precuror. X-ray diffraction spectra, collected on a Rigaku D-Max diffractometer with CuK~

radiation, was used to check for crystalline phases. ESCA spectra, which monitored the rare earth oxide dispersion, were measured on a Perkin Elmer 5600 XPS hemispherical electron analyzer using Mg Kt~ (1256 ev) radiation with a pass energy of 35ev. Samples were mounted on double sided tape. Peak positions were shifted to place carbon at 285 ev. Surface area and pore size distributions were measured by N 2 gas desorption. Platinum dispersions were measured by hydrogen chemisorption at 25~ assuming a 1:1 H:Pt surface stoichiometry. Chemisorption isotherms were obtained by measuring dihydrogen uptakes between 25 and 200 torr and extrapolated to zero pressure to obtain the total chemisorption uptakes. The strong chemisorption uptakes, which are reported here, were obtained by subtracting the back adsorption isotherm (i.e. the hydrogen chemisorption measured after evacuating the cell following the total adsorption isotherm) from the total uptake (6). Reaction tests The 2-methylpent-2-ene (2MP2) isomerization test was carried out as described previously (7). The formation rates and rate ratios of the product hexene isomers of this test reaction reflect the relative acid site concentration and strength of the catalyst,

565 respectively. The product hexene isomers formed include 4-methylpent-2-ene (4MP2), t-3methylpent-2-ene (t-3MP2), and 2,3 dimethylbute-2-ene (2,3 DMB2). 4MP2 requires only a double bond shift, a reaction occurring on weak acid sites. 3MP2 requires a methyl group shift (i.e., a stronger acidity requirement than double bond shift), whereas the double branched 2,3DMB2 product requires even stronger acidity. For a homologous series of solid acids, differences in t-3MP2 rates normalized with respect to surface area reflect the density of acid sites possessing strengths sufficient to catalyze skeletal isomerization. Since skeletal isomerization rates generally increase with increasing acid strength, the ratio of methyl group migration rate to double bond shift rate should increase with increasing acid strength. The use of rate ratios, in lieu of individual conversion rates is preferable since differences in acid site populations are normalized. For the n-C 7 reforming and n-C 12 isomerization reactions the catalysts were run in a fixed bed micro reactor equipped with on-line GC analysis. The catalyst, together with a quartz powder diluent, was added to a 6 inch reactor bed. A thermocouple was inserted into the center of the bed. The catalysts were calcined at 350-500~ immediately prior to use and reduced in H 2 at 350-500~ for 1 hour. n-Heptane or dodecane (Fluka, puriss grade) were introduced via a liquid feed pump. The runs were made at 100-175 psi with a H2/n-heptane (or n-C12) feed ratio of 7 and a weight hourly space velocity of 6-11. Benzene hydrogenation was used to probe metal site activity. A 12/1 H2/benzene feed was passed over the catalysts at 700 kPa with a weight hourly space velocity of 25. The temperature was set to 100~ and the conversion of benzene to cyclohexane was measured after 2 hours at temperature. The temperature was then increased at 10~ increments and after two hours, the conversion remeasured. R E S U L T S AND D I S C U S S I O N Samples with different loadings of rare earth and yttrium oxides on MS25 Si-AI were prepared and characterized. Figure 1 shows the x-ray diffraction spectra of a sample of 25% Nd203/Si-AI compared with the unmodified Si-AI.

0

IO

20 two

30

40

50

~ea~m(des,~)

Figure 1. X-ray diffraction spectra of Si-A1 and 25% Nd/Si-AI. In both cases, we observe an amorphous pattern; no crystallites of rare earth oxide appear even at 25% wt. loading. This indicates that oxide particles remain less than -30,~, in diameter. The surface area, pore volume and pore size distribution of the starting Si-AI support also change on impregnation. Table 1 lists the values for yttria-modified samples of

566 MS25. The results show that as the yttria disperses over the entire surface, the smaller pores of Si-AI are blocked first making a fraction of the surface inaccessible to physisorbed N 2. Furthermore, the average pore size increases as a result of preferential blockage of small pores (Figure 2). Table 1. Change of Surface Area and Pore Volume For Yttria-modified Si-A1 (75/25% wt.) surface area (m2/g) 302 235 208 183

Si-A1 4%Y203/Si-AI 9%Y203/Si-AI 15%Y203/Si-A1.

,e~ r.,,n,~ ,~ :~,Jo-lis~vl' ~ ,,,..---.I ~ L'p*,, ,,t,=~ 0.~-0.~,,sgj

20 16

;i02-AI, Oj t "~

~12

pore volume (cc/g) 0.78 0.67 0.65 0.52

, , ~ ' ~

4.Y.O~SIO.-AI.O.

4 sio,.~o;' "f...... /_ I

o

-

0

10

-

20

30

40

SO r,

60

?t

$0

90

""

100

(~)

Figure 2. Pore size distribution for Y203-modified Si-AI (75/25% wt. SIO2-A1203) We examined a series of Nd203-1oaded silica-aluminas by ESCA to further monitor rare earth oxide dispersion. As seen in Figure 3, the ESCA Nd/(Si+A1) signal, which measures the surface concentration of Nd (relative to (Si and AI), to about a -30,& penetration depth) increases linearly with increasing Nd20 3 loading. This result strongly suggests that N d 2 0 3 continues to wet and spread on the surface (i.e. remain below monolayer coverage) up to 30% wt. loading. 0.25

o 0.2i o 0.15 + 0.1-

0.0S 10

15

20

2$

30

wt~ NdzO~ on SiO,-AI203

Fig. 3. Surface Nd ESCA Signal Increases Linearly with Loading The titration of silica-alumina acid sites with neodymium oxide is clearly seen in Figure 4. The relative acid strength, as estimated by the 2,3DMB2/4MP2 ratio during 2MP2

567 isomerization, systematically decreases with increasing N d 2 0 3 loading. The ratio does not change dramatically for the lowest neodymia loadings under these conditions (250~ 1 hr run time) since it lies near its equilibrium value. Under less severe, lower temperature isomerization

Fig. 4. SiO2-A1203 Acid Strength Decreases with Nd203 Addition conditions, larger differences occur at low rare-earth oxide loading levels. Comparative data for chlorided and fluorided aluminas is also shown in Figure 4 and indicate that the acid strength of rare earth modified silica-aluminas can be adjusted to match that of conventional halide-modified transitional aluminas. We investigated variations in thermal treatments on rare-earth modified Si-AI as well as using different rare earth oxides. We are including Y203 with the rare earth oxides, since it behaves similar to them. In Figure 5, we show that 25% wt. Nd203/Si-A1 and 18% wt. Y203/Si A1 which contain comparable similar atomic loadings have nearly identical t-3MP2/4MP2 rate ratios and both ratios are just slightly larger than that shown by 0.9% wc C1/AI20 3.

Fig. 5. Equimolar Rare Earth Oxide Concentration Display Similar Acid Strengths An interesting difference occurs with the stronger acid sites. Figure 5 shows that 0.9% wc CI/A120 3 contains more of the stronger acid sites, as measured by the 2,3DMB2/4MP2 ratio than the 25% N d 2 0 3 or 18% Y 2 0 3 catalysts. We suspect that this difference in distribution in site strengths results from preferential titration of stronger acid sites on Si-AI by the basic rare earth oxides. This subtle difference in acid strength distribution between halided aluminas and rare-earth oxide modified silica-aluminas should have an impact on their catalytic properties. We also see (Fig. 6) that raising the calcination temperature of rare earth modified silicaaluminas from 500 to 600~ may slightly sinter the rare earth oxide. Although the x-ray pattern

568 still remains amorphous, sintering would lower the effective surface coverage of the neodymium oxide titrant and consequently increases support acidity. This speculation is consistent with the data shown in Fig. 6.

Fig. 6. Apparent Acid Strength of 25% Nd203/SiO2-AI20 3 Increases with Calcination Temperature Published studies have shown that differences in basicity among rare-earth oxides is small; significantly less than differences among alkaline earth oxides (8). Furthermore, basicity does not change uniformly across the periodic table. A sequence of base strengths for the different rare earth oxides has been proposed: La > Pr = Nd > Sm > Gd -- Eu > Tb -- Ho --- Er > Dy --- Tm = Yb -- Lu > Ce (9). As seen in Figure 7, aside from CeO 2, most of the rare earth oxides we tested did not show substantial differences in basicity in the tests used here. Only impregnation with ceria did not appreciably titrate the acidity, perhaps as a result of the lower basicity of the Ce +4 ion.

Fig. 7. Different Rare Earth Oxides at Equimolar Loadings on SIO2-A120 3 We measured the dispersion of Pt (impregnated from a chloroplatinic acid precursor, calcined at 450~ and reduced at 500~ on a series of Nd203-1oaded silica-aluminas (Fig. 8). We find, unexpectedly, that dispersion increases with increasing rare earth oxide loading up to about 18% N d 2 0 3, where it plateaus at between 40 and 50%, compared to 10% with unmodified Si-AI. This compares with dispersions o f - 6 0 - 8 0 % measured on similarly Pt-loaded transitional A1203 catalysts. Transmission electron micrographs confirmed the decrease in particle size with rare earth content on Si-A1.

569

0.8

0.57

0.6

~

OA

i

0.2

5

10

15

20

25

30

% Nd203 In 0.3%Pt/Nd~O~/Si-A!

Fig. 8 Pt dispersion on Nd203-modified SIO2-A120 3, 0.3% wt. Pt, catalysts reduced 450~

As an additional probe of metal activity, we monitored benzene hydrogenation activity. As seen in Figure 9, Pt-containing rare earth catalysts have lower hydrogenation activity than chlorided alumina catalysts; this result reflects inhibition of metal activity on these supports relative to conventional transitional alumina supports. Whereas the acid strength can be adjusted close to that of chlorided and flourided aluminas, metal activity is somewhat inhibited on these catalysts relative to halided aluminas. This inhibition is not due to dispersion, and perhaps indicates a SMSI interaction between Pt and the dispersed Nd20 3 phase. 100 ., 80~0.3%Pt/0.9%CI/AI

60 ~9 40

"

)r

i

.3%Pt/25%Nd20~Si-A1

~ 20 ~-

: 7%Nd2Ob/Si.Ai

100

120

140

160 ('C)

180

200

Temperature

Fig. 9. Benzene hydrogenation on Nd203-modified SIO2-A120 3 (100~ kPa; 25 WHSV).

12/1 H2/C6H 6, 700

We compared Pt-containing rare-earth Si-AI catalysts with 0.3% Pt/0.9% C1/AI203 under n-heptane reforming conditions. The direct metal catalyzed dehydrocyclization of n-C 7 produces toluene, a desirable product. The formation of methyl hexanes proceeds primarily through a bifunctional mechanism. Both metal and acid sites can crack heptane or methyl hexanes with CH 4, C2H 6, and n-C4H10 originating from metal hydrogenolysis activity, while i-C4H10 arises principally from acid catalyzed beta-scission reactions. Small quantities of propane can form from

570 either metal or acid catalyzed reactions. As seen in Table 2, total conversion under identical conditions is about two times higher over 0.3% Pt/0.9%C1/AI20 3 than over 0.3% Pt/25%Nd203/Si-AI catalysts. Table 2. Heptane reforming at 175 psi, 6/1 H ~/n-C 7, 500~ 11 WHSV, 10 h on stream. toluene C5methyl hexane total selectivity (%) selectivity (%) selectivity (%) conversion

(%) 0.3%Pt/0.9%Ci)A1203 0.3 %Pt/25.%Nd~O3/Si-AI 0.6%Pt/Si-AI

79 44 53

27 11

4

18 28 78

.,.

51 57 15

,.

This is surprising in light of the close matching of acid strengths suggested by the 2MP2 reaction test, the comparable support surface areas, and similar platinum dispersions. Selectivites also differ dramatically in this reaction. Toluene selectivity is lower on rare earth Si-A1 than on chlorided alumina, with the rare earth catalyst preferentially forming C 7 isomers. The rare earth oxide catalyst shows slightly lower cracking selectivity than chlorided alumina and substantially lower cracking selectivity than Pt/Si-A1. These selectivity differences, indicate an inhibited metalcatalyzed dehydrocyclization activity, and when coupled with the lower observed benzene hydrogenation activity suggest that metal activity on the rare earth silica-alumina catalysts is inhibited relative to halided alumina catalysts. Apparently, platinum can readily dehydrogenate/hydrogenate n-C 7 on rare earth silica-alumina catalysts, although it does the more demanding direct dehydrocyclization poorly. The inhibition of metal activity relative to halided alumina catalysts limits direct metalcatalyzed dehydrocyclization reactions during heptane reforming but should not interfere with hydroisomerization reactions. Direct paraffin reactions require stronger acidity than olefin reactions, since paraffins are weaker bases. In classical bifunctional reactions involving paraffins, the metal initiates the reaction by dehydrogenating the paraffin to a more reactive olefin. This olefm reacts on an acid site to form a carbenium ion which in turn rearranges (isomerizes), and then decomposes (via proton loss) to an isomerized olefin. The latter is easily rehydrogenated to an isomerized paraffin over the metal. The bifunctional sequence normally has olefin isomerization on the acid site as the rate determining step, which gives an overall negative 1st order H 2 rate dependence. The acid strength and metal site activity must be "balanced" to avoid excessive metal or acid cracking. Alternative pathways found on stronger acid sites which involve intermolecular hydride transfer to heal carbocations producing isomefized paraffins (10) do not occur to any extent on any of the intermediate strength acids discussed here. We compared Pt/silica-alumina, yttria-modified silica-alumina, and fluorided alumina for n-C 12 isomefization. Not surprisingly, increasing yttria content lowers catalyst activity at a fLxed space velocity (Fig. 10). The 9% Y203/Si-AI catalyst compares closely to the l%F/AI203 catalyst in activity. Of the catalysts evaluated here, the 9%Y203-1oaded Si-Al had higher isomefization selectivity at equal conversion (Fig. 11).

571

8o

p'

0.02

-

"

o.o5

-o.~J

o.il

o.'14

I/Wl/SV

Fig. 10. n-C12 conversion over 0.3% Pt on silica-alumina, yttria-modified silica-alumina and FAI20 3 at 100 psi, 12"1 H2/n-C12, 350~

-

8s

O...,,

65

55

,

9 20

, 30

-

. 40

.

. 50

. 60

.

%n-~] Converslen

70

80

90

Fig. 11. i-C12 selectivity over 0.3% Pt on silica-alumina, yttria-modified silica-alumina and FAI20 3 at 100 psi, 12"1 H2/n-C12, 350~ CONCLUSIONS The addition of rare earth oxides to amorphous silica-alumina produces a modified support with a dispersed rare earth oxide surface phase and with acid properties similar to either chlorided or fluorided aluminas. X-ray diffraction does not show rare earth oxide crystallites (>30,~) up to 30% wt. rare earth oxide loadings. The ESCA (rare earth)/(Si + A1) signal increases linearly with increasing rare-earth oxide loading, strongly suggesting that the rare earth oxide continues to spread on the surface (i.e. remain below monolayer coverage) up to 30% wt. loading. The surface area and pore volume in the support decrease, but the average pore diameter increases as the dispersed rare-earth oxide blocks smaller pores, making a fraction of the surface inaccessible to physisorbed N 2. The olefin model compound reactions show that acid strength can be adjusted close to that of chlorided and fluorided aluminas with the addition of 520% wt. of rare earth oxides to the amorphous SiO2-A120 3 support. During benzene hydrogenation and n-C 7 reforming tests, metal activity on platinized rare earth oxide-modified silica-alumina remains inhibited relative to platinized halided-alumina catalysts, even at high rare earth oxide loadings, where the measured H 2 chemisorption and thereby the assumed dispersion equal that of halided aluminas. This reduced metal activity limits the direct metal-catalyzed dehydrocyclization reactions during heptane reforming, but facile platinum catalyzeddehydrogenation/hydrogenation reactions are not disturbed, so that hydroisomerization readily proceeds.

572 ACKNOWLEDGMENTS We thank Lenny Yacullo and John Ziemiak for their experimental assistance. REFERENCES 1. Misono, M. and Okuhara, T., Chemtech, 23-29 (1993). 2. Thomas, J.M., Scient. Amer., 266 (4) 112-118 (1992). 3. Arata, K., Trends Phys. Chem., 2 1-24 (1991). 4. Washington Bulletin, National Petroleum Refiners Association, Feb. 18, 1994. 5. Degnan, T.F. and Kennedy, C. R., AICHE Journal 39(4), 607-11 (1993). 6. Sinfelt, J. H., Carter, J.C., and Yates, D.S.C., J. Catal. 2..44,283-9 (1972). 7. Kramer, G.M. and McVicker, G.B., Acct. Chem. Res. 19, 78-86 (1986). 8. Maitra, A.M., Appl. Catal., A 85(1), 27-46, (1992). 9. Maitra, A.M., J. Therm. Anal. 36(2), 657-75, (1990). 10. Iglesia, E., Soled, S. and Kramer, G., J. Cat. 144, 238-45 (1993).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

573

W h a t N M R Has Told Us about Solid Acidity J. F. Haw a and J. B. Nicholas b aChemistry Department; Texas A&M University; College Station, TX 77843 USA bEnvironmental Molecular Sciences Laboratory; Pacific Northwest Laboratory; Richland, WA 99352 USA A wide variety of NMR methods are being applied to understand solid acids including zeolites and metal halides. Proton NMR is useful for characterizing Brensted sites in zeolites. Many nuclei are suitable for the study of probe molecules adsorbed directly or formed in situ as either intermediates or products. Adsorbates on metal halide powders display a rich carbenium ion chemistry. The interpretation of NMR experiments on solid acids has been greatly improved by the integration of theoretical chemistry and experiment.

1. OVERVIEW AND PERSPECTIVE In 1962 Olah reported the 13C NMR spectrum of the t-butyl cation in superacid solution, [ 1] and NMR was thenceforth the experimental method of choice for studies of intermediates in solution acid chemistry. The inhomogeneous nature and diversity of solid acid systems will ensure that no one experimental technique will so completely dominate as NMR has in solution studies, but the contributions and potential of NMR to solid acid studies are clearly such as to put it on an equal footing with reaction studies, infrared, TPD, diffraction methods and calorimetry. NMR of solids is a very diverse collection of methods, and the practice of applying it to chemisorption is changing dramatically as a result of advances in theoretical chemistry including reliable chemical shift calculations. The wide application of NMR to solid state chemistry grew out of the revival of magic angle spinning in the mid 1970's. This line narrowing technique is very effective for dilute or low ~,, spin 1/2 systems including most examples involving 13C, 31 p, 19F, 29Si and 15N as well as 1H at the chemical concentrations found in typical solid oxides including zeolites. Magic angle spinning is also partially to highly effective in narrowing the central transition for many quadrupolar nuclei including 27A1. More sophisticated line narrowing techniques involving more complicated mechanical averaging or elegant application of multil?le quantum evolution have extended the practice of high resolution solid state NMR to include 170, 23Na and many other nuclei that may sometimes be of interest in studies of catalytic materials. In addition to sample rotation, a particular solid state NMR experiment is further characterized by the pulse sequence used. As in solution NMR, a multitude of such sequences exist for solids; many exploit through-space dipolar couplings for either signal enhancement, spectral assignment, internuclear distance determination or full correlation of the spectra of different nuclei. The most commonly applied solid state NMR experiments are concerned with the measurement of spectra in which intensities relate to the numbers of spins in different environments and the resonance frequencies are dominated by isotropic chemical shifts, much like NMR spectra of solutions. Even so, there is considerable room for useful elaboration; the observed signal may be obtained by direct excitation, cross polarization from other nuclei or other means, and irradiation may be applied during observation or in echo periods prior to

574 observation to selectively remove or reintroduce dipolar interactions. Molecular dynamics on very diverse time scales also reveal themselves to suitable NMR techniques probing line shapes and relaxation times. Other kinds of NMR experiments applied to solid acids and other catalysts have little resemblance to the more familiar forms of spectroscopy. For example, diffusion in a magnetic field gradient interferes with refocusing of spin echoes, and this is applied in diffusion measurements. Other applications of gradients lead to a variety of imagining experiments which can be used to profile macroscopic inhomogeneities. Thus, the NMR characterization of a solid acid can range from, for example, obtaining a simple proton or silicon MAS spectrum using a routine method to an elaborate negotiation between the spectroscopist and the spi~ systems naturally contained within or introduced to probe the acid. The reader is referred the recent book by Bell and Pines [2] for a more complete overview of the various methods and objectives in NMR studies of solid acids and other heterogeneous catalysis. In the present contribution we illustrate the application of 1H, 13C and 19F MAS NMR to two archetypal solid acids, BrCnsted sites in zeolites and solid metal halides such as aluminum chloride and bromide powders which exhibit "Lewis superacidity". An important characteristic of the more recent work is the integration of quantum chemical calculations into the design and interpretation of the NMR experiments.

2.

P R O T O N NMR OF ZEOLITE BRONSTED SITES IN HZSM-5

We first illustrate the application of NMR to solid acids with some purely experimental results and then motivate theoretical methods and the integration of theory and experiment. Figure 1 shows 1H MAS spectra of a dehydrated sample of zeolite HZSM-5. Spectra obtained in the vicinity of room temperature seem to suggest no more than two isotropic signals as well as a series of spinning sidebands. The isotropic peak at 2.0 ppm is due to silanols, while that at 4.3 ppm to BrCnsted sites. The apparently singular isotropic shift for the Br~nsted sites in this material was long used to argue that all such sites are "identical", and indeed many non spectroscopic methods support the idea that the acid strengths are uniform in HZSM-5. Carefully inspecting the downfield (high shift) side of the 4.3 ppm resonance in the spectrum obtained at 296 K suggests a very broad line shape component. When spectra are run at much lower temperatures, this sharpens into an additional isotropic peak at 6.9 ppm. We used 1H{27Al } double resonance experiments to demonstrate that both the 6.9 and 4.3 ppm signals have strong dipolar coupling to aluminum; thus, it would appear that there are at least two NMR distinguishable Brcnsted sites in HZSM-5 [3]. The exact nature of the new site is under investigation in several laboratories. Figure 1. Proton MAS NMR spectra of dehydrated HZSM-5 acquired over a range of temperatures. Note in particular the broad signal at 6.9 ppm which is resolved at lower temperatures.

!''"

25

I~ : " 1 : : ' ' 1 ' : : : 1 ' " ' 1 ~

20

15

10

::':

5

0

-5

!::::1

-10 -15 ppm

575

3.

ZEOLITE PROBE MOLECULES AND REACTION INTERMEDIATES

It is often said that the property of acidity is manifest only in the presence of a base, and NMR studies of probe molecules became common following studies of amines by Ellis [4] and Maciel [5, 6] and phosphines by Lunsford [7] in the early to mid 80s. More recently, the maturation of variable temperature MAS NMR has permitted the study of reactive probe molecules which are revealing not only in themselves but also in the intermediates and products that they form on the solid acid. We carried out detailed studies of aldol reactions in zeolites beginning with the early 1993 report of the synthesis of crotonaldehyde from acetaldehyde in H Z S M - 5 [8] and continuing through investigations of acetone, cyclopentanone [9] and propanal [ 10]. The formation of mesityl oxide 1, from dimerization and dehydration of

H

1

2

acetone, was a very useful observation, because Farcasiu had recently proposed the use of the 13C shifts of this molecule as a measure of liquid acid strength [ 11 ]. In solution, the 13C shifts of I reflect the protonation equilibrium and the contributions of the resonance structures for the protonated ketone 2. In the acidic zeolites, acetone-2-13C exhibits a downfield shift (that is not nearly so large as that seen in 100 % H2SO4 solution), and 1 gives shifts that can be interpreted with Farcasiu's solution measurements to suggest that HZSM-5 has an effect on mesityl oxide that is similar to the effect of ca. 70 % H2SO4 [9]. Indeed, a large number of purely experimental observations of carbenium ions and related species underscore the fact that zeolites are strong Brcnsted acids but not superacids. Some of the more interesting "probe molecules" studied on solid acids by NMR have been carbenium ions generated in situ. These include the indanyl carbenium ion formed from the reaction of styrene dimers in HZSM-5 [ 12]. Although this result convincingly demonstrated the formation of a free cation as persistent intermediate species, the high stability (low acid strength) of this cation suggests that results like this are exceptional. 4.

THEORY, EXPERIMENT

AND H A M M E T T I N D I C A T O R S

The probe molecules of greatest historical interest in catalysis are the Hammett indicators [ 13]. The difficulty of making reliable visual or spectrophotometric observations of the state of protonation of these species on solids is well known. We have recently carried out the first NMR studies of Hammett indicators on solid acids [14]. This was also the occasion of the first detailed collaboration between the authors of this article, and theoretical methods proved to strongly compliment the NMR experiments. The Hammett story is told after first reviewing the application of theoretical chemistry to such problems. Central to the application of any physical method in chemistry is the process of modeling the relationship between the observables and molecular structure. However often one does this, it is rarely an exact process. One can rationalize almost any trend in isotropic chemical shift as a function of some variation in molecular structure - after the fact, but the quantitative prediction of such trends in advance defies intuition in most nontrivial cases. Even though the NMR spectrum is a function

576 of molecular structure, molecular structure is not a function of the NMR spectrum. Even if one correctly "assigns" spectral data to a gross chemical structure, the proposed structure rarely specifies exact bond lengths and angles, and the testable predictions that one makes from it are necessarily qualitative. Theoretical chemistry provides the means to model the relationship between structure and spectra, and this modeling process is revolutionizing the practice of interpreting NMR spectra. One strategy is as follows. 1. NMR experiments are interpreted to suggest a qualitative model of the chemisorbed state, the structure of an intermediate or some other finding. 2. The appropriate level of theory is then used to test the reasonableness of this finding as well as alternate explanations. 3. Finally one calculates spectroscopic observables (e.g., chemical shift tensors, vibrational frequencies) for the refined structure and compares to experiment. Agreement between the values of a variety of observables, determined by both theoretical and experimental means, provides powerful verification of our interpretation. Clearly, the integration of experiment and theory is much preferable to "hand-waving arguments" for the assignment of unusual chemical shifts to exotic species, as has often been done. There are many considerations that critically affect the accuracy of a quantum mechanical calculation, the most obvious of which are the degree to which the atomic coordinates of the theoretical model relate to the experimentally observed compound, the flexibility of the basis set and the extent to which electron correlation is accounted for [ 15]. For chemisorption studies in zeolites, cluster models of appropriate size have been used to elucidate a wide range of zeolite chemistry [ 16]. The perturbation theory of Mr and Plesset is the most straightforward way of improving a Hartree-Fock calculation; such treatments truncated at second order (MP2) usually provide much of the correlation effect at a reasonable computational cost. As an alternate means of including electron correlation, density functional theory (DFT) methods have been demonstrated to give results comparable to high level postHartree-Fock calculations with considerably less computational cost. Recent studies have validated the DFT approach in the computation of geometries, vibrational frequencies, energetics, and other molecular properties. More important to catalytic applications, DFT has also been shown to give good results for reaction barriers and transition state geometries in proton transfer reactions. Reliable NMR chemical shift calculations are now possible with a variety of means, including the IGLO, LORG, GIAO approaches, and with the inclusion of electron correlation by either MF2 or density functional methods. We used DFT to optimize the geometries of various Hammett bases on cluster models of zeolite Br~nsted sites. For p-fluoronitrobenzene and p-nitrotoluene, two indicators with strengths of ca. -12 for their conjugate acids, we saw no protonation in the energy minimized structures. Similar calculations using the much more strongly basic aniline analogs of these molecules demonstrated proton transfer from the zeolite cluster to the base. We camed out 19F and 15N experimental NMR studies of these same Hammett indicators adsorbed into zeolites HY and HZSM-5. Figure 2 compares the results of theory and experiment for the specific case of pfluoronitrobenzene. Inspection of the calculated structure shows that the proton is still on the zeolite, and the 19F shifts are more like chloroform solution than superacid solution. Furthermore, when the 19F chemical shift was calculated for the theoretical structure, it was found to agree with the experimental result.

577

Figure 2. The Hammett indicator p-fluoronitrobenzene on the BrCnsted site of zeolites. Left Theoretical structure calculated with DFT at the BLYP/DNP level. Right - Experimental 19F MAS NMR spectra on HY (top) and HZSM-5 (bottom). 5.

N M R S T U D I E S OF M E T A L H A L I D E SOLID ACIDS

We have discovered that a wide variety of carbenium ions and related electrophilic species can be prepared by the direct contact of suitable precursors with metal halide powders at reasonable temperatures and in the absence of solvent, permitting direct observation of reaction chemistry and measurement of the principal components of the 13 C chemical shift tensors [ 17]. The Friedel-Crafts alkylation and acylation reactions, first reported in 1877, are among the most important in chemistry. The generalized reaction, as surveyed in Olah's books [18, 19] proceeds on a wide variety of catalysts; the most familiar is AIC13, but other metal halides including ZnCI2 are valued for their lower activity and sometimes greater selectivity. Metal halide powders or molten salts are also employed as catalyst components in several industrial processes. We readily prepared the t-butyl 3 and acylium 4 cations on aluminum chloride powder from the corresponding halides. 3 and 4 are, respectively, the archetypal Friedel-Crafts alkylation and acyclation intermediates. From the 13C spectra of these and other cations, we were able to measure the principal components of the chemical shift tensors. The isotropic shift in solution is the average of the three principal components; it should be clear

i /C+

H

H3C

3

m C+

~CH 3

I 4

578 that correlations between structure and spectra should be far more meaningful when made for the principal components instead of their average alone. We have also been carrying out extensive theoretical calculations of the structures and chemical shift tensors of the cations observed in the experimental studies. As we develop greater confidence in the methodology we will be able to explore subtle effects due to ion pairing and other interactions with the media. This work is also valuable in that it assists our ongoing efforts to understand the nature of electropMlic species in zeolite solid acids. Even acetone shows appreciable shifts following contact with metal halides - the magnitude of which correlates with the expected strength of the Lewis acid-base interaction.

~.. 231 ppm

J"

I

227 ppm

Figure 3. 75.4-MHz 13C NMR spectra of a c e t o n e - 2 - 1 3 C on various substrates. The isotropic shift is strongly dependent on the strength of the Lewis acid. Note that ZnCI2 and zeolite ZnY yield the same shifts.

Zn !

231 ppm ZnY zeolite C P, 198 K . ~.4-

245 ppm

Alll 3 I '''

400

i l ' '""'I 'I

350

" ''

300

'"I ' ' "'

250

I "'""l"'""i'""i'"~""l

200

150

' ''"i'l

100

' '''

50

I

0

ppm Figure 3 shows 13C MAS spectra of acetone-2 -13C on various materials. Two isotropic peaks at 231 and 227 ppm were observed for acetone on ZnC12 powder, and appreciable chemical shift anisotropy was reflected in the sideband patterns at 193 K. The 231 ppm peak was in complete agreement with the shift observed for acetone diffused into ZnY zeolite. A much greater shift, 245 ppm, was observed on A1CI3 powder. For comparison, acetone has chemical shifts of 205 ppm in CDC13 solution, 244 ppm in concentrated H2SO4 and 249 ppm in superacid solutions. The resonance structures 5 for acetone on metal halide salts underscore the similarity of the acetone complex to carbenium ions. The relative contributions of the two canonical forms rationalizes the dependence of the observed isotropic 13C shift on the Lewis acidity of the metal halide.

579

CH 3

H

T

H3

H3 C

~O

.

.

.

n+ x

Mn+X n

m

n

5 There is good reason to believe that the potential of NMR studies of carbenium ions on solid metal halides exceeds that of corresponding studies in superacid solutions. Of course the advantages of working in solids include the possibility of very low temperatures and the mass transport restrictions of frozen media. Thus, Mehre and Yannoni were able to characterize the sec-butyl cation in frozen SbF5 by NMR [20] and Schleyer and coworkers have obtained infrared evidence of the allyl cation in the same medium [21 ]. So far, we have been successful in every case in which we have tried to duplicate known solution carbenium ion chemistry on appropriate metal halide solid acids. For example, we have recently prepared a number of acomplexes by the reaction or alkyl halides on aluminum halides. Figure 4 shows the case for the formation of the toluenium -13C7 ion. 32 ppm

O

+ CH3Br

139 ppm AIBr 3 "~

m~. 201 ppm AI2Br7"

178 pp n ~ 50ppm 32 ppm 178ppm

201ppm | 213 K I ''

250

~ ''

139ppm

Figure 4. 75.4 MHz 13 C NMR spectrum of benzene13C 6 and bromomethane13C reacting on A1Br3 powder. The spectrum shows the formation of t o l u e n i u m - 1 3 C 7 ion as indicated in the scheme. The resonance at 129 ppm is due to an excess of benzene- 13C6"

5

Q I'

200

'''

I'''

150

'I''

100

''

I'

50

'''I

0

ppm

REFERENCES .

2. .

G. A. Olah, Angew. Chem., Int. Ed. Engl., 34 (1995) 1393. A. Bell and A. Pines (eds.), NMR Techniques in Catalysis, Marcel Dekker: New York, 1994. L. W. Beck, J. L. White and J. F. Haw, J. Am. Chem. Soc., 116 (1994) 9657.

580

o

.

.

.

8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21.

W. H. Dawson, S. W. Kaiser, P. D. Ellis, and R. R. Inners, J. Am. Chem. Soc., 103 (1981) 6780. G. E. Maciel, J. F. Haw, I.-S. Chuang, B. L. Hawkins, T. E. Early, D. R. McKay and L. Petrakis, J. Am. Chem. Soc., 105 (1983) 5529. J. F. Haw, I.-S. Chuang, B. L. Hawkins and G. E. Maciel, J. Am. Chem. Soc., 105 (1983) 7206. P. Chu, D. D. Mallmann, and J. H. Lunsford, J. Phys. Chem., 95 (1991) 7362. E. J. Munson, and J. F. Haw, Angew. Chem., 105 (1993) 643. T. Xu, E. J. Munson, and J. F. Haw, J. Am. Chem. Soc., 116 (1994) 1962. T. Xu, J. Zhang, E, J. Munson and J. F. Haw, J. Chem. Soc. Chem. Comm. (1994) 2733. D. Farcasiu and A. Ghenciu, J. Am. Chem. Soc., 115 (1993) 10901. T. Xu and J. F. Haw, J. Am. Chem. Soc., 116 (1994) 10188. L. P. Hammett and A. J. Deyrup, J. Am. Chem. Soc., 54 (1932) 2721. J. F. Haw, J. B. Nicholas, L. W. Beck, T. R. Krawietz, and D. B. Ferguson, J. Am. Chem. Soc., in press. W. J. Hehre, L. Radom, P. v. R. Schleyer, and J. A. Pople, Ab Initio Molecular Orbital Theory; Wiley & Sons: New York, 1986. J. Saurer, Chem. Rev., 89 (1989) 199. T. Xu, P. D. Torres, L. W. Beck and J. F. Haw, J. Am. Chem. Soc., 117 (1995) 8027. G. A. Olah, Friedel-Crafts Chemistry; Wiley & Sons: New York, 1973. G. A. Olah and/~. Moln~, Hydrocarbon Chemistry; Wiley & Sons: New York, 1995. P. C. Myhre and C. S. Yannoni, J. Am. Chem. Soc., 103 (1981) 230. P. Buzek, P. v. R. Schleyer, H. Vancik, Z. Mihalic, and J. Gauss, Angew. Chem., Int. Ed. Engl., 33 (1994) 448.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) I l th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

581

N o v e l M i c r o p o r o u s Solid "Superacids": C s x H 3 - x P W 1 2 0 4 0 (2 < x < 3 ) t T. Okuhara a, T. Nishimurab,$, and M. Misonob aGraduate School of Environmental Earth Science, Hokkaido University, Sapporo 060, Japan bDepartment of Applied Chemistry, Graduate School of Engineering, The University of Tokyo, Bunkyo-ku, Tokyo 113, Japan

It was confirmed that pore size of the acidic Cs salts (CsxH3-xPWl2040) was controlled by the Cs content. For example, Cs2.5H0.5PW12040 had mesopores and micropores larger than 8.5 /~ in diameter, and C s 2 . 2 H 0 . s P W 1 2 0 4 0 micropores of about 7/~. Shape selective catalysis was clearly observed for the latter catalyst. The acidic Cs salts as well as H3PW12040 possess very strong acidity and were more catalytically active than zeolites and SO42-/ZRO2 for decomposition of esters and alkylation in liquid-solid reaction system. 1. I N T R O D U C T I O N Heteropolyacids are good catalysts for acid-catalyzed reactions and have been used in large-scale catalytic processes [1-5]. The formation of Cs salts brings about high surface area [6-9] and high catalytic performance for both acidcatalyzed reactions [9,10] and oxidation reactions [11,12]. We indicated previously that H 3 P W 1 2 0 4 0 and its acidic Cs salt possess "superacidity", as far as the H a m m e t t indicator test and temperature programmed desorption of NH3 are concerned [13]. Recent data of calorimetry of NH3 absorption by Lefevre et al. [14] demonstrate that H3PW12040 is a very strong acid. Excellent catalytic performances of H3PW12040 and its Cs salts, especially Cs2.5H0.5PW12040, have been shown for acid-catalyzed reactions such as direct decomposition of ester, alkylation [13], acylation [15], skeletal isomerizati0n [16], etc. tCatalysis by Heteropoly Compounds. XXVII. Part XXVI: S. Shikata, T. Okuhara, and M. Misono, J. Mol. Catal., 100 (1995) 49. SPresent address: Mitsui Toatsu Chemicals, Inc., Kasama-cho, Sakae-ku, Yokohama 247, Japan.

582 Shape selective catalysis as typically demonstrated by zeolites is of great interest from scientific as well as industrial viewpoint [17]. However, the application of zeolites to organic reactions in a liquid-solid system is very limited, because of insufficient acid strength and slow diffusion of reactant molecules in small pores. We reported preliminarily t h a t the microporous Cs salts of H 3 P W 1 2 0 4 0 exhibit shape selectivity in a liquid-solid system [18]. Here we studied in more detail the acidity, micropore structure and catalytic activity of the Cs salts and wish to report that the acidic Cs salts exhibit efficient shape selective catalysis toward decomposition of esters, dehydration of alcohol, and alkylation of aromatic compound in liquid-solid system. The results were discussed in relation to the shape selective adsorption and the acidic properties. 2. E X P E R I M E N T A L Acidic Cs salts of H3PW12040, CsxH3-xPWl2040 (abbreviated as Csx), have been prepared by titration of the aqueous solution of H 3 P W 1 2 0 4 0 with an aqueous solution of Cs2CO3 [13]. To the solution of H3PW12040 (8 x 10 -2 tool din'3), the appropriate amount of aqueous solution of Cs2CO3 (Cs: 2.0 x 10 -1 tool dm "3) was added dropwise at a rate of about 1 ml rain "1 at room temperature. The resulting white colloidal solution was allowed to stand overnight at room temperature and then was evaporated at 318 K to the solid. An isotherm of N2 adsorption and surface area were obtained by using a N2 adsorption system (Micromeritics ASAP 2000). The adsorption capacities for various molecules having different molecular size were measured by microbalance (Shimadzu TG-30) connected directly to an ultrahigh vacuum system [18,19]. The adsorption of isopropylacetate or cyclohexylacetate was measured in a liquid-solid reactor at 303 K using decane as solvent. The shape selective catalysis was examined by choosing five kinds of the reactions, Eqs. (1) - (5), that is, dehydration of 2-hexanol, decompositions of three kinds of esters and alkylation of 1,3,5-trimethylbenzene with cyclohexene. Catalytic reactions were performed in a three-neck flask (about 100 ml) after the catalyst (about 100 rag) was pretreated in a He flow at 573 K. Reactions were

/vr,

Hexenes

+

H20

(I)

--OCOCH3

,/%

+

CH3COOH

(2)

--OCOCH3

Butenes

+

CH3COOH

(3)

+

CH3COOH

(4)

OH

<~-OCOCH3

O

(5)

583 carried out with 8 mmol (in 142 mmol of nonane) of 2-hexanol at 373 K, 216 mmol of 1,3,5-trimethylbenzene and 10 mmol of cyclohexene at 343 K, isopropylacetate (17 mmol in 142 mmol of undecane) at 373 K, sec-butylacetate (15 mmol in 142 mmol of undecane) at 373 K, and cyclohexylacetate (14 mmol in 112 tool of nonane) at 373 K, respectively. 3. R E S U L T S AND DISCUSSION 3.1. Acidity a n d catalytic activity The acid strength of typical solid acids, which has been measured by different methods, is summarized in Table 1. Hammett indicator test indicated t h a t H 3 P W 1 2 0 4 0 and C s 2 . 5 H 0 . 5 P W 1 2 0 4 0 (abbreviated as Cs2.5) possess the "superacidity", since HO < 13.16 [10,13]. The higher acid strength of these heteropoly compounds than HZSM-5 and SIO2-A1203 was also confirmed by temperature programmed desorption [13]. Recently, Lefebvre et al. found in adsorption (absorption) calorimetry of NH3 that the anhydrous H 3 P W 1 2 0 4 0 gave the initial heat of 195 kJ tool "1 and claimed that H 3 P W 1 2 0 4 0 has superacidic sites [14]. In the case of HZSM-5, the initial heat of adsorption was reported to be about 150 kJ tool "1 [20], indicating that HZSM-5 was a weaker acid than H3PW12040. Fogash et al. [21] observed that small amounts (- 5 ~tmol g-l) of strong acid sites, of which the initial heat was about 190 kJ tool -1, were present on SO42-/ZRO2 and the heat of NH3 adsorption was 150-165 kJ tool -1 at the adsorption of about 20 pmol g-1. Temperature programmed desorption of NH3 shows that the order of the acidstrength is as follows: SO42-/ZRO2 > H3PW12040 - Cs2.5H0.5PW12040 > HZSM-5 > SIO2-A1203. This order is consistent with the results of Hammett indicator test. Table 1 Acid strength of various solid acids Solid acid HO a - AHads(NH3) b Tdes(NH3) c Ref. /kJ mol-1 /K H3PW12040 -13.16 195 -850 13, 14 C s2.5H0.5PW 12040 - 13.16 -830 13 SO42"/ZRO2 -14.52 165 - 1000 e 13, 21 (190) d HZSM-5 -12.70 150 -670 13, 20 Si02-A1203 -12.70 145 -600 13, 22 aAcidity function measured by Hammett indicator test. bInitial heat of NH3 adsorption (absorption). CDesorption temperature of NH3; the figures show the temperature of the desorption peak at the highest temperature, dAt less than 5 ~tmol g-1 of NH3 adsorbed, eDesorbed as N2 by the reaction with catalyst [13]. -

584 Table 2 demonstrates the high catalytic performance of Cs2.5. For the decomposition of cyclohexylacetate (Reaction (4)), the activity (per unit weight) of Cs2.5 is about 40 and 130 times as high as SO42"/ZRO2 and HZSM-5, respectively, and furthermore it is more active than H2SO4 by a factor of two orders of magnitude. Cs2.5 is also highly active for the alkylation of 1,3,5trimethylbenzene with cyclohexene (Reaction (5)). The high activity of Cs2.5 has been inferred to be due to the high acidity (the acid strength and amount) and the soi~ basicity of the polyanion [13]. Table 2 Catalytic activities of various acids for decomposition of cyclohexylacetate and aUTlation of 1,3,5-trimethylbenzene in liquid-solid reaction system Catalyst Surface area Reaction rates a m 2 ~-1 Reaction (4) b Reaction (5) c Cs2.5H0.sPW12040 135 130 58 H3PW12040 5 50 18 SO42"/ZRO2 93 "3 5 HZSM-5 332 1 < 0.5 SIO2-A1203 458 0.6 1 H2SO4 ammol g-1 h-1. b a t 373 K. CAt 343 K.

1

3

Figure 1 shows more detailed data of the change of the surface area as a function of the Cs content, x in CsxH3-xPW12040. The surface area of the acid form (x = 0), 5 m 2 g-l, decreased as x increased to 2 (the surface area of Cs2 was only 0.5 m 2 g-l). However, the surface area greatly increased as x exceeded 2 and

Figure 1. Surface area and surface acidity of CsxH3-xPWl2040.

585 became higher t h a n 130 m 2 g-1 when x > 2.5. This change was in contrast with the acidic Na salts; the surface area decreased monotonically as the Na content increased [23]. The acid amount on the surface, which is called surface acidity hereafter, was estimated from the surface area and the formal concentration of proton attached to polyanion. Here, the number of polyardons was calculated by assuming t h a t the hemisphere of each polyanion (11 A in diameter) contributes to the surface area. The number of protons was then estimated on the assumption t h a t each polyardon on the surface possesses (3 - x)/2 protons. Since solid state NMR revealed t h a t protons of the acidic Cs salts distribute almost uniformly t h r o u g h the whole bulk [24], this estimation of the" surface acidity m a y be reasonable. The surface acidity thus estimated is shown in Figure 1. The surface acidity decreases at first with the Cs content, but sharply increases when x exceeds 2. The maximum appeared at x =2.5. In Figure 2, the catalytic activities of Csx ( x = 0, 2.2, 2.3, 2.5, 2.7, 2.9 and 3) for the decomposition of isopropylacetate (Reaction (2)) are plotted a g a i n s t the surface acidity. A good correlation between the activity and the surface acidity was observed for these acidic Cs salts. As described above, the acid s t r e n g t h of Cs2.5 was very similar to that of H 3 P W 1 2 0 4 0 , indicating t h a t the effect of the formation of acidic Cs salts little influences the acid strength. Therefore, it is reasonable that the catalytic activity of Csx for this reaction is proportional to the surface acidity, that is, the number of protons on the surface. HZSM-5 and SO42/ZrO2 were almost inactive for this reaction under the same reaction conditions. ._

80

0

c~ loo

v

Cs2.2

~ L-

~" b~

v-

.5

.7

,-

. mo.

_

20

80 60 40

.

, 40

.

, 60

8O

Figure 2. Catalytic activities of CsxH3-xPW12040 for decomposition of isopropylacetate as a function of the surface acidity. The reaction was carried out at 373 K in liquid-solid reaction system.

6 0 x: 0

E

tt~

v

I

C

9 9

0

40 E

-

I -

o

Surface Acidity/llmol g-1

<1 '7

I.. .

-

~ 2o

Ow_~Cs3_ , 0 20

~1

"7

t_

cc

80

Cs2.3

1

I I

"7

E

~" 4 0 ._o

~

~Cs20

I

v

Cs2.3/'

E 60 ~-

-

I

20

c,L ',

-2

0

.1 C's2.2

9

i l l v

40

~

i

60

e-o

o

9

80

N

Surface Acidity/l~mol g-1

Figure 3. C&talytic activities of CsxH3-xPWl2040 for decomposition of cyclohexylacetate and the alkylation of 1,3,5-trimethylbenzene as a function of the surface acidity.

586

3.2. Pore structure and shape selective catalysis In Figure 3, the rates of Reactions (4) and (5) are plotted against the surface acidity. In contrast with the results in Figure 2, Cs2.2 (or Cs2.1) was inactive for the decomposition of cyclohexylacetate and the a l k y l a t i o n of 1,3,5trimethylbenzene, while the other acidic Cs salts exhibited activities corresponding to the surface acidities. The critical molecular sizes of the reactant molecules are about 5, about 6 and 7.5/k for isopropylacetate, cyclohexylacetate and 1,3,5trimethylbenzene, respectively [18,25,26]. Therefore, the abnormally low activities of Cs2.2 and Cs2.1 for the decomposition of cyclohexylacetate and the alkylation of 1,3,5-trimethylbenzene may be attributed to the shape selective behavior; if Cs2.2 and Cs2.1 have small pores, the larger molecules like cyclohexylacetate and 1,3,5-trimethylbenzene may be unable to enter the pores and may not react there. In order to elucidate the pore structure of Csx, the adsorption-desorption isotherm of N2 was first measured. Typical results are given in Figure 4. H3PW12040 exhibited a Type II isotherm (according to the IUPAC classification [27]), indicating that this is a non-porous adsorbent. Cs2.5 showed a Type IV isotherm, similar to Cs3 [8], in which a hysteresis was observed for the desorption branch. The Type IV isotherm is usually observed for materials having mesopores [28]. In addition, Cs2.3, Cs2.7, and Cs2.9 gave the Type IV isotherms, showing that these have mesopores in addition to micropores. On the other hand, as shown in Figure 4, the isotherm of Cs2.2 (72 m 2 g-l) was of Type I [27], where most of the adsorption took place at the very low pressure region, indicating that these salts have only micropores (according to the IUPAC definition, the pore diameter is less than 20 A) [28]. Cs2.1 also gave the Type I

80 I&. IZ

v

Cs2.5 60

E r 0

40

E

Cs2.2

,<

20 t,__

0 '13

,<

0

0

0.2

0.4

0.6 P/Po

0.8

1.0

Figure 4. Isotherms of N2 adsorption on CsxH3-xPWl2040 at 77 K. The catalysts were pretreated at 573 K in vacuum. ~ 9adsorption branch, 9desorption branch. H3 denotes H3PW12040.

587 isotherm. It was further confirmed t h a t HY zeolite (JRC-Z-HY4.8, Reference Catalyst of Catalysis Society of Japan) showed the Type I isotherm. The data of N2 adsorption indicate that Cs2.2 and Cs2.1 were microporous. The pore size of Cs2.2 and Cs2.1 cannot be determined by the N2 adsorption, so t h a t their pore sizes were estimated from the adsorption of molecules h a v i n g different molecular size. Table 3 compares the adsorption capacities of Csx for various molecules measured by a microbalance connected directly to an ultrahigh vacuum system [18]. As for the adsorption of benzene (kinetic diameter = 5.9 A [25]) and neopentane (kinetic diameter = 6.2 A [25]), the ratios of the adsorption capacity between Cs2.2 and Cs2.5 were similar to the ratio for N2 adsorption. Of interest are the results of 1,3,5-trimethylbenzene (kinetic diameter = 7.5/~ [25]) and triisopropylbenzene (kinetic d i a m e t e r = 8.5 A [25]) . Both a d s o r b e d significantly on Cs2.5, but little on Cs2.2, indicating that th, pore size of Cs2.2 is in the range of 6.2 -7.5/~ and that of Cs2.5 is larger than 8.5 A-in diameter. In the case of Cs2.1, both benzene and neopentane adsorbed only a little. Hence the pore size of Cs2.1 is less t h a n 5.9/~. These results demonstrate t h a t the pore structure can be controlled by the substitution for H + by Cs +. Table 3 Adsorption data for CsxH3-xPW12040 (x = 2.1, 2.2 and 2.5) Molecule Kinetic Pressure Temp. Amount of adsorption (cross diameter//~ b /Torr /K /j~nol g-solid -1 section/A2) a (p/p0) c Cs2.1 Cs2.2 Cs2.5 N2 3.6 137 77 487 861 1648 (16.2) (0.18) Benzene 5.9 21 300 10 124 232 (30.5) (0.20) Neopentane 6.2 99 273 5 179 390 (37.2) (0.19) 1,3,5-TMB e 7.5 0.6 300 11 237 (41.1) (0.20) 1,3,5-TIPB f 8.5 0.009 300 15 236 (59.4) (0.20)

Ratio d 0.52 0.53 0.46 0.05 0.06

aCalculated from the molecular weight and density of liquid, bFrom refs. 25 and 26. e'I~e ratio of the partial pressure at equilibrium (P) to the saturated vapor pressure (P0) is given in parentheses, dAdsorption amount on Cs2.2 divided by t h a t o n Cs2.5. el,3,5-trimethylbenzene, fl,3,5-triisopropylbenzene. Figure 5 shows the catalytic activities of small-pore Cs2.1 and Cs2.2 relative to Cs2.5. Cs2.5 catalyzed all reactions with considerably high activities (the reaction rates are shown in the parentheses in Figure 5). On the other hand, although Cs2.2 was as active as Cs2.5 for the dehydration of 2-hexanol and decomposition of isopropylacetate, it was much less active for the decomposition of cyclohexylacetate and the alkylation of 1,3,5-trimethylbenzene. In the case of

588

Cs2.1, the activity was observed only for dehydration of 2-hexanol. The activities per unit surface area showed the same trend as Figure 5. The critical sizes of the reactant molecules were estimated and are shown in Figure 5, where the figures for 2-hexanol, isopropylacetate, sec-butylacetate and cyclohexylacetate are estimated by MM2 from Pauling's atomic radius and molecular model [18]. Therefore, the unique catalysis of Cs2.2 is understood if one assumes that it is active only for small molecules. In other words, this catalyst exhibits "reactant shape selectivity", where the catalyst differentiates the reactants according to their size. The decomposition of cyclohexylacetate was very slow on Cs2.2 (Fig. 5) , although the molecular size of cyclohexylacetate (~6.0 A) is smaller than the pore size of Cs2.2 (6.2 ~ 7.5 A). Probably the adsorption or diffusion of cyclohexylacetate in the pore is restricted by the coexisting solvent molecule, as the pore size is only slightly greater than the size of reactant.

Figure 5. Relative catalytic activities of CsxH3-xPWl2040 (x=2.1, 2.2 and 2.5) for various kinds of reactions in liquid-solid reaction system, aMolecular size. bpore size. Catalytic activity was estimated from the initial rate of the reaction. The activity of Cs2.5 for each reaction is taken to be unity. The figures in the parentheses are the reaction rates in the unit of retool g-1 h-1. In order to confirm further the restriction of the adsorption by the pore in the liquid-solid reaction system, the adsorption of the reactants was measured at 303 K in the liquid- solid system, at which no reaction took place. Figure 6 provides the time courses of the adsorption of cyclohexylacetate and isopropylacetate on Cs2.2 and Cs2.5. Both isopropylacetate and cyclohexylacetate adsorbed on Cs2.5 (Figure 6a). The ratio of the adsorption amount was about 1.5 times which is

589 close to the ratio of the molecular cross section of the two molecules. On the other hand, the adsorption of cyclohexylacetate was about one-tenth t h a t of isopropylacetate on Cs2.2 (Figure 6b). This result shows t h a t the shape selectivity exists in adsorption as well and adsorption of molecules into the pore is difficult even when the molecular size is less than the pore size in the liquid-solid reaction system, if the difference is small. This explains the low activity of Cs2.2 for the decomposition of cyclohexyacetate.

"o (D IL_

o

0.5

a

0.4

e ~ 0.3

"E

0.1

:3 0

E <

v.nJ..

0

,

,

2

,

,

4

,

,

6

,

u

0 Time/h

2

4

6

8

Figure 6. Time courses of the adsorption of isopropylacetate (A) or cyclohexylacetate (O) on Cs2.5H0.5PW12040 (a) and Cs2.2H0.8PW12040 (b). Results of Figures 4 and 5 suggest that the pore size of these materials can be precisely and nearly continuously modified by the Cs content, and accordingly the catalytic function is controlled. The pore size of Cs2.1 was smaller than that of Cs2.2, as was estimated to be less than 5.9 ,~ from the adsorption of benzene. In accordance with this, Cs2.1 had an activity for the dehydration of 2-hexanol, but was inactive for other reactions, irrespective of its considerably high surface area measure by N2 (55 m2g'1). Line-width of XRD peaks of these acidic Cs salts (x = 2.1~ 2.2, and 2.5) show t h a t the size of the primary crystallites was about 120 A and the primary crystallites themselves are nonporous [29]. Therefore, the pores observed in the present study correspond to the interparticle voids (not intraparticle). The pore structure and the mechanism of the formation of pores will be discussed in our forthcoming paper. As was stated above, the very strong acidity (and probably together with the organophilicity of the pore wall) makes these salts very active catalysts in liquidsolid organic reaction systems. We wish to emphasize that this is the first example for the shape selective catalysis of heteropolyacids at least to our knowledge.

590 REFERENCES

1. M. Misono, Catal. Rev.-Sci. Eng., 29 (1987) 269; 30 (1988) 339. 2. M. Misono, Proc. 10th Intern. Congr. Catal. (eds. L. Guczi, F. Solymosi and P. Tetenyi), Elsevier-Amsterdam, Akademiai Kiado, Budapest, 1993, p. 69. 3. Y. Ono, Perspectives in Catalysis (eds. J. M. Thomas and IL I. Zamaraev), Blackwell Scientific Publ., London, 1992, p. 431. 4. Y. Izumi, tL Urabe and M. Onaka, Zeolite, Clay, and Heteropoly Acid in Organic Reactions, Kodansha, Tokyo and VCH, Weinheim, New York, 1992. 5. I.V. Kozhevnikov, Russ. Chem. Rev., 62 (1993) 473. 6. S.J. Gregg and M. M. Tayyab, J. Chem. Soc. Faraday Trans., I, 74 (1978) 348. 7. M. Furuta, K. Sakata, M. Misono and Y. Yoneda, Chem. Lett., 1979, 31. 8. J.B. MacMonagle and J. B. Moffat, J. Coll. Interface Sci., 101 (1984) 479. 9. S. Tatematsu, T. Hibi, T. Okuhara and M. Misono, Chem. Lett., 1984, 865. 10. M. Misono and T. Okuhara, CHEMTECH, 23 (1993) 23 and references cited therein. 11. M. Misono, K. Sakata, Y. Yoneda and W. Y. Lee, Proc. 7th Intern. Congr. Catal., Tokyo, 1980, p. 1047, Kodansha, Tokyo and Elsevier, Amsterdam, 1981. 12. S. Nakamura and H. Ich~ashi, Proc. 7th Intern. Congr. Catal., Tokyo, 1980, Kodansha, Tokyo and Elsevier, Amsterdam, 1981, p. 755. 13. T. Okuhara, T. Nishimura and M. Misono, J. Mol. Catal., 74 (1992) 247. 14. F. Lefevre, F. X. Liu-Cai and A. Auroux, J. Mater. Chem., 4 (1994) 125. 15. Y. Izumi, M. Ogawa, W. Nohara and K. Urabe, Chem. Lett., 1992, 1987. 16. tL Na, T. Okuhara and M. Misono, J. Chem. Soc., Faraday Trans., 91 (1995) 367. 17. S. M. Csicsery, Pure Appl. Chem., 58 (1986) 841. 18. T. Okuhara, T. N i s b i m ~ a and M. Misono, Chem. Lett., 1995, 155. 19. K. Inumaru, T. Okuhara and M. Misono, J. Phys. Chem., 95 (1991) 4826. 20. D.J. Parrillo, C. Lee, R. J. Gorte, D. White and W. E. Farneth, J. Phys. Chem., 99 (1995) 8745. 21. IL B. Fogash, G. Yaluris, M. R. Gonzalez, P. Ouraipryvan, D./~ Ward, E. I. Ko and J. A. Dumesic, Catal. Lett., 32 (1995) 241. 22. H. Taniguchi, T. Masuda, IL Tsutsumi and H. Takahashi, Bull. Chem. Soc. Jpn., 53 (1980) 362. 23. T. Hibi, Takahashi, T. Okuhara, M. Misono and Y. Yoneda, Appl. Catal., 24 (1986) 69. 24. T. Okuhara, T. Nishimura, H. Watanabe, tL Na and M. Misono, AcidBase Catalysis II, Kodansha, Tokyo and Elsevier, Amsterdam, 1994, 419. 25. D.W. Breck, Zeolite Molecular Sieves, John Wiley and Sons, New York, 1974. 26. M.E. Davis, C. Saldarriaga, C. Montes, J. Graces and C. Crowder, Zeolites, 8 (1988) 362. 27. tL S. W. Sing, D. H. Everett, R. A. W. Haul, L. Moscow, R. A. Pierott, J. Rouquerol and T. Siemieniewska, Pure Appl. Chem., 57 (1985) 603. 28. S.J. Gregg and K. S. W. Sing, Adsorption, Surface Area and Porosity, 2nd Edn, Academic Press Inc., London, 1982. 29. N. Mizuno and M. Misono, Chem. Lett., 1987, 967.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis -40th Anniversary Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

591

C o m p a r i s o n of t h e r e a c t i v i t i e s of H 3 P W 1 2 0 4 0 a n d H 4 S i W 1 2 0 4 0 a n d t h e i r K +, N H 4 + a n d alkylation.

Cs + s a l t s

in liquid phase

isobutane/butene

N. Essayem, S. Kieger, G. Coudurier and J.C. V~drine. Institut de Recherches sur la Catalyse, CNRS, UPR associ~ ~ l'UCB Lyon, 2 avenue Albert Einstein, F-69626 VILLEURBANNE C~dex, France. Abstract

Activity in n-butane isomerization reaction of various alkaline salts of H3PW12040 and H4SiW1204o was shown to be strongly dependent on the strength and number of accessible protons whereas the stability with time on stream was correlated to the presence of mesoporosity. For the liquid iC4/C4 = continuous alkylation reaction, the strength and the n u m b e r of acid sites appeared less important than the existence of mesoporosity indicating that the diffusion of the reactants and of the products plays an important role in this reaction.

1. INTRODUCTION C4 aliphatic alkylation presents a great i n t e r e s t in petroleum refinery processes [1]. This reaction is a great challenge for heterogeneous catalysis since previous studies realized on classical acid solids (zeolites, SO42-/ZRO2) exhibited a rapid deactivation of the solid attached with a high production of byproducts particularly those produced by dimerisation of butenes [2-7]. A recent study [8] has claimed that c~esium salts of 12-tungstophosphoric acid, Cs2.5Ho.5PW12040 was much more active than the parent acid H3PW12040 and also than 8042-/ZRO2 in aliphatic alkylation. Further, in a recent work [9] focused on the behavior of CsxH3-xPW12040 in gas phase isomerisation of nbutane, we have pointed out that over the enhancement of the protonic surface acidity, the partial exchange of H § with Cs + cations produced solids with different textures. The aim of the present work was the investigation of the catalytic reactivity of different salts (K +, NH4 § Cs +) of H3PW12040 and H4SiW12040 with various compositions in continuous liquid phase alkylation and its comparison with nbutane isomerisation reaction in gas phase. The results are discussed in terms of strength and number of protonic sites and the presence of mesoporosity is shown to be important for the production of trimethylpentanes in the alkylation reaction.

592 2. EXPERIMENTAL H 3 P W 1 2 0 4 0 , n H 2 0 and H 4 S i W 1 2 0 4 0 , n H 2 0 samples were prepared in a classical way [10]. Alkaline or am m oni um salts were prepared as previously described [9] by adding the stoichiometric amounts of the aqueous solution (5M) of the corresponding chloride to an aqueous solution of H 3 P W 1 2 0 4 0 or H 4 S i W 1 2 0 4 0 (0.1M). The resulting suspension was kept under vigourous stirring for 24 h. Then, the solids were recovered after 2 washings and separation by centrifugation. The samples were characterized by chemical analysis (induced coupled plasma an d atomic absorption techniques apparatus), nitrogen adsorption i s o t h e r m s (at 77 K), XRD p a t t e r n s (Siemens diffractometer and CuKa radiation), SEM observations (Hitachi $800 apparat us of the University C. Bernard, Lyon I) and TGA-DTA (Setaram 92-12 apparatus). The IR spectra were recorded with a Bruker IFS 48 FTIR spectrometer. n - b u t a n e isomerization reaction was carried out in a differential flow microreactor in the following conditions: reaction temperature = 473 K, total gas pressure = 105 Pa, 5 % n-butane in H2, total flow rate = 18 cm3.min-1. Prior to reaction, the c a t a l ys t (300 mg) was p r e t r e a t e d in N2 at the chosen temperature for 2 hours. Liquid iC4/C4 = continuous alkylation reaction was carried out in the following conditions" the two reactants iC4 and C4 = were introduced by mean of two HPLC pumps and were kept under a pressure of 6.105 Pa. An additional circulation of high pressure He was used in order to impose the total pressure during one run, usually 20.105 Pa. The liquid products were depressurized and vaporized before being analysed by on-line gas chromatography equipped with a semi-capillary column (GS Alumina, 30mx0.53mm/Altech). Prior to the reaction run, the catalyst was treated in flowing He at the chosen temperature for one hour and then cooled to the reaction temperature. The usual conditions of reactions were as follows: reaction temperature = 333K, total pressure = 20.105 Pa, molar ratio iC4/C4 = = 21, total liquid flow rate: 0.2 cm3-min -1, catalyst weight: 500 mg. 3. CATALYSTS CHARACTERIZATION 3.1. Chemical composition a n d textural features The chemical compositions of the samples, obtained from chemical analyses are reported in Table 1. In order to check the chemical analyses, the mother and washing liquors were collected, analysed and their acidity was titrated. In all cases, the alkaline cations were detected only as traces. The acidimetric titration allowed us to determine the HPA amount remaining in the solution. On the other hand, the samples separated after precipitation and washings were weighted in order to calculate the precipitate yields. The results are reported in table 1 where the samples are designated as MxY (M being the alkaline or ammonium cation, Y the heteroatom, x the stoichiometry deduced from chemical analyses. This table shows that whatever be the nature of the cation, the neutralization of the acid favoured the precipitation of the neutral salt leading to a precipitate with a stoichiometry higher than that expected. In good agreement with these

593 results, the yields of precipitation increased with the M]HPA stoichiometry of synthesis reaching almost 100 % for the stoichiometry of the n e u t r a l salts. The BET surface a r e a values are also reported with the d i s t r i b u t i o n of porosity b e t w e e n microporosity (pore d i a m e t e r <1.8 nm) deduced from N2 adsorption i s o t h e r m s (t-curves) and mesoporosity (pore d i a m e t e r > 1.8 nm). The following t r e n d is observed: for high atomic M]HPA ratio used for the precipitation, the p r e c i p i t a t e s exhibited high surface a r e a m a i n l y due to microporosity. However, depending on the n a t u r e of the counter cation a n d also of the previous ratio values, the textural characteristics were not similar. In particular, it is i n t e r e s t i n g to note the presence of mesopores for (NH4)2.4P, Cs2.9P, Cs2.7P and Cs2.4Si samples. Table 1 Chemica!..compositions and textural properties of the samples Sample Stoichiometry a Precipitate S BET Synthesis condition

Precipitate composition chem Acid. c

yield b

m2-g-1

(%)

Surface area (m2.g-1) Pore diameter (nm) 0-1.8 >1.8

Cs2.9P Cs2.7P Csl.9P

3.0 2.5 2.0

2.9 2.7 1.9

1.9

100 95

147 138 71

108 94 64

39 44 7

K3P K2.2P

3.0 2.5

3.0 2.2

3.0 2.6

100 98

156 132

139 111

17 21

(NH4)2.8P (NH4)2.4P

3.0 0.8

2.8 2.4

3.0 -

100 41

121 60

111 20

10 40

Cs2.9Si Cs2.4Si

3.0 2.0

2.9 2.4

2.9 -

96 76

185 189

167 149

18 40

aAtomic ratio: alkaline cation/central atom (P or Si). bMolar ratio: precipitate HPA/initial HPA. Cdeduced from acidimetric titration of the m o t h e r solutions. 3.2. S t r u c t u r a l f e a t u r e s of the s~mples The XRD p a t t e r n s of a l k a l i n e or a m m o n i u m salts of H 3 P W 1 2 0 4 0 or H4SiW12040 were typical of the cubic structure. No additionnal XRD peak was observed i n d i c a t i n g the absence of other phases. The same u n i t cell was measured for the different Cs salts of H3PW12040 (ao = 1.184 nm) and for the Cs salt of H4SiW12040 ( ao = 1.178 nm), in agreement with the value reported in the literature for the neutral salts [9,11] whatever be the Cs content. Since the values of the unit cell p a r a m e t e r of the acidic salts were not i n t e r m e d i a t e between those of the neutral salts and of the p a r e n t acid (ao = 1.216 and 1.217 nm, respectively, for H3PW12040 and H4SiW12040), one m a y propose t h a t there is no formation of solid solutions but r a t h e r a mixture of the neutral salt and of the acid highly dispersed on the neutral salt.

594 3.3. I ~ A - T G A a n a l y s e s DT-TG curves for all highly exchanged Cs and K salts showed nor weight loss neither thermal peak up to 1023K, except a weight loss accompanied by an endothermal peak before 423K due to physisorbed water, indicating the very high thermal stability of these samples. On the other hand, for the ammonium salts, aider the first weight loss, a gradual weight loss was observed above 673K corresponding to NH3 departure and an exothermal peak (without weight loss) at 873 K due to the cristallisation of a tungsten oxide stemming from the decomposition of the Keggin unit. The thermograms of the different acidic salts were similar to those already described for CsxH3-xPW12040 compounds [9]. In addition to the initial weight loss present in the neutral salts, a second loss was observed beetwen 623 and 823 K. The first loss is due to the departure of physisorbed and/or crystallization water and the second one to a deprotonation step [9] which allowed us to calculate proton content values. These values are in good a g r e e m e n t with chemical analysis data. Comparison with H3PW12040 and H 4 S i W l 2 0 4 0 thermograms suggests that the acidic salts may be a mixture of the neutral salt of high thermal stability and of the pure acid which was deprotonated between 623 and 823 K. SEM observations of the solids before and after TGA experiments support this assumption.

Figure 1. SEM photographs of Cs2.4H1.6SiW12040 (a)before and (b) after thermal treatment up to 1023 K. As it can be seen on Figure 1, the C2.4Si sample before TG analysis was composed of e l e m e n t a r y particles (9-30 nm) being agglomerated in large spherical shaped particles (100-1000nm) with a spongy aspect. Such a morphology was already described for Cs salts of H3PW12040 [9,12,13]. After TG analysis up to 1023K, the SEM observation revealed the presence of a new shape of large cristallites while the initial spherical agglomerations remained but the e l e m e n t a r y particles a p p e a r e d more individualized. F u r t h e r m o r e , XRD pattern of this resulting solid showed the peaks characteristic of the neutral salt and additional peaks attributed to the tungsten oxide, W24068.

595 So, it is proposed that initially Cs2.5SiW12040 was composed of elementary particles of Cs3HSiW12040 linked together by H4SiW12040 which acts as a binder. Upon thermal treatment, H4SiW12040 was decomposed into its oxide components which were majoritarily large cristallites of W 2 4 0 6 8 , , while the particles of Cs3HSiW12040 were not modified. 3.4. I n f r a r e d s t u d y of t h e r m a l stability a n d acidity of the salts of H3PW12040 a n d tt4SiWl2040. The infrared spectra, in the range 1200-400 cm -1 of the acids and of all their salts exhibited the same characteristic absorption bands due to the vibrations of the PW120403- or SiWl20404- anions as already well described by RocchiccioliDeltcheff et al [14,15]. The in-situ infrared study of H3PW1204o, H4SiW12040 and their caesium, potassium and ammonium salts evacuated at various temperatures revealed that these solids were stable up to 673 K since all the bands characteristic of the polyanions were still present at this temperature. Upon heating above 573 K, the terminal W-Ot band was observed to split into two bands at 980 and 962 cm -1, as already mentionned [16]. e

d

T

c

|

.....

9

9

27%

b

Z

a

sSo

eb0

460

WAVENUMBER ( c m " )

200

a-c

d

~ 23% Z

f_

200

/j, ...-i! 1060

s60

'

WAVENUMBER ( c m " )

600

Figure 2. FTIR spectra of Csl.gP sample evacuated at (a)" 298 K, (b): 323 K, (c): 373 K, (d): 473 K and (e): 573 K. In the 1800-1400 cm-lrange, as shown in Figure 2 for CSl.9HI.IPW12040 sample, a progressive evolution of the bands due to the deformation vibrations of H20 occured. The two bands observed initially at 1620-1622 and 1423 cm -1, shifted to 1720-1704 and 1463 cm -1, respectively, during evacuation between 298 and 373 K while a band at 1300 cm -1 progressively increased up to 473 K. These three bands disappeared upon evacuation at 573 K concomitantly with the modification of the W-Ot band. The same bands were observed for all the samples with, nevertheless, different intensities (for example the optical

596 density per g of the b a n d at 1720 cm -1 varied from 8.4 for H3PW12040 to 3.5 for CSl.9Hl.lPW12040). The band n e a r 1720 cm -1, often observed for heteropolyacids by IR spectroscopy [17-20] a n d inelastic n e u t r o n s c a t t e r i n g [20,21] was a t t r i b u t e d to the 5as(H20) vibration of the (H502 +) ion. The other two bands at 1463 a n d 1300 cm -1 were n e v e r m e n t i o n e d for HPA compounds b u t were observed for H U O 2 P O 4 , 4 H 2 0 [22] and were a t t r i b u t e d to t e r m i n a l H-O d e f o r m a t i o n v i b r a t i o n s of (H502+). In the range of the OH vibrations, the t r a n s m i t t a n c e of the samples was very low b u t it is possible to observe two bands a t 3440 and 3470 cm -1 which progressively decreased upon dehydration and disappeared above 473 K. W h e n NH 3 was adsorbed on the samples d e h y d r a t e d at 473 K, the intense b a n d d u e to NH4 + ions a p p e a r e d at 1427 cm -1 while the b a n d at 1720 disappeared confirming the protonic n a t u r e of the (H502 +) species. Thus, we have s t r o n g evidences for the presence of ( H 5 0 2 +) ions in all samples after dehydration above 298 K. These species were stable up to 473 K w h a t e v e r be the cation or the heteroatom, except NH4 +, for which these species disappeared between 423 and 473 K. 4. C A T A L Y T I C A C T M T Y The activity and decay behaviour of t h e different porous heteropolycompounds were compared in two reactions r e q u i r i n g strong acid sites; t h e n - b u t a n e i s o m e r i z a t i o n and the i s o b u t a n e / 2 - b u t e n e alkylation. A l t h o u g h these two r e a c t i o n s are i m p o r t a n t in the p e t r o l e u m refining i n d u s t r y , n-butane isomerization is often used as a "test reaction" since it is known t h a t this reaction requires very strong acid sites and only a limited n u m b e r of oxides are active in this reaction, under mild conditions (T = 473 K). 4.1. C a t a l y t i c activity in gas p h a s e n - b u t a n e i s o m e r i z a t i o n It was previously shown that Cs salts of H3PW12040 were active at 473 K for such a reaction, in p a r t i c u l a r CSl.9Hl.lPW12040 due to its higher r e m a i n i n g proton content and its relatively high surface area [9]. For all samples, a strong dependence of the t e m p e r a t u r e of p r e t r e a t m e n t on the initial activity and on the catalyst decay was observed. For example, the initial conversion of Csl.gP p r e t r e a t e d at 473 K, was reduced by about one half after a p r e t r e a t m e n t at 573K and the deactivation was more pronounced. F o r t h i s reason, the a c t i v i t y of the s a m p l e s was m e a s u r e d a f t e r a p r e t r e a t m e n t at 473K. In Table 2, their catalytic performances are given. Since all samples deactivated with time on stream, a deactivation coefficient R % = 100.(rate at 4 min - rate at 150 min)/rate at 4 min, was calculated and its values are reported in Table 2. First, it is be to noted t h a t all Cs salts of H4SiW12040 were inactive which may be due to the lower acidity strength of the protonic species of H4SiW12040. [23,24]. However, surprisingly, the initial activity of H4SiW12040 was three times h i g h e r t h a n t h a t of H3PW12040 but, after 35 min of time on stream, the two acids exhibited the same low activity.

597 Secondly, for a same cation, it is observed t h a t the intrinsic activity and the deactivation coefficient increased when the atomic M/HPA ratio decreased, t h a t is to say when the proton content increased. In consequence, t h r e e of the salts u n d e r study, n a m e l y CSl.9P a n d (NH4)2.4P, and K2.2P to a lesser extent, exhibited a higher initial activity t h a n all the others. Table 2. Catalytic activity in n-butane isomerization at 473 K after 4 min of time on stream. Samples

Rate of isobutane formation

Deactivation factor

9 10-10mol.s-l.m-2

9 10-8 mol.s-l.g-1

R(%)

Cs2.gP Cs2.7P Csl.9P

0.4 1.0 44.0

0.7 1.5 31.8

31 49 76

K3P K2.2P

1.7 7.9

2.6 10.4

21 42

(NH4)2.8P (NH4)2.4P

4.4 47.0

5.3 28.4

50 59

all CsxSi

0

0

93.0 266.0

3.7 10.7

H3P H4Si

23 92

C a t a l y s t weight = 0.3 g. Reaction t e m p e r a t u r e = 473 K. Total flow rate = 18 cm3.min -1. 5% nC4 in H2. Pretreatment: 2 h at 473K under flowing nitrogen. It is to be noted that among the most active samples, (NH4)2.4P was the most stable with time on stream. That may be due to its lower proton content but more probably to a lower poisoning effect of coke deposits due to the presence of mesoporosity instead of microporosity. 4.2~ Catalytic activity in liquid phase i - b u t a n e ~ u t e n e alkylation As shown in Table 3, after a p r e t r e a t m e n t performed at 333 K, the activity of the K3P sample increased with time on s t r e a m (TOS), giving rise to a high p r o d u c t i o n of d i m e t h y l h e x a n e s (DMH) and of olefins (C8=). After a d e h y d r a t a t i o n performed at 423 K, the conversion of C4 = and the selectivities t o w a r d s TMP were initially high. As generally observed in the a l i p h a t i c alkylation reaction with solid acids, the decrease of the catalyst activity was accompanied by a concomitant decrease of the selectivity in TMP a n d an increase of the selectivities in DMH and olefins (C4 = dimerization) indicating

598 t h a t with the catalyst decay the dimerization reaction is favoured as r e g a r d to the alkylation. The same sample p r e t r e a t e d at 503 K, exhibited a lower initial activity and showed moreover a fast and drastic deactivation. Table 3 Influence of the t e m p e r a t u r e of p r e t r e a t m e n t on the catalytic properties of K3PW12040 in iC4/C4 = liquid phase alkylation. Pretreat. conditions

TOS

C4 = conversion

Selectivities (%)

T (K)

(rain)

(%)

C6+C7

TMP

DMH

C8 =

333

5 35 65

2 5 17

0 0 0

80 0 0

20 70 70

0 30 30

423

5 35 65

86 76 49

18 12 9

77 62 44

5 16 36

0 10 11

503

5 35

10 0

4 .

18

12

65 .

.

.

Catalyst weight = 0.5 g. Total flow rate = 0.2 cm3.min -1. Molar ratio iC4/C4 = = 21. Reaction temperature = 333 K. Pretreatment: flowing He, at 333, 423,503 K. The activity in aliphatic alkylation of all the alkaline and a m m o n i u m salts of H4SiW12040 and H3PW12040 are compared in Table 4. A deactivation factor, R %, m e a s u r i n g the decrease of the catalyst activity after 65 min is also reported The results cannot be discussed, as for the isomerization reaction, in simple t e r m s of a m o u n t and s t r e n g t h of the acidity. Indeed, the CsSi compounds, inactive in isomerization, were active in alkylation which proves t h a t this reaction does not require so strong acid sites. It is also striking to observe that, for a same cation (except the a m m o n i u m samples), in contrast with isomerization reaction, the activity increased w h e n the atomic M/HPA ratio increased. This m e a n s t h a t the a m o u n t of acid sites is less a d e t e r m i n i n g p a r a m e t e r t h a n the morphology. The most active samples for n-C4 isomerization, (NH4)2.4P and CSl.9P, showed opposite reactivities in liquid alkylation. The first one gave rise to a high production of TMP while the second one was only initially slightly active. The m a i n difference between these two samples concerned t h e i r p o r o s i t y " (NH4)2.4P was mesoporous while CSl.gP was mainly microporous. Then, one may suggest t h a t the presence of mesoporosity is essential for the accessibility of the r e a c t a n t s to the acid sites and the desorption of the products. As a consequence the catalytic activity seems more governed by the textural features t h a n by the acidity. As a general trend, the samples which were, at the same time, active and stable for the alkylation reaction, exhibited a mesoporosity equivalent to about 40 m2.g -1.

599 It is striking to observe t h a t the most initially active samples were the most stable with time on stream in the n-butane isomerization. Table 4 C a t a l y t i c activities of alkaline and a m m o n i u m salts of H 4 S i W 1 2 0 4 0 and H3PW12040 in continuous liquid iC4/C4 = alkylation at 333 IC T O S = 5 rain. Samples

Selectivities (%)

C4 =

Deactivation

conversion (%)

C6+C7

TMP*

DMH

C8 =

R (%)

Cs2.9P Cs2.7P CSl.9P

53 72 5

7 2 0

36 8 38

39 60 32

18 29 30

75 0 100

K3P K2.2P

86 5

18 0

77 29

5 54

1 17

43 0

(NH4)2.8P (NH4)2.4P

5 57

0 35

13 59

79 3

8 3

0 0

Cs2.9Si Cs2.4Si

42 17

0 0

74 24

7 53

19 23

93 0

H3P H4Si

84 2

0 0

58 17

13 71

27 12

98 50

C a t a l y s t weight = 0.5 g. Total flow rate = 0.2 cm3.min-l.Molar ratio iC4/C4 = = 21. Reaction t e m p e r a t u r e = 333 K. Pretreatment: l h under flowing He at 423 K. *TMP distribution: 60 to 80 % for 2,3,4-TMP, 15-20 % for 2,2,4- and 2,2,3-TMP, 05 % for 2,3,3-TMP. However, the importance of the acidity and of the accessibility to the sites can not explain the behaviour of the K3P sample. Indeed, the K3P sample exhibited a high performance in TMP formation and was only slightly less stable t h a n (NH4)2.4P. This p o t a s s i u m salt, was n e i t h e r p a r t i c u l a r l y active in n-C4 i s o m e r i z a t i o n nor m e s o p o r o u s . It e x h i b i t e d p r i n c i p a l l y micro- a n d macroposities. Several e x p e r i m e n t s on different p r e p a r a t i o n s of the K3P sample have been carried out to check the reliability of this result. 5. C O N C L U S I O N In a g r e e m e n t with previous suggestions [9], the p r e s e n t r e s u l t s have confirmed t h a t whatever be the n a t u r e of the cation or of the heteroatom, the acidic salts appeared in the form of pure acid embedded in a n e u t r a l salt. The conditions of precipitation and more p a r t i c u l a r l y the initial stoichiometry determine the texture and the composition of the resulting precipitate.

500 On all the samples, the protonic species (H502 +) were evidenced by infrared spectroscopy, for dehydration t e m p e r a t u r e s varying between 323 and 473 K, which corresponds to the range where the samples are the most active in the isomerization and alkylation reactions. This suggests that the efficient sites for the two reactions are (H502 +) species since the thermal stability range of this ion corresponded to the thermal p r e t r e a t m e n t range for optimal catalytic properties. Moreover it was observed that the behaviour of the samples in the two reactions was very different. This is probably due to the fact t h a t in gas phase n-C4 isomerization, the activity was favoured on solids which presented both a high number of accessible protons and a significant surface area while in aliphatic alkylation the presence of a mesoporosity appeared to be essential for the accessibility of the reactants to the acid sites and for the desorption of reaction products.

A. Corma and A. Martinez, Catal. Rev. Sci. Eng., 35 (1993) 483. 2. J. Weitkamp, in "Catalysis by zeolites", B. Imelik et al. (ed), Elsevier, Amsterdam, 1980, Stud. in Surf. Sci. and Catal., 5 (1980) 65. F.W. Kirsch, J.D. Patts and D.S. Barmby, J. Catal., 27 (1972) 142. 4. E. Garwood and P.B. Venuto, J. Catal., 11 (1968) 175. 5. A. Corma, A. Martinez and C. Martinez, J. Catal., 146 (1994) 185. A. Corma, V. Gomez and A. Martinez, Appl. Catal., 119 (1994) 83. 7. C. Guo, S. Yao, J. Cao and Z. Qian, Appl. Catal., 107 (1994) 229. 8. T. Okuhara, M. Yamashita, K. Na and M. Misono, Chem. Letters, (1994) 1451. N. Essayem, G. Coudurier, M. Fournier and J.C. Vedrine, Catal. Lett., 34 (1995) 223. 10. M.T. Pope, Heteropoly and isopolyoxometallates inorganic cooncepts, Vol. 8, Springer, Berlin, 1983. 11. J.A. Santos, Proc. Roy. Soc.A, 150 (1935) 309. 12. M. Mizuno and M. Misono, Chem. Lett. (1987) 967. 13. J. Gregg and M.M. Tayyab, J. Chem. Soc. Faraday Trans.I 74 (1978) 348. 14. C. Rocchiccioli-Deltcheff, R. Thouvenot and R. Franck, Spectrochim. Acta, Part A 32 A (1976) 143. 15. C. Rocchiccioli-Deltcheff, M. Fournier, R. Franck and R. Thouvenot, Inorg. Chem., 22 (1983) 207. 16. K.Y. Lee, N. Mizuno, T. Okuhara and M. Misono, Bull. Chem. Soc. Jpn., 62 (1989) 1731. 17. B.W.L. Southward, J.S. Vaughan and C.T. O'Connor, J. Catal., 153 (1995) 293. 18. J.G. Highfield and J.B. Moffat, J. Catal., 88 (1984) 177. 19. A. Bielanski, A. Maecka and L. Kubelkova, J. Chem. Soc., F a r a d a y Trans. I, 85 (1989) 2847. 20. G.J. Kearley, R.P. White, C. Forano and R.C.T. Slade, Spectrochim. Acta, 46 A (1990) 419. 21. U.B. Mioc, Ph. Colomban, M. Davidovic and J.T. Tomkinson, J. Mol. Struct., 326 (1994) 99. 22. G.J. Kearley, A.N. Fitch and B.E.F. Fender, J. Mol. Struct., 125 (1984) 229. 23. Y. Yzumi, R. Hasebe and K. Urabe, J. Catal., 84 (1983) 402. 24. B.K. Hodnett and J.B. Moffat, J. Catal., 88 (1984) 253. o

o

o

o

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

601

Coupling of Alcohols to Ethers: Reaction Pathway*

the Dominance of the Surface Ss2

Kamil Klier, Qun Sun, @ Owen C. Feeley, # Marie Johansson, and Richard G. Herman Zettlemoyer Center for Surface Studies and Department of Chemistry, 7 Asa Drive, Lehigh University, Bethlehem, PA 18015, U.S.A. Coupling of alcohols to ethers, important high value oxygenates, proceeds on acid catalysts v/a general pathways that uniquely control product composition, oxygen retention, chirality inversion, and kinetics. The dominant pathway is the S~2 reaction with competition of the alcohols for the surface acid sites. This is exemplified by formation of methyl(ethyl) isobutylether (M(E)IBE) from methanol(ethanol)/isobutanol mixtures, retention of oxygen (aSo) of the heavier alcohol, and optimum rate as a function of concentration of either reactant alcohol. The SN2 pathway in the confinement of zeolite pores exhibits additional features of a near-100% selectivity to dimethylether (DME) in H-mordenite and a near-100% selectivity to chiral inversion in 2-pentanol/ ethanol coupling to 2-ethoxypentane in HZ~M-5. A minor reaction pathway entails olefin or carbenium intermediates, as exemplified by the formation of methyl tertiarybutyl ether (MTBE) from methanol/isobutanol mixtures with oxygen retention of the lighter alcohol. Calculations of transition state and molecular modeling of the oxonium-involving pathways dramatically demonstrate how the reaction path selects the products. 1. INTRODUCTION In the synthesis of higher alcohols from H2/CO , methanol and 2-methyl-l-propanol (isobutanol) are the predominant products formed over alkali-promoted Cu-based catalysts [1-4]. These alcohols can be transformed into other chemicals and fuel additives, e.g. high octane ethers. Following the discovery by Nunan et al. [5] that the sulfonic acid resin-catalyzed reaction of methanol with isobutanol yielded MIBE as the dominant product relative to DME, diisobutylether (DIBE), and MTBE, an important general question arose as to how the nature of the catalyst controls the selectivity in alcohol coupling reactions over inorganic and organic solid acid catalysts, v/z.

*This research was partially supported by the U.S. Department of Energy-PETC. @Present address: DuPont CR&D, Experimental Station, P.O. Box 80356, Wilmington, DE 19880-0356. #Present address: Exxon Research & Engineering Co., Corporate Research, Route 22E, Annandale, NJ 08801.

602

ROR'

ROR

R '=, Rn '=

R'OR'

Here R is a methyl or ethyl group, R ' is C4§ including branched alkyl, R ' - is the olefin derived from dehydration of R ' OH, and R . ' - is an oligomer of R ' -- It is demonstrated here that these pathways can be controlled by the properties of the acid catalyst utilized, i.e. in terms of nature of acid sites, strength of acid sites, shape selectivity, and choice of reaction conditions, especially temperature. 2. EXPERIM

AL

The catalyst testing was carried out in a gas phase downflow stainless steel tubular reactor with on-line gas analysis using a Model 5890 Hewlett-Packard gas chromatograph (GC) equipped with heated in-line automated Valco sampling valves and a CP-sil 5 or CP-sil 13 capillary WCOT column. GC/MS analyses of condensable products, especially with respect to O-isotopic distribution, was also carried out using a CP-sil 13 capillary column. For analysis of chiral compounds, a Chirasil-CD capillary fused silica column was employed. Experiments were carried out using isotopically labelled methanol (97% 1SO) and ethanol (98% 1SO) purchased from MSD Isotopes. Anhydrous isobutanol was purchased from Aldrich Chemical Co., Inc. and contained the natural abundances of oxygen isotopes, i.e. 99.8% 160 and 0.2% XSo. Nafion-H was obtained from C. G. Processing, Inc. and Amberlyst resins were provided by Rohm and Haas. The ZSM-5 zeolite was provided by Mobil Research & Development Corp. H-Mordenite, montmoriUonite K-10, and silica-alumina 980 were obtained from Norton, Aldrich, and Davison, respectively. 7-Alumina was prepared fi'om Catapal-B from Vista. For probing the nature of the acid sites by X-ray photoelectron spectroscopy (XPS), the samples were evacuated before gaseous pyridine was adsorbed. Excess pyridine was desorbed at 151YC, and then samples were pressed onto a sample stub under N 2 and loaded into the SCIENTA ESCA-300 instrument without exposure to air. Sample charging was minimiTed by using a flood gun while acquiring the experimental data. 3. RESULTS AND DISCUSSION 3.1. Selectivity Control by Catalysts A wide variety of solid acid catalysts has been examined using the methanol/isobutanol reaction mixture to establish activity and selectivity patterns for alcohol coupling and dehydration reactiom (Table 1).

603 Table 1 Selectivity of products (mol%) observed over 5.0 g of organic resin catalysts at 90"C and inorganic catalysts at 175"C (except as noted) at 0.10 MPa from methanol = isobutanol = 1.72 mol/kg eat/hr in a Nz/He carrier gas with a flow rate of 16 mol/kg cat/hr. Catalyst DME ..... Butenes MIBE M;FBE C s Ethers Octenes iii

Amberlyst-~-15 Amberlyst-35 Amberlyst-36 Purolite C-150 Bio-Rad AG50W NaIion-H MS

25.6 25.5 27.0 20.3 19.1 16.4

24.9 36.4 29.9 28.1 14.6 7.9

32.5 27.9 32.5 36.8 50.9 61.9

H-Mordenite a SO42"/ZrO2 Silica-Alumina MontmoriUonite HZ~M-5 y-Alumina y-Alumina b

92.0 6.6 3.4 6.4 13.1 42.1 23.1

7.6 82.3 82.3 70.7 77.1 <0.1 52.7

_.e 3.1 4.2 5.4 9.3 57.1 24.2

8' 1 4.7 5.4 6.3 5.4 1.1

8.9 5.3 5.9 8.6 10.1 12.7

'

__ 0.4 0.4 2.0 0.6 1.4 2.6 5.8 <0.1 <0.1 . . . . . . . . . . . . .

N.Afl N.A. N.A. N.A. N.A. N.A. --6.9 8.2 7.3 --

9150oC; b250oC; c... indicates not present within detection limits; dN.A. = not analyzed.

The patterns that have emerged are the following: (i) H-mordenite is a selective catalyst for dehydration of methanol to DME, while isobutanol is not significantly converted to any ethers or dehydrated to isobutene over H-mordenite under the conditions used. (ii) Sulfated zirconia is an efficient, highly selective catalyst for the dehydration of isobutanol to isobutene, with methanol dehydration to DME suppressed. (iii) Organic polymeric catalysts are generally more active than inorganic catalysts, for which higher temperatures, where the resin catalysts are unstable, are required. (iv) Although not shown, the overall order of activities for the dehydration reactiom over the inorganic catalysts investigated is SO42/ZrO2 > HZSM-5 zeolite > Hmordenite > SiO2/AI20 3 > H-MontmoriUonite >> y-AI20 3 [6]. 3.2. Nature of the Active Acid Sites High resolution XPS Nls analysis of adsorbed pyridine provides a surface sensitive tool capable of distinguishing Lewis and BrOmted acid centers. Refinement of XPS intensities by application of photoelectron cross-sections and inelastic mean free path allows for quantitative evaluation of surface acid centers. Adsorption of pyridine on a Lewis acid site increases the Nls binding energy from the free pyridine value of 398.0 eV [7] to ~400 eV due to the increase of the positive charge on the N atom in proximity of cationic Lewis acid centers. Adsorption on BrOnsted acid sites results in a larger shift of the Nls binding energy to about 401.5 eV due to the formation of a pyridinium ion.

604 Reactions exemplified above were found to ~ on catalysts containing strong Brcnsted acid centers but not on weak Lewis acid centers such as those on zirconia. Only Brr acid centers were found on sulfonic acid resin catalysts, e.g. Nafion-H. Poisoni-~ the Br~nsted centers by incremental exchange of the protons with K + yielded after full exchange an inactive catalyst for alcohol coupling, as well as for alcohol dehydration, where quantitative XPS demonstrated there was one K § acid group. Sulfate-doped zirconia was found to have both Lewis and Brr acid sites on the surface, where the sulfate groups were the carder of the protonic site. The ratio of Lewis to Brcnsted sites was found to vary depending on heat treatment, but the percentage of sulfate groups associated with a proton was found to be constant at 12-17%. 3.3. Oxygen Retention/Rejection During Ether Synthesis Over Resin Catalysts *SO-labelling studies were carried out to determine if there were a common intermediate to these two ethers. Experiments were carried out with *SO-methanolp60isobutanol = 1.0/3.2 and *SO-ethanolp60-isobutanol = 1.0/5.0 reactant mixtures such that conversion levels were <5%. Observed isotopic distributions are given in Table 2. Table 2 Percent isotopic composition (_+2 mol%) of O-containing products formed over the Amberlyst-35 catalyst (0.5 g) from the reaction of*SO-methanol (or 1SO-ethanol) with ~60isobutanol at 110~ 1 MPa, and GHSV = 14,200 ~/kg catal/hr. Isotope MIBE (EIBE) MTBE (ETBE) DME (DEE)

tsO

2 (5)

93 (96)

94 (93)

160

98 (95)

7 (4)

6 (7)

These results show that MIBE and EIBE derived their oxygens from isobutanol, while the oxygen in M'I~E and ETBE originated from methanol. This demonstrates that MIBE (EIBE) and MTBE (ETBE) were not formed from a common intermediate and that MTBE (ETBE) was not the product of isomerization of MIBE (EIBE). In each experiment, 42-45% of the reacted isobutanol was observed as isobutene, and the tertiary butyl ethers result from coupling of isobutene with methyl or ethyl oxonium. The flow of 1SO into the isobutyl ethers is consistent with surface-held methyl or ethyl oxonium associated with a Bronsted acid center undergoing an SN2 rear attack by isobutanol that is just leaving its bonded state on a neighboring sulfonic acid group by proton elimination, with H21SO being the leaving group and the MIBE (EIBE) retaining the ~60 of isobutanol. The reverse attack of isobutyl oxonium by methanol or ethanol is sterically hindered. An analogous experiment carried out with lSO-methanolp60-ethanol (1.0/1.4) gave rise to methylethylether containing 50% 1SO and 50% 160, along with DME and DEE containing > 98% *SO and *60, respectively, indicating no steric preference in the coupling of methanol with ethanol and that no si~ificant isotopic scrambling occurred. 3A. Shape Selective Reactions of Methanolflsobutanol Mixtures The probe reaction utilized a 1/1 molar mixture of methanol and isobutanol over H-mordenite, a strongly acidic zeolite comprised of linear one-dimensional channels made up of 12-ring 6.5 by 7.0 A windows [8]. There is a side-pocket system in H-

605 mordenite with an entrance 8-ring 3.7 x 4.8 A window. At a depth of 8 A into the pocket, there is a bridging oxygen that forms a pair of 8-rings, resulting in a distorted 2.6 x 5.7 A window to each side of the Y-branch. This causes comtrictiom of 2.6 A that prevent small molecules from traveling through the 8-ring channels. Small molecules can enter into the side-pockets only from the main channels. The narrower channel systems of I-t7_~M-5 (5.4 x 5.6 A and 5.1 x 5.5 A [9]) intersect to form a space of ~9 ,~ The results in Table 3 show that H-mordenite has a high selectivity and activity for dehydration of methanol to dimethylether. At 1500C, 1.66 mol/kg catal/hr or 95% of the methanol had been converted to dimethylether. This rate is consistent with that found by Bandiera and Naecache [10] for dehydration of methanol only over Hmordenite, 1.4 mol/kg catal/hr, when extrapolated to 150"C. At the same time, only 0.076 mol/kg catal/hr or 4% of the isobutanol present has been converted. In contrast, over the HZSM-5 zeolite, both methanol and isobutanol are converted. In fact, at 175"C, isobutanol conversion was higher than methanol conversion over HT_~M-5. This presents a seemingly paradoxical ease of shape selectivity. H-Mordenite, the zeolite with the larger channels, selectively dehydrates the smaller alcohol ha the 1/1 methanol/ isobutanol mixture. HZSM-5, with smaller diameter pores, shows no such selectivity. In the absence of methanol, under the same conditions at 150"C, isobutanol reacted over H-mordenite at the rate of 0.13 mol/kg catal/hr, higher than in the presence of methanol, but still far less than over HZSM-5 or other catalysts in this study. Table 3 Space time yields of products formed over H-mordenite and HZ~M-5 from a methanol/isobutanol = 1/1 reactant mixture (1.72 mol/kg catal/hr of each) at 0.1 MPa. Catalyst Temp. Space Time Yields (mol/kg catal/hr) (~ DME Butenes MIBE MTBE |

i

H-Mordenite

HZ~M-5

90 125

0.060

--

--

--

0.66O

--

--

--

150

0.830

0.068

--

--

90 125 150 175

0.005 0.071 0.261 0.185

0.001 0.169 0.339 1.086

0.012 0.350 0.134 0.131

-0.004 0.003 0.005

The differences in molecular size and shape of methanol and isobutanol need to be considered in relation to the course of the SN2 pathway in accounting for variances in reactivity. The diameter of methanol is about 3.8 ~ The isobutyl group is much larger, having a diameter of =5.0 ~ The hydroxyl group in isobutanol does not add much to its size for passage through the straight pore systems in mordenite or ZSM-5. The plausible cause of shape selectivity to DME in H-mordenite is the presence of the active protons within the side-pockets of mordenite that are accessible only to methanol. The protonated methanol molecule, a methyl oxonium ion, undergoes rearattack by a second methanol molecule entering the side-pocket from the main channel

606 in an SN2 reaction- A water molecule is the leaving group and a dimethyloxonium ion is the immediate product. Upon retur-ing the proton to the cation site, the product dimethylether molecule exits the side-pocket window to enter the main channel. The reaction path for methanol to DME based on molecular modeling [11] is depicted in Figure 1A. Isobutanol cannot pass into the side-pocket either to dehydrate to butenes or be activated to isobutyl oxonium. Furthermore, a rear-attack of methyl-oxonium in the side-pocket by isobutanol is sterically inhibited since that isobutanol would have to turn 90~ from its trajectory through the main channel for SN2 coupling with the methyloxonium. This would expose the longest dimension of isobutanol, 6.7 A, to the narrowest part of the main channel, 6.5 ~ In this configuration, the Van der Waals radii of both isobutanol and the H-mordenite channel walls would overlap, as depicted in the molecular model Figure lB. Additionally, even ff isobutanol could get into this position and attack the activated methanol in the side-pocket, the product methylisobutylether molecule, which is considerably longer than isobutanol, would have to back out of the side-pocket into the main channel, turning 90~ in order to travel out of the main channel. This is impossible within the confines of the channel.

Figure 1. Molecular graphics representations of [A] SN2 attack of a methanol molecule on a methyl oxonium ion in the side-pocket of the mordenite structure and [B] the size limitation of the bulky isobutanol molecule that prevents it from turning in the main channel to react with the methyl oxonium ion in the side-pocket. 3.5. Chiral Inversion of Ether Synthesis by the S ~ Pathway The S~, bimolecular, nucleophilic substitution reaction, according to Ingold's terminology [12], is accompanied with inversion of the configuration around asymmetric carbon- In solution phase with an acid catalyst such as H2SO4, it is genera~y believed that the synthesis of ethers from secondary and tertiary alcohols follow the SN1pattern, while synthesis from primary alcohols follow the S~2 pathway [13]. So far, however, no evidence for or theoretical analysis of surface-catalyzed Ss2 reactions has been reported. Solid acid-catalyzed S~2 reactions are demonstrated here for the coupling of chiral alcohols to form ethers, ROH + R ' OH =~ R O R ' + H20, with a very high degree of

607 chiral inversion, v/z. 97% over the HZ~M-5. This was accomplished by coupling optically active (R- or S-) 2-pentanol with ethanol. Although primary alcohols are more likely to couple through an SN2 reaction pathway than secondary alcohols, 2-pentanol was chosen because utilizing a secondary alcohol would be more general than that of coupling two primary alcohols. In addition, if the reaction involves the formation of carbenium ion, the intermediate would inevitably undergo rearrangement to form a more stable intermediate. In the case of ethanol coupling with 2-pentanol, the formation of 3ethoxypentane would be indicative of the earbenium ion reaction pathway. Under the reaction conditions employed here, the dehydrative coupling of ethanol and 2-pentanol produced diethylether and di-2-amylether as self-coupling products, 2- and 3-ethoxypentane as cross-coupling products, and 2-pentene but not ethene. For crosscoupling v/a an SN2 process, the 2-ethoxypentane product can in principle be formed either by ethanol attacking the activated 2-pentanol or vice versa, where the attacking alcohol would retain its oxygen and the resultant ether would have inverted chirality if it were formed by ethanol attacking the proton-activated 2-pentanol. On the other hand, the ether would retain ehirality of the starting 2-pentanol for 2-pentanol attack on the proton-activated ethanol. Using 1SO-ethanol, the relative contributions from these two distinct routes was determined, and the 1SO content and chirality of the cross-coupling products are presented in Table 4. The true configuration inversion I for ether formation due to SN2 attack of the activated 2-pentanol by ethanol was calculated from the oxygen isotope fraction Z in 2-ethoxypentane (Table 4, col. 6) and the ratio X/Y of the R- and S-isomers of 2-ethoxypentane (Table 4, col 5) using the relations of X = ZI and Y = Z(1- I) + (1- Z) = (1- Z1), that yields I(%) = 10011/(Z(1 + Y/X))]. Table 4 Product selectivities (+_2 tool%) (taking into account 2- and 3-ethoxypentane only) from the reaction of leO-ethanol and S-2-pentanol (160) over Nafion-H and HZ~M-5 catalysts at 100*C and 1 MPa and in concentrated H2SO4 solution at 100*C and 0.1 MPa, where the true inversion (I) was calculated by using the equation given above. Produr Selectivity (%)a R-2-EP/ S-2-EP 180/(160+ 1SO) Inversion Acid Ratio (%) (%) Catalyst S-2-EP R-2-EP 3-EP HZSM-5 Nafion-H H2SO4

14.0 32.9 33.6

86.0 60.9 64.0

0.0 62 2.4

(x/Y)

(z)

O)

86.0/14.0 60.9/32.9 64.0/33.6

89 84 82

97 77 80

aS-2-EP, R-2-EP, and 3-EP stand for S-2-, R-2-, and 3-ethoxypentane, respectively. The unreacted ethanol and the diethylether product retained > 98% of 1SO from the starting 1SO-ethanol, indicating that no isotope scrambling occurred. Data in Table 4 demonstrate that 1SO was retained in the mixed ether and ethanol attack of the acidactivated 2-pentanol via an axial SN2 rear-attack was the predominant synthesis pathway. Evidently, the shape selectivity induced by the ZSM-5 zeolite channel structure (Figure 2) plays an important role in achieving the remarkably higher configuration inversion

608 than either the Nafion-H or H2SO 4. The cross coupling of ethanol and 2-pentanol could not proceed inside either channel of HZSM-5 because once the 2-pentanol is adsorbed at the acid site, there is not enough space for ethanol to attack the activated 2-pentanol at the rear of the asymmetric carbon. However, the intersection of the channels can accommodate the transition state of the coupling reaction as illustrated in Figure 2 [14]. The transition state geometry was optimized by using the Spartan program at the R t ~ / S T O - 3 G level with the constraints that the O-C-O angle at the activated asymmetric carbon atom was linear and the two C-O bond distances (represented by dashed lines between the asymmetric carbon atom and the oxygen atoms from the attacking ethanol and the leaving water molecule) were 2.0 ,/~ This result indicates that acid sites that catalyze the dehydrative coupling of ethanol and 2-pentanol are located at the intersection of the two channel systems.

F'~mre 2. Transition state complex in the ethanol + 2-pentanol SN2 reaction activated by the proton at the channel intersection of HT~M-5 [14]. The zeolite pore structure is represented as a wire-frame section of the intersecting channels produced by the MAPLE V software package. The zeolite proton that activates the 2-pentanol molecule is marked with *.

In contrast, the Nafion-H catalyst bears a similarity to acid solution, very likely due to its flexible backbone carrying the sulfonic acid groups. The OH group of 2-pentanol is the preferred leaving group, after being activated by the surface H + and subjected to concerted nucleophilic attack by the light alcohol. The minor non-inversion path (23% over Nafion-H and 3% over HT_~M-5) can be accounted for by a less efficient carbenium ion (C +) or olefin (C z) intermediate mechanism. This minor path is corroborated by the observation of the 3-ethoxypentane side product, which could only be formed v/a carbenium ion or olefinic intermediates, over Nafion-H and in the liquid H2SO 4. However, 3-ethoxypemane was not observed with the H Z S M - 5 catalyst, demonstrating that its formation was suppressed by shape selectivity. The 3-ethoxypentane is more branched and would pass through the HZSM-5 channel at a slower rate than the 2-isomer even if formed by carbenium ion or olefm reaction at the channel intersections. Indeed, reaction over H7_~M-5 of ethanol with 3-pentanol resulted in 73% of the product being 2ethoxypentane (as a racemic mixture, along with 27% 3-ethoxypentane), while in HeSO 4 solution, 95.6% of the C7 ether product was 3-ethoxypentane [14]. These experiments demonstrate that the surface-catalyzed SN2 reaction is far more efficient than either the C + or C- pathway for the dehydrative coupling of alcohols over the solid acid catalysts tested. High selectivity to configurationally inverted chiral ethers ensues, especially in the case of the HZSM-5 catalyst, in which the minor C + or C: paths were further suppressed by "bottling ~ of 3-ethoxypentane by the narrow zeolite channels.

609 3.6. Kinetic Characteristics of SN2 Surface-Catalyzed Reactiom All evidence presented herein points to the generality of the SN2 path for the surface acid-catalyzed coupling of alcohols to ethers. The product composition, the ~SO flow from reactants to products, and chirality inversion demonstrate convincingly that mechanistic patterns of the SN2 reaction including oxygen retention, steric hindrances due to axial rear-attack of an alcohol activated by proton attachment by a second alcohol, and the accompanying conformational changes are in place. Additional features specific to surfaces, in contrast to solutiom, are (i) kinetics that display competitive adsorption of the two reactant alcohols on polymeric sulfonic acid catalysts [15,16] and (ii) shape selectivity in zeolites [11,14]. These features of the surface-catalyz~ S ~ paths have the following important consequences. ADO): On catalysts where dual-site competitive adsorption of the reactants occurs, the rate displays an opt/mum rather that "saturation" dependence on the concentration of either reactant alcohol. A comparison between the SN2 solution and surface kinetics is featured in Figure 3 to illustrate this point. Furthermore, selectivity is significantly influenced by relative strengths of the adsorption bond of the reacting alcohols. This is exemplified by the fact that methanol/isobutanol coupling over Nafion-H gives preferentially MIBE over DME (Table 1 and Ref. 15), while the Amberlyst catalysts give large quantities of DME (Table 1 and Ref. 16). The Langmuir-Himhelwood kinetic model is successful in casting the selectivity patterns in quantitative terms, and simply accounts for the preference to MIBE by relatively stronger isobutanol-to-methanol adsorption bond strength over Nation compared to the Amberlyst resins [15,16].

Figure 3. Schematic comparison of surface and solution kinetics for the SN2 path of ROH + R ' O H =~ R O R ' . Surface SN2 kinetics was taken from Ref. 16 and the solution SN2 kinetics was taken in the form of Rate = ktA]tA'I[H +]/(1 + Ir~,tA] + I~'[A' ]). Plots are of {Rate/ k[A][A'][H*]} v s [A] with Kb = lff 3 mol "1and assuming ILo' [A' ] < < 1, with A = ROH and A ' = R ' O H .

!

7 • 8O0 0r g

f

/,.

6o0 I

~" ~.. ,oo 7

[ .I "I"

~ 2o0 ~, ._. 0 0 tr

~o

7

I 0O0

2OOO

3OOO

4OOO

Concentration of ROH while keeping R'OH constant

Ad(ii): On catalysts with pores and cavities of molecular dimensions, exemplified by mordenite and ZSM-5, shape selectivity provides constraints of the transition state on the SN2 path in either preventing axial attack as that of methyl oxonium by isobutanol in mordenite that has to "turn the comer" when switching the direction of flight through the main channel to the perpendicular attack of methyl oxonium in the side-pocket, or singling out a selective approach from several possible ones as in the chiral inversion in ethanol/2-pemanol coupling in HZ~M-5 (14). Both of these types of spatial constraints result in superior selectivities to similar reactions in solutions.

610 4. SUMMARY Conclusive evidence has been presented that surface-catalyzed coupling of alcohols to ethers proceeds predominantly by the SN2 pathway, in which product composition, oxygen retention, and chiral inversion is controlled by "competitive double parking" of reactant alcohols or by transition state shape selectivity. These two features afforded by the use of solid acid catalysts result in selectivities that are superior to solution reactions. High resolution XPS data demonstrate that Brcnsted acid centers activate the alcohols for ether synthesis over sulfonic acid resins, and the reaction conditions in zeolites indicate that BrCnsted acids are active centers therein, too. Two different shapeselectivity effects on the alcohol coupling pathway were observed herein: transition-state constraint in HZ~M-5 and reactant approach constraint in H-mordenite. None of these effects is a molecular sieving of the reactant molecules in the main zeolite channels, as both methanol and isobutanol have dimensions smaller than the main channel diameters in ZSM-5 and mordenite. REFERENCF~ 1. 2. 3. 4. 5. 6. 7. 8. 9. 10.

11. 12. 13. 14. 15. 16.

IC Klier, R.G. Herman, and C.-W. Young, Preprints, Div. Fuel Chem., ACS, 29(5) (1984) 273. J.G. Nunan, C.E. Bogden, IC Klier, KJ. Smith, C.-W. Young, and R.G. Herman, J. Catal., 116 (1989) 195. J.G. Nunan, R.G. Herman, and K. KUer, J. Catal., 116 (1989) 222. R.G. Herman, in "New Trends in CO Activation," ed. by L Guczi, Elsevier, Amsterdam (1991) 265. J. Nunan, K. Klier, and R.G. Herman, J. Chem. Soc., Chem. Commun. (1985) 676. K. Klier, R.G. Herman, M.A. Johansson, and O.C. Feeley, Preprints, Div. Fuel Chem., ACS, 37(1) (1992) 236. R. Nordberg, R.G. Albridge, T. Bergmark, U. Ericsson, J. Hedman, C. Nordling, IC Siegbahn, and B.J. Lindberg, Arkiv Kemi, 28 (1968) 257. J.L Schlenker, J.l. Pluth, and J.V. Smith, Mat. Res. Bull., 13 (1978) 901; 14 (1979) 751. D.H. Olson, G.T. Kokotailo, S.L Lawton, and W.M. Meier, J. Phys. Chem., 85 (1981) 2243. J. Bandiera and C. Naccache, Appl. Catal., 69 (1991) 139. O.C. Feeley, M.A. Johansson, R.G. Herman, and IC Klier, to be submitted. E.D. Hughes and C.K. Ingold, J. Chem. Soc. (1935) 254; W.A. Cowdrey, E.D. Hughes, C.K. Ingold, S. Masterman, A.D. Scott, J. Chem. Soc. (1937) 1252. R.T. Morrison and R.N. Boyd, Organic Chemistry, 2nd Ed., Allyn and Bacon, Inc., Boston, (1972). Q. Sun, R.G. Herman, and K. Klier, J. Chem. Soc., Chem. Commun. (1995) 1849. J.G. Nunan, K. Klier, and R.G. Herman, J. Catal., 139 (1993) 406. L Lietti, Q. Sun, R.G. Herman, and IC Klier, Catal. Today, in press.

I I th International Congress on Catalysis - 40th Anniversary

611

Studies in Surface Science and Catalysis, Vol. 101

9 1996 Elsevier Science B.V. All rights reserved.

Characterization of two different framework titanium quantification of extra-framework species in TS-1 silicalites.

sites

and

L. Le Noc a, D. Trong On a, S. Solomykina a, B. Echchahed a, F. BUand a, C. Cartier dit Moulin b and L. Bonneviot a D6partement de Chimie a, CERPIC, Universit6 Laval, Ste-Foy, G1K 7P4, Qu6bec, Canada. Laboratoire pour l'Utilisation du Rayonnement Electromagn6tique b, CNRS-CEA-MENJS, Bfitiment 209d, 91405, Orsay Cedex, France.. The quantification of the extra-framework titanium species in titanium silicalites of MFI structure, TS-1, was performed using either XANES at the Ti K-edge or XPS Ti (2p) photolines. In addition, two different framework sites, [Ti(OH)(OSi)3 ] and [Ti(OSi)4], were characterized in dehydrated samples using Diffuse Reflectance UV-visible, multiple scattering analysis of EXAFS, ~H and 29Si NMR spectroscopies. 1. INTRODUCTION The need for heterogeneous catalysis using the environment friendly hydrogen peroxide as oxidant agent for mild oxidation reactions has grown considerably. More and more high surface area materials containing titanium (IV) ions incorporated in SiO2 matrices of various structural properties (amorphous gels, mesoporous materials of MCM or HMS types and microporous silicalites or zeolites) have been found active for those reactions such as epoxidation of olefins (Shell process) and hydroxylation of aromatics (Enichem process) [15]. Among these solids, only the microporous silicalites S-1 and S-2 of MFI and MEL structures containing Ti, i.e. TS-1 and TS-2 respectively, are active for the catalytic oxidation of alkanes into alcohols and ketones (oxyfunctionalization) [6,7]. Despite an increasing number of characterization studies on those materials, there has been no detailed description of the framework Ti site structure nor quantification of extra-framework species [8]. It is therefore not yet established if micropores are necessary to activate alkane oxyfunctionalization reactions or if a specific Ti environment is required. To tackle the problem, a panel of techniques were applied to a series of dehydrated TS-1 with a (Ti/Ti+Si) ratio varying from 0.4 to 4.6 %. 2. E X P E R I M E N T A L The S-1 and TS-1 were synthesized from gels containing a mixture of TEOS, tetraisopropoxytitanium, and TPAOH. H202 was only added for the synthesis of TS-1 containing 3.4wt% Ti [9]. Hydrothermal crystallization, filtration and calcination at 540~

612 followed standard recipes. The Ti loading was obtained by atomic absorption measurements. The X-ray diffraction patterns were recorded as reported earlier [9]. The diffraction angle was calibrated using Si as internal reference. The unit cell parameters were obtained from Rietveld refinement of the powder patterns. The XPS data were measured using a VG Scientific Escalab Mark II as reported earlier [10]; a binding energy of 103.3 eV for Si(2p) was chosen as reference. The UV reflectance spectra were collected on a Perkin-Elmer Lambda 5 spectrometer equipped with a reflectance attachment provided by Harricks, interfaced with an IBM computer. A titanium free silicalite was taken as a reference. The X-ray absorption data we're collected in transmittance mode at LURE (France) [9-11]. The XANES spectra were normalized following a standard procedure and the energy calibrated on the edge of a metal Ti foil. The EXAFS signal was extracted as reported earlier [9]. The structure of the [Ti(OH)(OSi)3] and [Ti(OSi)4] clusters were generated using Cerius 2 from Molecular Simulation TM. Their spatial coordinates were fed into the FEFF6 code for MS ab initio calculations. The N M R measurements were performed at RT using a Brucker ASX-300 MHz spectrometer with a 7 mm MAS probe. The zirconia rotors were filled with the same volume of powder, evacuated under vacuum at 480~ during 24 hours, filled with a dry gas and sealed. For 29Si NMR, pure oxygen was used to increase the spin-lattice relaxation rate [12,13]; the recycle time for full relaxation was then decreased from 60 (in dry air) to 15 s for S-1 and from 80 to 30 s for the [ 1.5ITS-1. A 5gs-rff2 onepulse sequence was applied while the rotor was spun at a rate of 5 kHz (no cross polarization). The filling procedure insured a proportionality within ca. +1% between the silicon spectrum area and the sample weight introduced in the rotor. This allowed the normalization of the Si and H spectra to the number of silicon atoms. The gas was changed to pure nitrogen for the proton measurements in order to decrease the spin-spin relaxation rate in comparison with the dead time of the probe (3040~ts). An 7r./2-x-rc pulse sequence was used with a recycling time of 4s. A series of 15 experiments at a z delay varying from 200 to 5000 gs was performed. The number of proton was calculated from the intensity of the signal extrapolated at time x = 0 s using a two exponential relaxation law to fit the data. Another series of experiments in static conditions was performed with x varying from 40 to 100 gs; in this case, a single exponential relaxation process was used to fit the data and calculate the number of protons. The intensity factor between the proton signal and the silicon signal was obtained from the pure silicalite sample for which the number of protons matches exactly the number of [Si(OH)(OSi)3] species, Q3 of the 298i spectrum. The number of Q3, n(Q3), was obtained from the integration of the half of the Q3 line between-90 and-104 ppm to avoid the overlap range between Q3 and [Si(OSi)4 ] Q4 lines (vide infra, Figure 5). The accuracy on the number of Q3 species was ca. 2% and ca. 3% on the proton signal intensity for both the TS-1 and the reference S-1 sample leading to a precision on the number of proton, n(H) of ca. 10%. 3. E X T R A - F R A M E W O R K Ti C H A R A C T E R I Z A T I O N AND Q U A N T I F I C A T I O N A limit of ca. 1.8% of Ti incorporation into the framework of the S-1 silicalite has been calculated by titration using Raman spectroscopy [14] or using voltametric measurements [ 15]. XRD and Rietveld calculation of the unit cell expansion led to higher limit of ca. 2.5 % [ 16]. The extra-framework Ti was identified as belonging to a segregated TiO2 anatase phase. Higher level of incorporation of as much as ca. 6 and 8 % were claimed, using either silicon

613 alkoxides with higher hydrolysis rate [17 ] or Ti alkoxides with lower hydrolysis rate [181 in their sol-gel synthesis route, respectively. In both cases, no unit cell expansion above the maximum value of ca. 5390 ./t 3 tbund bv Millini et ai [16] were reported and. no clear evidence of the absence of extra-framework Ti was brought by the authors. Thereti~re. the nmxirnum of Ti incorporation is still a controversial matter that prompted us to develop a new technique to quantify them. Recently, some of us showed that the whole XANES Ti K-edge of dehydrated [2.4]TS-I and [4.6]TS-1 samples can be fitted with a linear combination of the experimental edges of TiO2 anatase and a model organosiloxytitanium compound, HDPOSST, as octahedral and tetrahedral references, respectively [9,11 ]. The latter compound mimics a substitutional site with its [Ti(OSi)4] core. Since the edge profiles are characteristic of a given phase where the absorber atom is located [19], these fits confirmed anatase as the phase in which the Ti extraframework species develop and strongly support the hypothesis of isomorphous substitution of Ti to Si in TS-1. In addition, the fits provide a quantification of extm-ti'amework and framework species assuming that these species are properly represented by anatase and HDPOSST compounds respectively. The data clearly indicate that up to 1.5% Ti the samples are free of extra-framework species. Above 1.5%, the anatase phase develops and accounts for 25% of the Ti in the [2.4]TS-1 samples for a framework substitutional level of 1.7% in agreement with some of the literature data [ 14,15].

Fig. 1. Quantification of framework Ti sites and unit cell expansion versus Ti contents using XPS 2p transition lines ( O ), XANES profiles at Ti K-edge ( A ) and Rietveld refinement of the XRD powder patterns ( 1-1). Solid data points for TS-l prepared as in ref. [9].

614 XPS can also be used as a quantifying tool. It has been shown that a [1.6]TS-2, a titanium silicalite of MEL structure, exhibits a single Ti 2p3/2 line at a binding energy of ca. 459.8 eV assigned to framework species [10]. At higher loading, a second line appears at 458.3 eV and dominates the spectrum for a Ti content of ca. 4.2% which is assigned to octahedral extra-framework oxidic titanium phase. In TS-1, a Ti photoline associated to framework species also arises at the same position while the extra-framework sites are characterized by the same binding energy of ca. 457.7 eV as the binding energy found for anatase mixed with silica. The XPS spectra of TS-1 samples were simulated using the same set of curves as for anatase and the [1.5]TS-1. This gave excellent fits and provided a quantification of the extra-framework species reported in Figure 1. Though the quantification using XPS gives the same trend as XANES, there is a systematic underestimation of the framework species for the former technique. This is due to a low energy tail of the Ti 2p3/2 photolines present in all the spectra of low Ti content TS-1. This line may arise from i) a slight reduction under the XPS measurement conditions (15 h accumulation time) ii) TiOH groups according to a surface study of futile TiO2 [20] and, iii) minute amounts of surface segregated TiO2. Works are in progress concerning this matter. Anyhow, the unit cell expansion versus the titanium loading is fully consistent with the data of Millini et al.[ 16] and reinforces the idea that there is effectively no extra-framework species up to at least 1.5% Ti in our samples. 4. S T R U C T U R E OF T H E F R A M E W O R K SITES 4.1 UV reflectance spectra Up to now, no techniques were able to differentiate Ti tetrahedral species in glasses, Tisilicalites or Ti-MCM materials. Very recently, some of us showed that the electronic spectra in the UV range significantly differs for dehydrated TS-1 and dehydrated TiO2-SiO2 silica gel containing 1.5% Ti [21 ]. The titanium in a tetrahedral environment, according to XANES [8], exhibits a strong UV transition peaking at 40,900 cm -1. This transition falls in the predicted range of energy for the ligand to metal electron transfer (LMET) band according to the empirical optical electronegativities theory [22]. By contrast, the electronic transitions are blue shifted by about 10,000 cm ~ for the TS-1 samples and, are about 6 to 7 times more intense than in the glass. A proper dilution of the sample (here six times for the [1.5]TS-1, Figure 2) was necessary to obtain a reflectance spectra proportional to the true absorption using the Kubelka-Munk function (F(R,o) = (1-R~2/2R~. The electronic spectra are composed of three Gaussian components centered at 50,200, 44,060 and 40,290 cm -~ The relative intensity of the first two and more intense ones is constant for samples of increasing Ti loading up to 1.5%; above 1.5%, the UV absorption due to anatase in the 32,000-36,000 -1 cm range interferes with the electronic bands of the tetrahedral species. The third component of the fit increases with the titanium loading in and accounts for the differences on the low energy tail between the spectra (Figure 2, insert). The first two bands are assigned to Ti tetrahedral sites found only in the silicalite matrix while the third one to sites similar to those found in amorphous S i O 2 matrices. These spectral differences are related to different Ti-O-Si bond angle of the Ti sites. Indeed, an angle opening will shift the bridging oxygen hybridization from sp 3 to sp 2 and

615 eventually to sp. This mechanism will favor a n donation into the electronic empty "e" level of Ti in Tu symmetry. As a consequence, the non-bonding "e" level will split into a filled bonding "en" level and an empty "en*" anti-bonding level. The latter electronic level is the so-called LUMO implied in the LMET of the lowest energy ~-., (there is no d-d transition tbr a d ~! Ti 4* ion). A larger angle will lead to a higher ~ donation and a blue shift of the LMET. Along this reasoning, the dominant sites in TS-1 would be characterized by larger Ti-O-Si angles than in those in amorphous SiO2 matrices.

F 0.5 0.4

n~ u..

s

0.3 ,~ 9

_

0.3

~. 0.2 ~

9 /[15%]TS-1

~" 0.1~

~.~[0-5%]TS-1

o.oL,

v

0.2

,

0.1 0.0

50

45

40

35

wavenumber (103 cm -~) Figure 2. Diffuse reflectance UV spectra of dehydrated [1.5]TS-I experimental data and fit residue (dotted lines), gaussian fit components (solid lines); insert" details for two different TS-1.

4.2 EXAFS Ti-O-Si angle calculations The calculation of the Ti-O-Si angles are based on the multiple scattering analysis of the EXAFS signal at the Ti K-edge. For the first attempt published earlier, the only availaible experimental data to compare with the theorical data was the [2.4]TS-1 containing unfortunately extra-framework species [11]. The adjustment to the EXAFS signal of the [1.5]TS-I lead to a more reliable value of ca. 163 ~ wich is very close to the average angle of ca. 166 ~ of the HDPOSST compounds used as reference [23]. Using the same Debye-Waller factor tbr the oxygen and silicon shells to fit the EXAFS data of both the model compound and the [1.5ITS-1 sample, the number of Si second neighbors was found to be smaller than 3. This led us to the hypothesis of the existence of [Ti(OH)x(OSi)4.x] [11] or more likel.x~ [Ti(OH)(OSi)3] species [23] according to the possibility to substitute Si for Ti where Q~ [Si(OH)(OSi)3 ] naturally occurs in S-1 silicalite (vide infra). This type of so called open site is more likely to relax its structure and exhibits larger Ti-O-Si angles than closed sites of [Ti(OSi)4] type which fully undergo a compression stress from the lattice where it substitutes a smaller T site of [Si(OSi)4] type [23].

616

0.8

_1

'

I

'

I

'

I

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0.4

~

I

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'a~ -

o

0

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o

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12

I

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i

8

'

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i

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4

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.

, I,,,

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,

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,

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6

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i

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i

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---|

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!

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,

I , 2.6

!

t

! 3

,

!

,

I , 3.4

I

,? 3.8

R(A) Figure 3. X-ray absorption at the Ti K-edge of TS-I samples, experimental data (full lines) and ab inido calculation (dotted lines); b) to f) the calculation are compared with the [ 1.5]TS-1 experimental data; left hand side (dTi-O = 1.80 ,,~, dsi-O = 1.63 A, Ti-O-Si angle of c.a. 163 ~ and 3 Si second neighbors) and right hand side, mixture of sites 30% with angle of 154 ~ and 70% with angle of 165 ~.

617 Further MS calculations were performed to test the hypothesis of a mixture of sites with different average Ti-O-Si angles. For the sake of simplicity, the Ti(O)(OSi)3 and Ti(OSi)4 clusters were set as perfect tetrahedral sites with the same Ti-O and Si-O distance in both cases and the Debye-Waller factor kept identical to those found for the fit of HDPOSST. To find out what could be the angle difference between the two sites, the angle for the open site was set to the average value of ca. 166 ~ of the HDPOSST model compound while the angle was varied to obtain a satisfactory resemblance between the ab initio and the experimental data. To obtain a better fit the angle was finally decreased down to 165 ~ for the open site and found to be at least 154 ~ for the closed site with about 30-a:l0% of the titanium engaged in the latter sites. The ab initio EXAFS calculation based on the single site model reproduces most of the experimental features but does not give satisfactory results for the two first oscillations (Figure 3a). On the other hand, the double site model drastically improves the results, particularly for the second oscillation (Figure 3b). The fourier transform shows that the intensity of the second shell peak is better fit as well (Figure 3d) while the phasing match for this shell is maintained (Figure 3e and 30. Work is in progress to further improve the model.

Q4

a

v

or) t-

, I , , , , l l , , , I

90

t-

. m

100

....

110

I ....

-120 -130

Q3 integration -90

-95

-100

-105

chemical shift (ppmFI MS) Figure 4. Q3 line of the 29Si NMR spectra: a) S-l, b) [0.5]TS-land c) [1.5]TS-1; insert: total 29Si spectrum of the S-1 silicalite.

618

4.3 Characterization of TiOH groups from quantitative IH and 29Si NMR The 29Si MAS-NMR spectrum is typical of TS-1 (Figure 4, insert) [24]. The Q3 line corresponds to the Si(OSi)3OH sites. The silicon atoms linked to the titanium atoms are not resolved and resonate in the Q4 line. The latter line does not shift when the titanium loading increases; only a slight broadening (+0.2 ppm) is observed. This indicates that there are no changes in the average Si-O distances and Si-O-Si angles and, a slightly broader distribution of these parameters [25]. On the other hand, the Q3 line clearly decreases when the Ti loading increases (Figure 4). In fact, the number of silanol groups disappearing is approximately equal to the number of incorporated titanium atoms (Table 1). This observation prompted us to verify on the IH NMR spectra if this was due to the replacement of SiOH by TiOH groups. An orthorhombic titanium free silicalite (dashed line of fig. 5) exhibits two narrow proton resonance lines a and b, at 1.8 and 2.1 ppm, and a broad line c at ca. 2.5 ppm. All these signals correspond to silanol defects (7.7% of total Si) which can be removed under a strong hydrothermal treatment leading to the monoclinic, defect-free S-1. The latter only exhibits the known structured Q4 line due to the 24 inequivalent Si sites of the monoclinic structure. We found that the monoclinic S-1, exposed during one year to ambient air, evolves back to a defective structure (6% of defects) with the same a, b and c proton resonances as the original orthorhombic structure. It is therefore unlikely that these resonances comes from silicon vacancies. They are consequently assigned to terminal silanols located in different sites of the structure. This shows that the lattice relaxes by Si-O-Si bridge breaking leading to a retraction of the unit cell from 4365 to 4330/~3.

600 a

a

b/..

500 --- 4 0 0

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300

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.=_ 200

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.

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,

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i

,

,

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4

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,

~

,

1

,

,

,

,

1

,

,

1 3 2 chemical shift ( p p m / T M S )

0

-1

-2

Figure 5. IH NMR spectra of [1.5]TS-1 (solid line) and S-1 (dashed line); inserts a) signal intensity versus echo delays "r ([1.5]TS-1, solid line and S-l, crosses) and b) S-1 spectra acquired at various x delays.

619 The titanium silicalites exhibit very similar IH spectra to the defective S-1. Besides, a continuous increase of the a line at the expense of the b line is observed upon the Ti loading (Figure 5). It could correspond to either a redistribution of the silanol defects, or to the replacement of silanol defects by titanol defects at 1.8 ppm. The latter assumption is reasonable considering the ~H resonance of TiOH terminal groups in TiO2 anatase [26]. The total proton titration shows that the number of defects n(H) is always greater than the number of Q3 silanol groups (Table 1). That clearly demonstrates the existence of the TiOH groups in flae dehydrated TS-1. The predominance of the latter sites suggests that the structure preferentially relaxes by breaking of Ti-O-Si bonds instead of Si-O-Si bonds. Table. 1 Q3 defects and proton titration using Solid State NMR. static MAS Q3 Ti n(Q 3) n(H) n(H) per n(H) n(H) per -4 (10-4 mol) Ti atom (10-4 mol) Ti atom (%,+0.2) (%) (10 mol) S-1 7.7 0 3.28 3.28 3.28 0.4 + 1.4 [0.5]TS-1 7.1 0.51 3.00 3.16 0.7+ 1.4 3.08 0.7+0.5 [1.5ITS-1 6.4 1.52 2.68 3.19 0.8 + 0.5 3.10

5. C O N C L U S I O N Both XANES and XPS allow to titrate framework tetrahedral Ti and extra-framework TiO2 anatase species. There was no detectable extra-framework Ti up to 1.5wt% Ti. Open sites, [Ti(OH)(OSi)3], are the dominant framework sites according to UV-visible spectroscopy and ~H solid state NMR. They are characterized by a large average Ti-O-Si angle of ca. 163 ~ and two very intense ligand to metal electronic transitions. The closed sites, [Ti(OSi)4], are marginal at low loading. Their smaller Ti-O-Si angle than those of the open sites accounts for a higher structural stress undergone by these sites than the open sites. According to our late results during the referring process on amorphous TiO2-SiO2 system, the Ti-O-Si angle for closed sites is 30 ~ smaller compared to open sites. This is more than suggested here. Calculations taking into account this new information are in progress. 6. R E F E R E N C E S 1. 2. 3. 4. 5. 6.

M.G. Clerici, G. Bellussi and U. Romano, J. Catal., 129 (1991) 159. T. Sato, J. Dakka and R. A. Sheldon, Stud. Surf. Sci. Catal., 84 (1994) 1853. T. Blasco, M. A. Camblor, A. Corma and J. Perez-Pariente, J. Am. Chem. Soc., 115 (1993) 11806. P.T. Tanev, M. Chibwe, and T. J. Pinnavia, Nature, 368 (1994) 321. R. Hutter, D. C. M. Dutoit, T. Mallat, M. Schneider and A. Baiker, J. Chem. Soc., Chem. Comm. (1995) 163. D . R . C . Huybrechts, L. De Bruycker and P. A. Jacobs, Nature, 345 (1990) 240.

620 7.

8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26.

A. Bittar, D. Trong On, L. Bonneviot, S. Kaliaguine and A. Sayari, Proc. 9th Int. Zeolites Conf., R. von Ballmoos, J. B. Higgins and M. M. J. Treaty eds., ButterworthHeinemann, Boston, 1 (1993) 453. S. Bordiga, S. Coluccia, C. Lamberti, L. Marchese, A. Zecchina, F. Boscherini, F. Genoni, G. Leofanti, G. and Petrini, G. Vlaic, J. Phys. Chem., 98 (1994) 4125. D. Trong On, S. Kaliaguine and L. Bonneviot, J. Catal., 157 (1995) 235. D. Trong On, L. Bonneviot, A. Bittar, A. Sayari and S. Kaliaguine, J. Mol. Catal., 74 (1992) 233. C. Cartier dit Moulin, C. Lortie, D. Trong On, H. Dexpert and L. Bonneviot, Physica B, 208&209 (1995) 653. D. J. Cookson and B. E. Smith, J. Magn. Reson., 63 (1985) 217. J. Klinowski, T. A. Carpenter and J. M. Thomas, J. Chem. Soc., Chem. Commun., (1986) 956. G. Deo, A. M. Turek, I. E. Wachs, D. R. C. Huybrechts and P. A. Jacobs, Zeolites 13, (1993) 365. S. Castro-Martins, A. Tuel and Y. Ben TMdt, Stud. Surf. Sci. Catal., 84 (1994) 501. R. Millini, E. Previde Massara, G. Perego and G. Bellussi, J. Catal., 137 (1992) 497. A. Tuel and Y. Ben Tagtrit, Applied Catal. A, 110 (1994) 137. A. Thangaraj, M. J. Eapen, S. Sivasanker and P. Ramasamy, Zeolites, 12 (1992) 943. P. Behrens, J. Felsche, S. Vetter, G. Schulz-Ekloff, N. I. Jaeger and W. Niemama, J. Chem. Soc. Chem. Commun., (1991) 678. T. K. Sham and M. S. Lamnas, Chem. Phys. Letters, 68 (1979) 426. D. Trong On, L. Le Noc and L. Bonneviot, J. Chem. Soc., Chem. Comm., accepted. C. K. J~rgensen, Prog. Inorg. Chem., ed. S. J. Lippard, Wiley, New York, 12 (1970) 101; J. A. Duffy, Structure and Bonding, 32 (1979) 147. L. Le Noc, C. Cartier dit Moulin, S. Solomykina, D. Trong On, C. Lortie, S. Lessard and L. Bonneviot, Stud. Surf. Sci. Catal., 97 (1995) 19. A. Tuel and Y. Ben TMrit, J. Chem. Soc., Chem. Commun., (1992) 1578. G. Engelhardt and D. Michel, High-Resolution Solid-State NMR of Silicates and Zeolites, John Wiley & Sons, Chichester (1987). V. M. Mastikhin and A.V. Nosov, React. Kinet. Lea., 46 (1992) 123.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

621

I n f l u e n c e o f sulfur d i o x i d e on the selective catalytic r e d u c t i o n of N O by d e c a n e o n C u catalysts. F. Figueras 1", B. Coq 1, G. Mabilon 2, M. Prigent 2 and D. Tachon 1 1Laboratoire de Mat6riaux Catalytiques et Catalyse en Chimie Organique, URA 418 du CNRS, ENSCM, 8 rue de l'Ecole Normale, 34053 Montpellier Cedex, France. 21nstitut Franqais du P6trole, BP 311, 92506 Rueil Malmaison, France Abstract The selective catalytic removal of NO in oxygen rich atmospheres has been investigated in the presence of sulfur dioxide on a series of Cu catalysts. The reactivities correlated with the reducibility of Cu species determined by temperature programmed reduction with hydrogen. Without sulfur dioxide in the feed, the activity is related to the reducibility of Cu species. The addition of SO2 to the solid shifts the TPR peaks to higher temperatures. The magnitude of this effect is lower for acid zeolites such as MFI and B E A . Sulfation results in a small inhibition of the reactivity for deNOx in the case of Cu/AI203, no or little change in the case of Cu/zeolites, and a promotion of activity in the case of Cu/TiO2 and Cu/ZrO2. The oxidation of decane on Cu/TiO2 and Cu/ZrO2 is inhibited by SO2 at low temperatures, but remains close to 100% in presence or absence of SO2 on CuffiO2 above 600K. In the case of Cu/ZrO2 the addition of SO2 increases the rate of oxidation above 640 K. The positive effect of SO2 on deNOx is attributed to the promotion of a bifunctional mechanism in presence of strong acid sites.

1. I N T R O D U C T I O N The introduction, in most industrial countries, of very stringent regulations for nitrogen oxides emissions has forced the development of catalysts for the abatement of NOx emission from vehicles. If viable solutions now exist for stoichiometric gasoline engines with the so-called three-way catalysts, the future of diesel and lean burn engines may depend on the discovery of new catalysts active in the presence of large excess of oxygen [1]. In fact, the selective catalytic removal of NO in presence of excess oxygen remains a challenge. Most of the current studies involve C1-C4 hydrocarbons as reductants and zeolites as catalysts, among which Cu-exchanged MFI zeolites are considered as one of the most active [2]. The reductant shows a complex influence in this reaction: it has been thus reported that a Cu/ZrO2 catalysts are active with propene but show low activity with propane as reductant [3]. For a practical use reduction by higher alkanes would be attractive, since it would be easier to handle in a vehicle. It is well known also that higher alkanes suffer radical gas phase oxidation above 723 K. Therefore, their use requires catalysts active and selective for deNOx at lower temperatures. The mechanism of NOx elimination is still debated: a redox mechanism involving Cu ions is probable, and isolated Cu cations exchanged into MFI [4,5] or mordenite [6] have been found to be more active than CuO clusters. It must be emphasized, however, that acid zeolites exhibit good activity at high temperature, and acid mechanisms have been proposed [7-10]. In presence of Cu this acid mechanism disappears probably due to the decrease of the acidity of mordenite upon Cu exchange [6]. According to * present address: Institut de Recherchessur la Catalysedu CNRS,associ6 h l'Universit6 C. Bernard, 2 avenue A. Einstein, 69626 Villeurbanne Cedex, France. D.T aknowledgesa PhD grant from liP.

622 Sasaki et al. [8] this acid mechanism could proceed by nitration of a product of the oxidation of the hydrocarbon, followed by autoreduction of this intermediate to nitrogen. Poisoning of deNOx catalysts by SO2 could also be a problem since diesel fuels contain small amounts of sulfur compounds. Only a few studies deal with this subject [11-13]. It appears from the literature that for Cu catalysts the use of MFI as a support reduces the inhibition by SO2. Support effects also appear in the case of Co since Co/MFI is much less sensitive to SO2 than Co/ferrierite [13]. Since this support effect may be related to acidity, it becomes important, to investigate the influence of SO2 on the properties of Cu catalysts supported on SiO2, A1203, MFI, BEA and unpromoted or sulfate promoted TiO2 and ZrO2. These latter have been reported active for deNOx [14]. 2. E X P E R I M E N T A L

2.1. Preparation of the solids The supports used are listed in Table 1. Commercial supports were used except for zirconia. In that case a high surface area sample was prepared by precipitating a solution of zirconyl chloride at a constant pH of 10, then drying at 400 K. Cu on alumina, silica or zirconia catalysts were prepared by impregnating these solids with a solution of Cu acetate in distilled water. The amount of metal was altered by changing the concentration of the Cu solution. Zeolites were exchanged from their Na forms, by a solution of Cu acetate. After exchange, washing and drying at 400 K, the solids were calcined at 773 K. The chemical analyses were performed after dissolution of the solids by the Service Central d'Analyse (Solaize, France) using plasma atomic absorption spectroscopy.

2.2 Characterizations Surface areas were determined from the adsorption isotherms of nitrogen at 77 K, using a Micromeritics ASAP 200 instrument. Powder X-ray diffraction patterns were obtained with a CGR theta 60 instrument using CuKa monochromated radiation. Reducibility and the amount of Cu species were determined by temperature programmed reduction (TPR) with H2 (H2/Ar: 3/97, vol/vol). The experimental set up has been described previously [6]. Accessibility to Cu sites was determined by temperature programmed desorption of NO (NO TPD), using an experimental setup similar to that used for TPR, except the detector was a quadrupole mass spectrometer (Balzers QMS421) calibrated on standard mixtures. The samples were first activated in air at 673 K, cooled to room temperature in air, and saturated with NO (NO/He: 1/99, vol/vol). They were then flushed with He until no NO could be detected in the effluent, and TPD was started up to 873 K at a heating rate of 10 K/min with an helium flow of 50 cm 3 rain -1. The amount of NO held on the surface was determined from the peak area of the TPD curves.

2.3 Catalytic properties They were determined in a fixed bed reactor, at a GHSV of 63 000 h-1 with a reaction mixture composed by oxygen (9 vol%), n-decane (1200 ppm), NO (1000 ppm) and helium (qs 100%). For the experiments where SO2 was added to the feed, a special reaction mixture containing 20 ppm of SO2 was used. The solid could also be presaturated with SO2 by injection in the carrier gas, using a six port chromatographic valve equipped with a 200 ~tL loop. The products were analyzed continuously by sampling on line to a quadrupole mass spectrometer Balzers QMS 421 equipped with a Faraday detector, and nitrogen and CO were discriminated by analyzing on line on a GC equipped with on 13X molecular sieve column and a catharometer. While CO traces were detected among the products when Cu/MFI is used as catalyst, no CO was formed on Cu/ZrO2 either sulfated or not.

623 Table 1 Characteristics of the supports used. - Support

,

Origin

BEA MFI ~,Alumina Silica Titania Zirconia

'

Reference

Zeocat ~ t Procatalyse Rhone Poulenc Degussa home made

PB-1 PZ-3/30H GFS=200 Z175MP P25 ,

'''i'''1'''1'''1'''1'''1'''

100

9

9

Siirface (m2/g) 758 200 175 40 300

9

9

=,

9

9

9

oo

80 v

60 ~40 rj O

20

--"

u

9

o o

oo~ 0

k-,t

633

653

673

693

713

733

753

773

area Si/AI 9.85 15

Cationic form Na + H+

The catalytic conversion of NO was investigated first in absence of catalyst (blank). The results reported in Fig. 1 show that the homogeneous gas phase oxidation of the alkane starts at 650 K. No reduction of NO is observed in the homogeneous process, then the production of N2 can be ascril:r_xi to the catalytic reduction. The catalytic properties were determined by temperature programmed reaction (ramp: 2 K rain-l). The temperature was increased from 523 to 673 K and back.

Temperature (K) Figure 1. Conversions of the alkane (filled points) and of NO (open points) in the blank experiment. The catalytic properties were characterized in a simplified manner by two parameters: the maximal conversion CM and the corresponding temperature TM. The selectivity of NO conversion to N2 is always very high (> 98%). The formation of NO2 is marginal on these Cu catalysts. 3. R E S U L T S

3.1. Chemical composition The chemical analyses of the samples are reported in Table 2. The X-ray diffraction spectra of these solids do not show the presence of any Cu oxide phase with size larger than 3-4 nm, excepted in the cases of Cu(3)SiO2 and Cu(4)ZrO2 on one side, Cu(146)Na(6)FAU-10 and Cu(146)Na(28)MFI-15 on the other side, where the lines characteristic of CuO do appear with a line broadening corresponding to a particle size of about 4 nm.

624

3.2.

Redox

properties

The thermograms of Cu reduction in silica-, alumina-, titania- and zirconia-supported catalysts show only one peak, the maximum of which is reported in Table 3. The amount of hydrogen consumed by the reduction corresponds, within experimental error, to the theoretical amount required for the reaction: Cu 2+ + H 2 - - > Cu + 2H + In agreement with Shimokawabe e t al. [15], the desorption profiles of NO show two peaks at around 450 and 670 K, which correspond to Cull-NO and (CulI-O)-NO species respectively. The latter NO species usually represent no more than 10% of the total amount. Table 2 Chemical analyses of the different catalysts Catalyst

Support

Si/AI

wt% Cu

Ctt/Al

Na/AI

C'u exch

Cu(46)Na(22)BEA CufI28)Na(7)BEA Cu(48)Na(x)MFI Cu(146)Na(2~)MFI Cu(E)Al203 Cu(3)AI203 C'u(1)SiO2 Cu(3)SiO2 Cu(1)TiO2 Cu(E)TiO2 Cu(4)Zr02

BEA BEA MFI MFI A1203 A1203 SiO2 SiO2 TiO2 TiO2 Zr02

11.8 9.7 15 15.4

1.80 4.95 1.05 4.40 2.60 2.95 0.80 2.55 1.60 2.40 4.20

0.23 0.64

0.22 0.07

69 135 48 175

-

-

0.73

0.28

(~)

The comparison of the results of TPR with those of NO TPD on Cu on silica and alumina suggests that a better dispersion of the Cu species induces a higher reducibility of Cu, as reported elsewhere [ 16]. On the other hand, a strong support effect appears in the case of Cu supported on zirconia or titania, which shows the easiest reducibility in spite of a low dispersion as judged from XRD. Table 3 Maximum temperatures and hydrogen consumption during the temperature programmed reduction of Cu/oxides by hydrogen, and NO taken up by Cu atom. Catalyst

Reduction temp.

H2/Cu

NO/Cu

Cu(1)SiO2 Cu(3)SiO2 Cu(E)AI203 Cu(3)AI203 Ct(2)TiO2 Cu(4)Zr02

550 630 600 540 440 470

1.05 1.07 1.11 0.97

0.30 0.03 0.04 0.12 -

(K)

.......

In the case of zeolites (Table 4), a very broad reduction peak is observed on the sample partially exchanged, with a poorly defined maximum. On the other hand, two reduction peaks appear for Cu exchanges higher than 100%. In both cases, however, the consumption of hydrogen corresponds to the reduction of Cu 2+ to metallic Cu. A similar situation was previously reported for Cu-MOR [6]. By

625 analogy with the results obtained with Cu-MOR, we assign the low temperature peak to the reduction of CuO clusters to Cu 0 and isolated Cu 2+ cations to Cu +, and the high temperature peak to the reduction of Cu + to Cu 0. TPO experiments show that the re,oxidation of Cu occurs between 400 and 460 K for Cu/AI203 and Cu/SiO2 respectively. Reoxidation occurs then at a lower temperature and is thus faster than reduction of the Cu surface. Table 4 Maximum temperatures and hydrogen consumption during the temperature programmed reduction of Cuzeolites by hydrogen. Catalyst

1rst peak

2/kl peak

Cu(46)Na(22)BEA a Cu(128)Na(7)BEA Cu(48)Na(x)MFI a Cu(146)Na(28)MFI

630 530 710 600

630 720

(K)

H2/Cu ., 1.02 1.05

(K)

CuO/Cu2'40/100 35/65 0/100 50/50

aCu was postulated to be fully dispersed in the 10w exe "l'maged zeolites. Hydrogen uptake was difficult to quantify precisely.

The effect of sulfur addition on the TPR profiles for some of these catalysts is reported in Table 5. The saturation of the solid by SO2 shifts the TPR profiles to higher temperatures and a second reduction peak appears at high temperature. The maximum temperatures measured for this second high temperature peak are reported in Table 5 (column 6). The effect of sulfur addition depends on the nature of the support, and zeolites appear less sensitive to SO2. 300

250 v

-

Cu(4)ZrO-no SO2 2 Cu(4)Z.rO-1 S02

200

2

- - , o - - Cu(4)ZrO-6S02 2

a. 1 5 0 c-

2 "0 I

100

50

300

400

500

600

700

800

900

Temperature (K) Figure 2. Temperature programmed reduction profiles of Cu on zirconia and sulfated zirconia with increasing amount of SO2.

626 TPR experiments were performed at different amounts of sulfur on a Cu/ZrO 2 (Fig. 2). SO2 addition induces a modification of the TPR profile: the reduction peaks are shifted toward higher temperatures. The total hydrogen consumption is decreased by 50% for the first SO2 dose, then remains constant. 3. 3 Catalytic properties In the absence of sulfur, for two samples of comparable dispersion (NO/Cu = 0.04), the activity is related to the reducibility of the CuO phase. The highest conversion (98% at 670 K) is achieved on CuBEA, which appears as active as Cu-MFI. In order to investigate the effect of sulfur, the catalytic properties were determined: i) in absence of SO2 in the feed before and after addition of SO2 to the catalyst, up to a molar ratio SO2/Cu= 3.6, ii) in presence of sulfur dioxide in the feed (20 ppm). Table 5 Influence of the presaturation of the solid by SO2 on the redox properties of some Cu catalysts. 'Catalysts

Temperatures of the reduction peaks (K) No addition of SO2 Presaturation with SO2

(so2/cu = 2.85) Cu(3)Al203 Cu(3)SiO2 Cu(4)ZrO2 Cu(146)Na(28)MFI Cu(128)Na(7)BEA

'frst peak 540 630 470 600 520

2 ~ peak 720 630

lrst peak 620 660 480 650 560

2nd peak 800 590 660

Due to the small amount of SO2 in the feed, sulfation of the catalyst by the charge is very slow, and the effect of sulfur appears clearly only when the catalyst is sulfated separately. Dosing of SO2 onto the solids shows that sulfur is indeed adsorbed at the surface since Cu/TiO2 retains about 0.4 wt% S and Cu-BEA 0.1 wt% S. It is interesting to notice that the effect of SO2 clearly depends on the type of support. The different catalysts are compared in Table 6. The addition of sulfur dioxide has no effect on activity in the case of zeolites and slightly inhibits CR/AI203, but promotes NO conversion on Cu(1)/TiO2 and Cu(4)/ZrO2. The activity is practically doubled for Cu(4)/ZrO2 with a small shift of the optimal temperature. The promotor effect of SO2 increases with the amount added to the reaction medium (Fig.3). An effect of the addition of sulfur dioxide has also been observed on the oxidation of decane with an increase of the activation energy e ~ t e d for such a poisoning. This addition leads to a noticeable decrease of the rate of oxidation at low temperature, where Cu sulfate is stable, but the effect becomes negligible at about 600 K. At this temperature, the conversion of decane estimated by the evolution of the peak e/m = 57, characteristic of the hydrocarbon, is close to 100% with Cu/TiO2 catalysts in presence or not of SO2 (Figure 4). With Cu/ZrO2 SO2 inhibits decane oxidation below 640 K. At 640 K a conversion of about 60% is observed in both the presence or absence of additive and an acceleration of oxidation is noticed at higher temperatures.

627

100

.

.

.

.

.

.

.

.

.

.

.

.

.

I

|

!

!

i

!

Cu/-nq

~80

--0--

0

v

~

E

,--60

Cu/'l'iO2 (20 ppm $ 9

CurnO~(sojcu/= 0.35) curnq (sojcu = 2.as)

0

>.

"40

0 o

0 z

20

0 450

500 55O 600 Reaction temperature (K)

6,5O

Figure 3 . Conversion of NO on Cu(1)/TiO2 in presence of different amounts of 502 added to the catalyst. The reaction is performed without SO2 in the feed for the test, and in presence of 20 ppm of SO2 in the other cases. Table 6 Influence of sulfur dioxide on the NO maximum conversion (CM at.TM) over Cu catalysts. Catalyst Pretreatment by sulfur dioxide no SO2 in the feed 20 ppm SO2 in the feed so2/cu - 0 SO2/Cu = 0 SO2]Cu _ 3 TM (tC) CM(%) TM (K) CM (%) TM (K') CM (%) 710 42 '-Cu(Z)/AI203 590 44 640 37 Cu(3)/SiO2 600 14 600 34 600 66 Cu(1)/TiO2 530 41 640 61 Cu(4)/ZrO2 620 31 Cu(146)Na(28)MFI 710 55 650 62 Cu(46)Na(22)BEA 770 70 770 70 770 72 680 100 CuQ28)Na(7)BEA 680 98 -

-

-

-

4. D I S C U S S I O N Reasonable NO conversion can be achieved using n-decane as reductant. In the absence of sulfur dioxide, the catalytic activity is roughly related to the reducibility of the Cu phase of Cu ions in zeolites: the reaction temperature needed to reach 20% NO conversion parallels that of the TPR peak (Table 7). This relation also practically holds for Cu on simple oxides, therefore a redox mechanism in which reduction of Cu 2§ cations is the slow step could account for the results.

628 The addition of sulfur dioxide completely changes the activity patterns since sulfur has different effects depending on the support. On Cu-BEA zeolite, the addition of SO2 has little effect, neither on activity or reducibility. Indeed only trace amounts of S are retained by the solid in that case, which is consistent with a low interaction of SO2 with the solid. On C-k~iO2 and Ca/ZrO2, the addition of SO2 decreases the reducibility of Cu species, but increases the catalytic activity. Therefore the former relation between activity and reducibility no longer holds. These changes of activity are not related to the changes of alkane oxidation, since at the temperature of the maximum activities of Ctt/TiO2 and Cu/ZrO2, the concentrations of alkanes remaining in the reaction mixture are comparable in both eases. For TiO2 and ZrO2, it is well known that sulfation induces a strong increase of acidity [ 17] and the participation of an acid mechanism could then account for this promotion of activity. This mechanism can be described as a bifunctional process: oxidation of NO to NO2 on Cu sites, and nitration of a product of the oxidation of deeane on the acid function(8). The preparation of the catalyst must have a great influence on the activity. This has been shown by the comparison of three Cu/TiO2 catalysts prepared in different conditions: one in which titania is first treated with sulfuric acid, then by Cu acetate (denominated Cu/SO4/TiO2, containing 0.5 wt% Cu, 0.6 wt% S), one in which Cu is

100 80 t-" 0 .~_.

'~ k_

60

r O

~

40

c-

'-'

20 V

~

. . . I . . . I . . . I . . . I , i i i . i i i . . . I . . .

0

460

480

500

520 540 560 T e m p e r a t u r e (Ix')

580

600

620

Figure 4. Conversion of decane as a function of the reaction temperature on Cu(1)/TiO2 in absence of SO2 either in the feed or the solid (Experiment 1), on the same solid but in presence of SO2 in the feed (Experiment 2), and with the same Cu(1)/TiO2 catalyst, presaturated by SO2 (SO2/Cu _ 3, Experiment

3). introduced first, then sulfates using the same procedure (SO4/Cu/TiO2, containing 0.42 wt% Cu, 0.35 wt% S), and a Cu(2)/TiO2 sample sulfated by dosing SO2 in the gas phase (containing 2.38 wt% Cu and 0.49 wt% S). The comparison of these samples reported in Fig. 5 shows that the activity is higher at the higher Cu content. The slow step would then be the oxidation of NO at the Cu surface, as is commonly observed [18]. This is consistent with the observation of only trace amounts of NO2 in the effluents.

629

Table 7 Comparison of the temperatures needed to obtain 20% NO conversion and of the reducibility expressed by the temperature of the first peak in TPR by hydrogen. Catalyst Cu(46)Na(Z2)BEA Cu(128)Na(7)BEA Cu(48)Na(x)MFI C'u(146)Na(28)MFI Cu(1)TiO2 Cu(2)Al203 Cu(4)ZrO2

T2X)%(K') 640 530 670 560 500 540 570

I rst TPR peak (K') 630 520 720 600 530 590 620 .

The same behaviour has been found with Cu/ZrO2. A highly dispersed Cu phase was obtained at the surface of zirconia by reacting the support with Cu acetylacetonate [ 19]. This procedure yields an active catalyst. This catalyst was selective for N2 formation at low temperature (< 550 K), but produced only NO2 when the temperature becomes higher than 650 K. However, the same type of catalyst prepared from sulfated zirconia did not produce NO2 but selectively reduces NO to N2 whatever the temperature, with a yield of about 40% at 670 K, and a GHSV of 70000 h -1, using only 300 ppm of decane.

4~/'%

i~.s

i

i

i

I

-o80

i

,

i

I

so/c

i

i

i 'I

rns

i

i

i

I

i

i

i

I

i

i

'i

i

i

,

,

cu/sorn%

cu(2)rno + so

0

E ~

cO

60

40 0

0 z

20 0

540

560

580

600

620

640

660

680

Reaction temperature (K) Figure 5. NO conversion on CuffiO2 catalysts prepared according different procedures. The feed is the standard one containing 20 ppm of sulfur dioxide.

In conclusion, Cu on TiO2 or ZrO2 show a unique and interesting bebaviour since their deNOx activity is promoted and not intfibited by the presence of sulfur in the feed. This effect can hardly be attributed to a selective inhibition of the oxidation of decane, and is better explained by the promotion of a bifunctional mechanism involving the acid sites created on the support by the reaction of SO2.

630 REFERENCES 1. B. J. Cooper, Platinum Metals Rev. 38 (1994) 2. 2. M. Iwamoto ; "Future Opportunities in Catalysis and Separation Technology" ; M. Misono, Y. MoroOka, S. Kimura Eds.; Elsevier, Amsterdam, 1990, p.121. 3. tC A. Bethke, D. Alt, M. C. Kung, Catal. Lett. 25 (1994) 37. 4. S. Sato, Y. Yu-u, H. Yahiro, N. Mizuno and M. Iwamoto, Appl. Catal. 70 (1991) L1. 5. Z. Chajar, M. Primet, H. Praliaud, M. Chevrier, C. Gauthier and F. Mathis, Appl. Catal. B 4 (1994) 199. 6. B. Coq, D. Tachon, F. Figueras, G. Mabilon and M. Prigent, Appl. Catal. B 6 (1995) 271. 7. H. Hamada, Y. Kintaichi, M. Sasaki, T. Ito and M. Tabata, Appl. Catal. 64 (1990) L1; H. Hamada, Y. Kintaichi, M. Sasaki, T. Ito and M. Tabata, Appl. Catal. 64 (1990) LS. 8. M. Sasaki, H. Hamada, Y. Kintaichi and T. Ito, Catal. Lett. 15 (1992) 297. 9. J. O. Petunchi and W. IC Hall, Appl. Catal. B 2 (1993) L17; J. O. Petunchi, G. Sill and W. K. Hall, Appl. Catal. B. 2 (1993) 303. 10. IC Yogo, M. Umeno, H. Watanabe and E. Kikuchi, Catal. Lett. 19 (1993) 131. 11. M. Iwamoto, H. Yahiro, S. Shundo, Y. Yu-u and N. Mizuno, Appl. Catal. 69 (1991) L15. 12. E. Kikuchi, K.Yogo, S. Tanaka, M. Abe, Chem. Lett. (1991) 1063. 13. Y. Li and J.N. Armor, Appl. Catal. B 5 (1995) L257. 14. H. Hamada, Y. Kintaichi, M. Tabata, M. Sasaki, T. Ito, Chem. Lett. (1991) 2179. 15 M. Shimokawabe, N. Hatakeyama, K. Shimada, K. Tadokoro and N. Takezawa, Appl. Catal. A 87 (1992) 205. 16. S. D. Robertson, B. D. McNicol, J. H. de Baas, S. C. Cloet and J. W. Jenkins, J. Catal. 37 (1975) 424. 17. K. Arata, "Advances in Catalysis", Academic Press, San Diego, 1990, p.165, and references therein. 18. K. A. Bethke, C. Li, M. C. Kung, B. Yang and H. H. Kung, Catal. Lett. 31 (1995) 287. 19. G. Delahay, B. Coq, E. Ensuque and F. Figueras, Catal Lett. in press.

j.w. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 1 l th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

631

CoPt C l u s t e r s in N a M o r d e n i t e s as C a t a l y s t s for S C R o f NO• L. Gtmerrez, A Ra~ootta,A. Boix and J. Petunchi.

lnstituto de Investigaciones en Cat~li~s y Petroquimica- INCAPE (FIQ,UNL- CONICET) Santiago del Estero 2829 - 3000- Santa Fe- Argentina.

A new material based on Pt and Co exchanged in NaMordenite for the selective catalytic reduction (SCR) of nitric oxide with methane in the presence of excess oxygen is studied. The incorporation of 0.5% weight of Pt and 2% weight of Co to the zeolitic matrix alter calcination and reduction on ~ flow for 1 h yields a solid converting 100% of NO to N 2 and, simultaneously, 100% o f C H 4 t o CO: with a CH4/ NO ratio = 3 and 2% of oxygen in the feed at 450~ When the oxygen concentration in the feed varies, the NO conversion goes through a maxinmm for 2% at 450~ The incorporation ofl~ also promotes Co reduc~flity; 1% is reduced to Co ~ in the monometallic sample and 13% in the bimetallic sample. XPS results reveal that in the calmed samples Co 2+is at exchange position and, after being reduced, there appear thinly dispersed Co ~particles and exchanged Co 2+ions. A greater reducibility and a shift ofthe maxima in the temperature-programmed reduction profiles suggest a Pt-Co interaction. In order to get an efficient catalyst for nitric oxide abatement it is necessary that the highly dispersed Co ~ and Pt ~ particles and the Co 2+and I-I+ ions at exchange positions be in intimate contact inside the mordenite channels. 1. I N T R O D U C T I O N Since the pioneer work of Li and Armor (1,2,3), who found that methane could be used as a selective reductant ofnirrogen oxide over Co-ZSM5, several solids have been used with the same purpose (4,5,6,7). Metals such as Co, Rh and Pt, exchanged in ZSM5 (8), were similarly studied. In the absence of O: the three metals have a high selectivity but when 25% of oxygen is added to the feed, platinum is non-selective, the NOto N 2conversion being only 3% with 25% ofNO 2production. Under these conditions Rh-ZSM5 presented 25% of NO conversion with approximately 10% of NO: at 450~ Kikuchi and Yogo (9) in their studies ofSCR of NO using different protonic zeolites found that activity and selectivity are functions of acidity. The same authors have reported that the incorporation of Cra or In to HZSM5 notably increases the NO to N: conversion. Li and Armor (10) also studied Ga-HZSM5 and they found that this solid presented a high selectivity to NO reduction, even at high temperatures. The NO to N 2 conversion kept constant over 450~ Pd-based catalysts on different acid supports including HZSM5 were studied by Loughran and Kesasco (11) and Nishizaka and lVfisono (12). These authors found a behavior ~ to that ofGa-HZSM5, atm~uting the activity and selectivity ofthose solids to a bifunctional metal-acid action All materials effective for nitric oxide abatement are sensitive to the presence of water in the feed; however, Inui et aL (13) have recently reported that a protonated Co-contained silicate was effective with CH4for NO reduction in conditions equivalent to diesel engines exhausts, the NO to N: conversion on this material not being altered by a I-I:O concentration in the feed ofup to 10%. This work studies a new catalyst based on Co and Pt exchanged in mordenite. It presents an

632 ~ l y s i s ofthe effect o f Pt on activity and selectivity in the selective NOreduction, and finally discusses an atten~t to characterize active sites. 2. EXPERIMENTAL

Catalysts Prepara#orL Catalysts were prepared by ionic exchange starting from NaMordenite Zeolon Si/AI = 6. Monomet~c solids were prepared using for the exchange ~xted solutions of Co(NO3)2 and Pt(NI-13)4('NO3)2,respectively. Bimetallic solids were obtained by successive exchanges beginning with Pt salt for PtCoMordenites, and with Co salt for CoPtMordenites. The same support was exchanged with a solution ofNH4NO 3to prepare NH4Mordenite which was then exchanged with the Pt and Co salts, respectively, to finally obtain PtCoH-Mordenite. The exchange time for each ~rnple was 24 h at room temperature and pH = 5-6, the zeolite/solution ratio employed was 2g/dm3 and all the solids were filtered, washed and finally dried at 120~ for 8 h. Monometallic ~mples with Co were pre-treated according to the standard method, i.e., heating at 2~ in oxygen flow up to 110~ with isotherm of 2 h, then up to 210 ~ with the same isotherm, and finally at 400~ for 8 h. Mono and bimetalfic solids with Pt were calcined following the method reported by Callezot et aL (14) consisting in r a j ~ g temperature at 0.5~ in oxygen flowup to 350~ keeping this temperature for 2 h. The PtCoNH, M samples was then treated in oxygen at 450~ for 8 hs. The Co/celite sax~le was prepared byinvregnating Co(NO3)2 on the low surface silica (17 m2/g). X-rayDiffraction Diffractogramswere obtained with a S h i m a ~ XD-D 1 insmmaent with monochromator using ChtKa~ radiation. It was operated in continuous scan mode at 0.5 ~(20) min". TPRExperiments. They were performed with 0.100g of catalyst with an Okura TP-2002 S insmnnent, with a heating rate of 8~ using as reducing gas a 4.8% H 2in argon flow. All samples were calcined according to the standard method, prior to each e x p ~ t . XPS Spectra. They were obtained with an Esca 750 Shimadzu instnnnent, using MgKcx radiation This spectrometer is driven by a computer system (Escapac 760)~41ich allows both the accunmlation and the processing of data. The spectra were obtained at room temperature. Platinum was analyzed in the area corresponding to the binding energy of the 4forbitals, because the signals corresponding to Pt4d overlap with the AUGER signals o f Na. The (Co/Si)s atomic ratios were calculated using the area under the Co2p, Pt4f and Si2p peaks, the Scofield photoionization cross sections the mean bee paths ofthe electrons and the i n s t a l function was given the ESCA manufacttlre:l'. The samples were subjected to reduction treatment with H 2in the reaction chamber directly attached to the ESCA spectrometer. CatalyticMeasurements. The reaction was carried out using 0.500 g of catalyst placed in a fixedbed flow reactor. This was a 12 mmi. d. tubular quartz reactor with an internal thermowelL The typical reacting mixture consisted of 1000 ppm of NO, 1000 ppm of CH, and 2-10% of 05, balanced at 1 atm with He (GHSV = 6500 h-~). The catalytic activity and the composition ofthe reacting gases were analyzed with a Varian 3700 chromatograph with 2 cokum~, one with a 5 A molecu~~ sieve, and the other containing Chromosorb 102. The NO x conversion (C~o) was calculated from the N, production. Selectivity was defined as S = Cso/Ccm where CcH4is the CH 4to CO x conversion. For all the solids analyzed, the only reaction products were N: and CO s, N:O was not detected and the carbon balance was always better than 98%. The conversion reported was determined after reaching steady state (usually after 1 h oftime-on-stream). A reaction cycle was defined as follows: The sample was kept on stream increasing temperature flora 200~ to 500~ staying 1 h at each ten~erature. After the highest temperature was reached the solid was treated in O 5 o.v. at 450~ and cooled up to 200~ in He flow. In some cases, after pre~treatment, the samples were reduced in pure ~ flow for 1

633 h at the desired temperature (350 or 500~ Table 1. Prepared Solids. Samples

'

Metallic Content

% b

Nonfiml unit cell composition

Co

Pt

Pt(o.5)M

2.51 -

0.50

Pt(o.~)Co(2.o)M

2.05

0.56

Co(2.o)Pt(o.5)M

2.96

0.57

Co,.63Pto.ogNa3.56(AdO2)7(si0 2)41

Ptr

1.73

0.50

Pt0.09Co o.9oH,.o,(AI 0 2)7(Si02 ),u

COo.o/Celite

3.00

-

Co(2.o)M

C o,.3Na4.4(AIO2)7(SiO2)4, Pto.09Na6.r~(A10,)7(SiO2)4,

Pto.o9C~

(SiO2)4,

aSubscripts represent theoretical % in weight of cation. bDeterminedby atomic absorption and CPS. 3. RESULTS In agreement with the results obtained by other authors, CoM was active for SCR o f NO xwith CH4 (3,5). In fact, Co~2.0)Mwith 42% of B.E.C. (Base Exchange Capacity) presented a ma3dmum NO to N: conversion of 31% at 450~ (Fig. 1), the selectivity at this temperature being 0.74. When the sample was reduced at 350 ~ and 500~ a decrease of NO conversion was observed (Table 2). The incorporation of 0.5% ofl~ (2.25% of B.E.C.) increased the initial activity ofthe oxidized sample. How100 ever, after various reaction cycles the NO to N 2 cow 90 version stabilized at a value close to that of Cot2.0)M but with a si~ificantly lower s e l ~ . 80 When Ptt0.~)Cot2.0)M was reduced at 350~ for =~ 70 1 h in I-[2flowthe result obtained was totally different. In effect, a si,.~nificant increase of NO conversion and selectivity was observed (Fig. 1). Besides, the sample remained stable during the successive reaction cycles and 30 also through time at constant temperature. The order in 9 ~ 20 which the exchange was performed did not have signifilo / cant importance on the N: production but increased the CH4to CO, conversion in the whole temperature range. 0 350 400 450 500 5 5 0 In order to explore the effect of acid sites on TEMPERATURE, ~ bimozllic solids a san~le was prepared fromthe annnonic Fig 1. Nitric Oxide ( I-l, II) and methane form of the mordenite. In this case, a decxease of about ((3 ,O ) conversion as a function of tem30% in the NO conversion was observed when it was perature on : Co<2.0)Moxidized (open evahmted under the same conditions (Table 2). symbols) and Pt<0.~)CO<e.0)Mreduced (close symbols) at 350~ samples. Reaction condition: NO: 1000 ppm, CH4:1000 ppm, O:: 2%balance to 1 atmHe. GHSV: 6500 h-i

634 Table 2 Co, Pt and CoPtMordenite Catalytic Behavior' Catalyst

TR b

(oc)

C~Io c

(%)

Tmax d

Cci.14 c

(oc)

(%)

T Lf

Sg

(oc)

Pt(o.s)CO(2.o)M

-

35

450

78

450

0.45

Pt(o.5)Co(2.o)M

350

74

450

56.0

450

1.32

Co~2.o)Ptto.~)M

350

80

500

100

375

0.80

Pt~o.5)COo.o)HM

350

57

500

92

400

0.62

Co~2.0)M

-

31

450

42

475

0.74

Cot2.o)M

350

14

450

44

475

0.31

Co(2.o)M

500

25

450

52

450

0.48

Pt(o.5)M

-

13

500

34

520

0.38

Pt~o.5)M

350

-

550

68

525

-

Ptto.5)M.Co<2.o)M h

350

19

450

69

400

0.27

Co~3.0)Celite

3 50

10

5 50

60.0

5 25

0.17

aReaction conditions: GHSV: 6500 h-1,NO: 1000ppm, CH4:1000 ppm, O2: 2%. bReduction temperature 1 h with I~. cMaximumNO to N: conversion, qVlaximumNO conversion temperature. "CH4 to CO2conversion ~ight-offtemperature. gCso/Cca4selectivity, hMechanical mixture. The catalytic activities expressed as turnover frequencies at 3 50~ shown in Table 3 further confirm that the active sites have a different activity according to their environments.

Table 3. Turnover frecuencies of PtxCoMordenitey for the nitric oxide reduction.'

Catalyst Co<2.o)M

Pretreatment

TOF x 10 3 ( S"l)b

Standard

0.1

Pt
St + 1 h I-I: flow at 350~

0.7

Pt~o.5)Co~2.o)HM

St + 1 h H e flow at 350~

0.2

aReaction conditions: NO: 1000 ppm, CH4:1000 ppm, O:: 2% balance to 1 atm He, temperature 350~ bTOF = Number of NO molecules converted per total cations per second.

Monometallic catalysts containing the same % of Pt as bimetallic ones were also evahmted. The calcined samples presented low activity, the maximum NO to N 2 conversion being only 13%, whereas when it was reduced at 350~ in I-I: flow, no NO conversion was observed, presenting activity only for CH 4 with 02 combustion. The results obtained suggest that Pt has a promoting effect on Co<2.0)Mactivity for SCR o f NO• In order to determine whether this effect needs the intimate contact of cations or if'it was only a consequence of the presence of metallic PL the following experiment was conducted: a mechanical

635 ngxture of both calcined monometallic solids (Co(2.0)M and Pt(0.5)M) was prepared. This mixture was reduced at 350~ for 1 h in I-I:flow and then catalytically evak~ed under the above desc'n,bed conditions. The results obtained are shown in Fig. 2. Note that the maxinmm NO conversion is lower than 20%, a 100% C H 4 conversion being obtained at 500~ This suggested that an intimate contact of active sites inside the mordenite channelis necessary to obtain the desired effect.When the effect ofO 2 concentration on the feed flow was studied the results shown in Fig. 3 were obtained. Note that while the NO conversion on Pt(o5)CO(zo)Mreduced at 350~ presents a maxinmm for 2% oxygen, the other samples keep a constant conversion value from the said concentration. It should be remarked that the reduced bimetallic sample presents a 60% NO conversion, even for 10% 02 in the feed flow. The same behavior was observed when the Co(2.o)Pt(o.5)Mwas tested. The effect of CH4 concentration on the SCR of N O on reduced bimetallic samples was also studied. Fig. 4 shows the results for Co(2.0)Pt(0.5)M Note that for a C H 4 / N O r a t i o = 3 , a 100% NO conversion at 450~ was obtained. It is well known that H20 is a reversible inh~itor (15) for the removal of NO by methane over CoZSMS. In order to determine whether a sinn~ effect was observable in the bimetallic samples, 3% H20 was cofed to a flow containing 1000 ppm of NO and 1000 ppm of CH4 with 2% O: at 450~ The results obtained for Co(z0)Pt(0~)M are summarized in Fig. 5. It is worth noting that an almost 60% decrease was verified in the NO conversion, while in the methane conversion the decrease was about 26%. Upon removal of H20 from the feed, the catalyst totally recovered its activity. 100

100

90

90 80

ao N

70

~

60

0 ~

z ~

0

70

60

so

50

40

40

30

9 r,j

30

20

20

10

10

0 350 400 450 500 550 TEMPERATURE, ~

Fig 2. Selective reduction of NO with CH 4 on Co(2.0)Mand Pt(0.~)Mmechanical mixture. ( I ) NO to N2,(O ) CH, to CO 2 conversion. Reaction conditions as in figl.

0

1

~

,

l

.

t

.

L

0 2 4 6 8 10 02 CONCENTRATION, %

Fig. 3. Nitric Oxide conversion as a function ofoxigen inlet concentration on Pt(o~Co(o,)M reduced ( O ) ; CO2.0M oxidazed ( I ) and CO2.0M reduced ( [21 )samples. NO: 1000 ppm, CH4: 1000ppm, balance to Iatm He, Temperature 450 ~ GHSV: 6500 h -~.

In the search for a better understanding ofthe Pt promoting effect, the different solids were studied by the XPS and temperature-programmed reduction techniques.

636 Table 4 XPS data ofmonometallic and bimetallic mordenites Catalyst

Treatment

Co(2.o)M Coo.0/Celite Ptt0.~~Co~2.0)M Pt,.0)Cot2.0)M

Co2p3~ Binding energies (eV)"

Co/Si b

Co 2+

Co ~

eRiC. 400~

783.1

-

0.09

red. 427~

783.4

-

0.05

talc. 400~

781.5

-

0.22

red. 380~

-

778.2

0.08

talc. 350~ red. 380~

783.1 783.3

778.3

0.08 0.06

talc. 350~

783.1

-

0.08

red. 380~

783.3

778.3

0.08

'The binding energy values are referredto Si2p(102,6 eV)except in SiO2where the referencewas Si2p (103.8 eV). bSurface Co/Si ratio calculatedfrom XPS data experimental. Surface Si/AI ratio for all the exchanged mordenites was 6.1. The binding merg~es (B.E. s) for mono- and bi- metallic solids are mamaafiz_edin Table 4.In Co~2.0)M the B.E. value obtained for Co2p3a corresponds to 783.4 eV, which is not modified after a prolonged reduction (4 hs) at 427~ This signal is assigned to ions of C@* located at exchange positions within the zeolitic structure. This value differs in about 0.8 eV fiom the one reported for Co :~ in Y-zeolite by Zsoldos et al. (16). It must be noted that there exists a difference of approximately 2 eV with respect to the binding energy 100 measured for the cobalt oxide supported in silica (Co/ 95 Celite). 90 In the calcined bimetallic sample the exchanged C@+ signal was observed, and after a static reduction (in the XPS chamber) in 1-I2atmosphere at 380~ the Co ~ 80 (783.2 eV) and Co ~ ( 778.3 eV) signals coexisted. 75 The signal of Pt4foverlaps that o f A12p; conse70 quently, it was not possible to accurately determine the 9 ~ 65 signals corresponding to Pt in the bimetallic samples with 60 0.5%. When sa_mpleswith higlaer contents(1 and 5%) were studied, binding energies ofPt ~ (73.1 eV) and Pt ~ 55 (71.9 eV) were observed on both calcined and reduced 501000 2 0 0 0 3000 solids. The surface Pt/Si ratios were 0.003 and 0.016 for the bimetallic samples with 1 and 5% o f l ~ respectively. C H 4 CONCENTRATION, ppm The TPR profiles ofthe monometal~c and bimetalFig 4. Nitric o x i d e S ) and methane (O) lic solids are shown in Fig. 6. Co~2.0)M,after calcination at conversion as a function of methane inlet 400~ presents a broad peak centered about 400~ concentration on Co. 20)Pt(0~)M reduced sample at 450~ N~: 10()0ppm, 02: 2% which can be assigned to a fraction of the cobalt supported as oxide on the external surface ofthe mordenite. balance to 1 atmHe. GHSV: 6500h ~ A smallerpeakwith a maximum at 539oC, probably due i

,

i

,

1

637 to the Co2+reduction at exchange positions, can be observed in the high temperatures region The degree of reduction of cobalt in the temperature range under analysis was lower than 2% in accordance to what was reported by other authors for CoY (16) and CoZSM5 (18). The monometallic Pt(0.5)Mprestated two maxima in the low tempera10 f H2Olout) H20(out) tures region (194 and 269~ and a 0{~ , i i i sharper peak at 560~ This type of ,~0 1(30 1,~0 TPR profile could be attn~outed to Pt TIME, rain exchanged at different sites in the Fig 5. Effect of water addition on the NO (11) mordenite stmomr The He consumpand CH4(O )conversion on Co(2.o)Pt(0.5)M tion calculated from each TPR curve reduced sample. NO: 1000 ppm, CH4: indicated that about one hydrogen 1000ppm, 02: 2%, H20: 3%balanceto 1 atm molecule was consumed per Pt ion as He. Tenverature 450~ GHSV: 6500 h -~. shownin Table 5. Thisimpliesthat most of the Pt remained in divalent state after calcination to 350~ The TPR profile ofthe bimetallic samples presented two clearly defined mmdma, one at 237~ and the other at 56 I~ The latter one was shifted to higher temperatures with respect to the m o n o ~ c samples, Note that Co ~increased approximately 10 times with respect to the monometallic sample. The diffxactogramsobtained from the mono and bimetallic mordenites after standard pretreatment were coincidem with those of the o d g h ~ Na-Mordenite. Additional data on surface area measurements and pore vohune confirm that there is no break ofthe support struca ture during the c a i r n treatment. No ~m~alcorresponding to cobalt or platinum species were observed in the diffractograms of calcined or in the reduced samples (up to 700~ which would indicate that the particles f o m ~ have a smaller sizethan the detection limit of the instrument (40 A approximately). x c Diffractogram~ comparable to the ones described above were obtained for all the mordenite exchanged smnples at'-. ter being under reaction stream for several hours.

I~9~)i~..~.,,...I--120 (in)

)-120(in)

-I

0

200 400 600 800

__'f_

4. DISCUSSION

The results obtained in this work (Fig. 1, Table 2) dearly show the promoting effect ofthe Pt incorporated to Fig.6. TPR profille of. a)Pt(o~)Co(zo)M, CoMordenite for the NO selective reduction with methane b) Pt(0 5)M, c) Co(2.o)M , d) Co(3.o)/Celite. on streamwith excess oxygen. The reduction ofthebimetallic sample with ~ at 350~ yields a solid which converts 100% of NO to N 2 at 450~ in a reacting stream with a CHffNO ratio of 3 and 2% ofO 2(Fig. 4). TEMPERATURE, ~

638 Table 5 Reducibifity of Co PtMordenites. Catalysts

T

Metal loading gmol b Co Pt

399

IV(pt+Co)

0.44

0.010

0.19 34.54

0.004 0.68

1.59

0.608

0.80 3.27

0.309 0.08

3.01 0.96

0.080 0.030

0.48

0.015

42.5

Co(2.0)M C o(3.0)/Ce]ite

l~mol H2 ~

539 388

50.8

-

-

2.6

194 Pt(o.5)M 269 560 237 Pt(o.5)Co(z.o)M

34.7

2.0

561 170 Pt(o.5)CO~z.o)HM

29.3 500

2.6

*Temperature at maximum TPR peak_ bMetal loading per 100 mg of solid. ~H2 consumption from TPR profile per 100 mg of solid.

The catalyst remained stable for a long while (over 50 h) under reaction conditions at the above temperature. In order to wy to understand this interesting behavior, let us first focus our attention on the b e h ~ o r ofmonomet~c cobalt and platinum mordenite samples. Sevel~ authors (3,6) studied catalysts for SCR based on Co zeolite~ In a recent review (19) it was reported that Co-Femerite is the most active catalyst for the said reaction followed by Co-ZSM5 and Co-Mordenite whereas Co supported on other materials such as TiO2, A1203, SiO2-A]203.0rsilicalite presented no appreciable activity in the 400 - 500~ ten~erature range, under the conditions studied, Le. NO: 1600 ppm, CH4:1000 ppm, 05: 2.5%, GHSV: 30,000h ~. The results obtained in this study for monometallic samples (Co(z0)M and Coo.0/Celite) (Fig. 1, Table 2) are in general agreement with the data previously descried. At this point a con~arison ofthe b ~ c mordenite against the most active Co monometallic zeolite could be relevant. Li and Armor (20) reported that Co-ZSM5 (4% Co) gives c.a. 100% of NO conversion at 7500 h-t GHSV and 400 ~ for a ratio CH4 / NO=2.6 and 2.5% of O2. The Pt~0.,)Co~z0)Min this study gives the same conversion at 450 ~ and CH4/NO ratio of 3 but with only half the CO loading. On Co-Fer 42 % of NO conversion was reported (7) at 450~ but at 30,000 h-~ GHSV and a CH4/ NO ratio of 0.6. ~ s sample a~ears more active than the Pt~0.~)Co(0.2)Mbut it also has over twice as much Co loading A study under more similar conditions coul improve the comparison Generally speaking, the results obtained with Pt monometallic samples (Table 2) are also in agreement with those previously reported for comparable systems. In fact, Li and Armor (3) and Butch and Scire (8) found that Pt exchanged ZSM5 was completely unselective towards the nitric oxide reaction with methane in the presence of oxygen. The latter authors did not find NO to N 2conversion between 200 and 600~ bu/t they found that a maximum 25% of nitric oxide converts to NO r In contrast, in the absence of oxygen Pt-ZSM5 converts 93% of NO to N 2 at 450~ (8). Besides, with more reactive hydrocarbons (ethene, propene and propane) Pt-ZSM5 and platinum on other supports have been re-

639 ported to be active in the SCK of N O at lower temperatures (21,22,23). Allthese results suggested that the catalytic behavior observed in the bimetallic samples is not a consequence of the addition of the individual effects of each cation. The results of the mechanical mixture of both monometallic solids (Fig. 2) in which the Co/Pt ratio was the same as in bimetallic samples confirmed the above hypothesis. Another important concept derived fromthe above is that both cations need to be in intimate contact inside the channel ofthe mordenite structure to generate an effective SCR catalyst. Sina-larresults were found by Laughran and Resasco (11) in their study of PdH-ZSM5. These authors, in their studies on Pd catalysts on different acid supports, found that the promoting effect ofacidi~ on Pd in the case of I~ZSM5 requires an intimate contact inside the zeolite channel This was not necessary when the support was SO4/ZrOr The promoting effect being discarded by the simple mechanical addition ofthe cations, let us analyze the results derived from the XPS and TPR studies of the bimetallic samples to try to shed light on such an effect. In the calcined bimetallic samples Co 2+is located at exchange positions according to the XPS results (Table 3). C02P3a binding energies of Co2+were the same in mono and bimetallic samples. The Co ~ signal detected in XPS spectra ofthe reduced samples is showing the promoting effect ofplatinum on the cobalt r e d u ~ (compare Co(z0)M and Pt(0.~)Co(2.0)Min Table 3). The B.E of Co ~ in PtCoM is comparable to that of Co~ This some how differs l~omthe results reported by Zsoldos et aL (16) who, in their studies of Pt(7.0)Co(z0)YZeolite, found a shift of approximately 0.7 eV in the Co 2+B.E. between the mono and the bimetallic sample. They attffouted this effect to the CoPt interaction and the subsequent formation of bimetallic particles. In our system this possl~oilitycannot by totally ruled out due to the low content of surface Pt of the sample used in this study (Surface Co/Pt > 20). The Pt4t~a signal obtained with samples with high platinum loading (1 and 5%) si~owsus that despite the low heating rate [0.5cc/min] used during the calcination ofthe sample, the decomposition of the tetramin platinum ion provoked the reduction ofthe ion at superticial level TPR results (Fig. 6) show that this did not occur in the bull The promoting effect of Pt on Co reducibility in the bimetallic sample is confirmed through TPR studies(Fig. 6 and Table 5). In the 150- 350~ temperature range, 6% ofcobalt was reducedto metallic Co against 1% in the monometallic one (Table 5). From what is descn%ed above, it can be concluded that after reduction at 350~ the solidthat the r e a ~ nfix~e inilially "sees" has Co and metallic Pt particles, highly dispersed in the zeolite matrix together with Co2+cations and protons generated during reduction at exchange positiong However under reaction conditions, due to the oxygen excess, Pt ~ and Co ~ could be oxidized to their oxides while Co+2and the proton are leit intact. The remaining Na+ cations from the parent mordenite are also present in the solid but the remits here obtained do not allow a disoLsfion ofthe possible effect of such cations on the catalytic behavior ofthe bimetallic mordenite. This complex system yields a very active, selective catalyst for the NO selective reduction. Several concurring causes could be respona~olefor such an effect" A promoting effect ofacidi~ in the formation ofNO2, fin~'larto the one suggested for GaHZSM5 (9,10) and PdHZSM5 ( 12,13). NO 2has been proposed by various authors as an important intermediate in the reaction scheme. Pt can also contffoute to the formation ofNO 2(vide supra). The results obtained with Pt(0.~)C%.0)HM(Tables 2 and 3) would suggest that there exists an opfinann ratio ofthe acid function. In such special environments Pt would decompose NO as suggested by Ansell et al. (23) for SCR of NO with C3H6on Pt/A1203. Any N20 likely to be formed, would be easily removed bythe Co sites (24). In this hypothesis, the 02 concentration of the feed could present an optinmm as the one shown in Fig. 3. Note that for the case of Co monometallic samples, such maxinmm is not observed (Fig. 3). - A sinergetic effect cant occur among the three types of sites in intimate contact, with Pt, activating CH4 the acid site favouring NO 2formation and Co2+conm'buting with its well-known capaci~ to adsorb NO (25). Further research is necessary in this system in order to elucidate whether any ofthe previous pro-

-

640 posals, so far speculative, are feasible. 5. CONCLUSIONS - PtCoMordenite is an active, selective catalyst for SCR ofNO xwith C H 4. This new catalyst proves to be stable in the temperature range studied (Figs. 1 and 5). - Pt promotes the Co redueibilityin Mordenite (Fig. 6 and Table 5). - Clt~ers of Co and Pt particles highly dispersed with Co 2+at exchange positions, and an adequate acid function mint necessarl~ be i ~ ~ e contact withinthe Mordenite channels to gmerate an efficientcatalyst for the N O with CH4 abatement (Figs. 1 and 2 and Table 2). - The presence of water vapour in the feed provokes an important decrease of activity which is totally restored when water is suppressed (Fig. 5). ACKNOWI.,EDGMENTS. This work was carried out within the framework of the Joint Study Project on Heterogeneous Catalysis: CENACA (National Catalysis Center) - JICA (Japan lmemational Cooperation Agency). Thanks are also given to CONICET and Universidad National del Litoral for partial support, and to Elsa Cnimaldifor her help in editing the En~li~ REFERENCES _

2345678910111213141516171819202122232425-

Y. Li and J.N. Armor, U.S. Patent 5,149,512 (1992). Y Li andJ.N. Armor, AppL Catat B1 (1992) L31. Y. Li and J.N. Armor, AppL CataL B2 (1993) 239. K Yogo, M. Umeno, I-[ Watanabe and E. Kikuchi, CataL Letters 19 (1993) 131. E W~zeL G.A. Sill andW.K. HaIL J. of Catal 149(1994)239. J. Vassano, E. lVlir6and J. Petunchi, Appt Catat, in press (1995). Y Li and J.N. Armor, J. of Catal 150(1994) 376. IL Butch and S. Scire, Appt Catat B3 (1994)295. E. Kikuchi and K. Yogo, CataL Today 22 (1994) 73. Y Liand J.N. Armor, J. Catal 145(1994) 1. C.J. Loughran andD.E. Resasco, Appl. CataL,in press(1995). Y. N ' ~ and M 1W~sono,Chem Letters (1994) 2237. T. Inui, T. Kaabayashi and S. Iwamoto, Carat Letters 27 (1994) 267. P. Cmllezot,A. Alarc~-Diez, J./~LDalmo~ A.J. Renouprez and B. ~ J. CataL 39(1975) 334. Y. Li, P. Battavio andJ. Armor, J. CataL 142(1993)561. Z. Zsoldos, G. Vass, G. Lu and L. Guczi, Appl Sure Scq_78(1994)467. J.C.Kimand S.J.Woo,ApptCatat, 39, (1988) 107. J.M. Stencel,V.U.S. Rao, J.ILD~ K.a.Rhee, A,G.I~ere and 1LJ.De Angelis, J.Cata[,( 1983)84,109. J.N.Armor, Catal. Today, 26(1995) 147. Y.Li and J.N.Armor, in H.E. Cuny-Hyde and 1LEHowe (Editors), Natural Gas Conversion II.Elsevier, Amsterdam (1994) 103. G. 7hang, T. YamagudlL H. Kawakanfi and T. Suzuki, App[ Catat B 1 (1992) L15. A. Obuchi, A. Olfi, M, Nakatma'a, A. Ogata, K ~fiz~o and R Ohuchi, Appl Catal. B2 (1993)71. G.P AnselL S.E. G o h m ~ J. Hayes,/LP Walker, IL Butch and PJ. Millington, CAPOC3, Brussds (1994) 295. Y. Li andJ.N. Armor, AppL Catal B3 (1993)55. Y.Li, T.L.Slager and J.N.Armor, J.Catal, 150(1994) 388.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.)

11th International Congress on Catalysis - 40th Anniversary

641

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

Decomposition of Nitrous oxide over ZSM-5 Catalysts Freek Kapteijn a, Guido Mul a, Gregorio M a r b ~ a, Jos6 Rodriguez-Mirasol b and Jacob A. Moulijn a a Industrial Catalysis, Department of Chemical Engineering, Delft University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands Fax: +31 15 278 4452 Email: [email protected] b Department of Chemical Engineering, University of Malaga, Spain

Abstract Comparative kinetic and in-situ DRIFT studies of the N20 decomposition over Co-, Fe- and Cu-ZSM5 have been performed. The implications of the presence of 02, CO, NO, H=O and SO= on the catalyst activity and stabilitiy and on the mechanism are evaluated. 1. I N T R O D U C T I O N Nitrous oxide has received increasing attention the last decade, due to the growing awareness of its impact on the environment, as it has been identified as an ozone depletion agent and as a Greenhouse gas [1]. Identified major sources include adipic acid production, nitric acid and fertilizer plants, fossil fuel and biomass combustion and de-NOx treatment techniques, like three-way catalysis and selective catalytic reduction [2,3]. Much interest is noted in the development of catalysts that decompose nitrous oxide into its elements, at rates and conditions that are compatible with the production sources. 2 N20

-.+ 2 N2 + 02

(ZXrH~

= - 163 kJ/mol )

(1)

Catalysts include oxides, mixed oxides (perovskites) and zeolites [3]. The latter, transition metal ion-exchanged systems, have been shown to exhibit high activities for the decomposition reaction [4-9]. Most studies deal with Fe-zeolites [5-8,10,11 ], but also Co- and Cu-systems exhibit high activities [4,5]. Especially ZSM-5 catalysts are quite active [3]. Detailed kinetic studies, and those accounting for the influence of other components that may be present, like 02, H20, NO and SO2, have hardly been reported. For Fe-zeolites mainly a first order in N20 and a zero order in 02 is reported [7,8], although also a positive influence of 02 has been found [11]. Mechanistic studies mainly concern Fe-systems, too [5,7,8,10]. Generally, the reaction can be described by an oxidation of active sites, followed by a removal of the deposited oxygen, either by N20 itself or by recombination, eqs. (2)-(4). k1

N20

+

*

"-'+

N20

+

O*

----)

N2

+

O*

N2

+

02

(2)

k2

+

*

(3)

642

k3

2 O*

~

02

+

2*

(4)

k_3

On kinetic grounds eqs. (2) and (3) have been proposed for Fe-Y and Fe-Mor [7,8], while eq. (4) is generally used for pure and mixed oxides to account for oxygen inhibition. The deposited oxygen, often denoted as extralattice oxygen (ELO), has special properties. It catalyses the 02/1802 oxygen exchange reaction over Fe-ZSM-5 at room temperature [6]. From N2180 studies over Fe-Mor it appeared that this oxygen was exchanged with lattice oxygen and with N20, eq. (5) [5]. The latter can be interpreted as a nonproductive step (2). N2180 +

O*

~

N20

+

180*

(5)

For each ELO two Fe ions are involved [12,13], but with N20 about 1 oxygen is deposited per 4-5 Fe ions [ 10]. On the other hand more than eight times the maximum ELO capacity could be exchanged with lattice oxygen [5], indicating that this deposited oxygen readily looses its identity. The isolated nature of the ions in the zeolite framework resembles the dilute solid solution systems, extensively studied two decades ago for the N20 decomposition [ 14,15], for which mechanistic proposals have been made [15], that may apply to zeolites, too [3]. In essence, the oxygen moves from the TM ion to a matrix oxygen and either reacts with another such oxygen to 0 2 , or with an newly incoming oxygen on the TM-ion. Simultaneously the identity of the oxygen may be lost by exchange with the framework oxygen. If the peroxyoxygen is not mobile the reaction is limited to the direct vicinity of the TM ion, which is suggested by the value of four exchangeable lattice oxygens per Fe-ion. The presence of reducing agents, e.g. CO, may enhance the oxygen removal, eq. (6) [16], thereby accelerating the N20 destruction. k6 CO .4O* ----) CO2 + * (6) The present study focuses on the comparison of the behaviour of three catalysts with different performance, viz. Co-, Cu- and Fe-ZSM-5, with emphasis on the kinetics and mechanism of the reaction. 2. E X P E R I M E N T A L 2.1. Catalysts. Cu-, Fe- and Co-exchanged ZSM-5 zeolites have been used as catalysts for the N20 decomposition. ZSM-5 zeolite (SIO2/A1203 =37.2) in the sodium form (ZEOCAT PZ-2/40 Na; Chemie Uetikon) was ion-exchanged, under vigourous stirring, using aqueous solutions (pH between 5.5 and 6) of Cu(I/) acetate (4.0 mM), F e ( ~ sulfate (3.7 mM) or Co(H) acetate (39.7 mM) at 293, 343 or 323 K, respectively. The zeolites were then filtered and washed thoroughly with deionized water at room temperature before drying at 383 K overnight. The metal content and exchange stoichiometries were determined by ICP-AES and AAS. The data are given in table 1. 2.2. Experimental setup and procedures. The experimental setup for N20 decomposition consisted of a gas mixing section, a reactor and a gas analysis section. A quartz fixed bed reactor of 5 mm I.D. was used, containing 20 mg of catalyst (106-212 mm) diluted with

643

Table I Catalysts used Sample Metal loading (wt.%) Cu-ZSM-5 4.0 Co-ZSM-5 1.64 F~-ZSM-5 1.3

Na exchanged (%) 11 5 3

Metal exchange level (%) 130 73 98

180 mg of SiC (106-212 mm), to assure plug flow, at a total pressure of 2.5 bar. The SiC diluent did not contribute to the N20 decomposition at the reaction temperatures studied. Prior to each run, the catalyst was subjected to calcination by heating the catalyst in He at 30 K/min to 923 K and maintaining this temperature for one hour. Subsequently, the temperature was decreased to the desired value and the feed mixture was passed over the bed. Temperature and feed composition were varied in a random order in the experiments. Generally, 40 to 50 min after a change of conditions the conversion levels were constant and considered as the steady-state. At least five analyses were averaged for a data point. The product gases were continuously analyzed for NO and NO2 using a chemiluminescent analyzer, and discontinuously for N20, N2, CO, CO2 and 02 by GC equipped with a thermal conductivity detector and an electron capture detector, specifically for the N20 analysis, using a Poraplot Q column and a molsieve 5A column for separation. Conditions. The influence of temperature, partial N20 and 02 pressure, space time W/FN2o, and gases like H20, SO2, CO and NO on the decomposition of N20 over the catalysts were studied. The temperatures varied between 625 to 873 K. The inlet partial N20 pressure ranged from 0.05 to 0.2 kPa, the 02 pressure from 0 to 10 kPa and the space time from 1.5x105 to 11.0• g*s/mol. The total gas flowrates were between 1 and 5 ml(STP)/s. NO, CO or SO2 were added in molar ratios of 0-2 with N20 (at 0.1 kPa). The H20 pressure amounted to 13.6 kPa by passing the feed gas through a saturator kept isothermal by means of a water bath. Parameter estimation. Integral reactor behavior was used for the interpretation of the experimental data, using N20 conversion levels up to 70%. The temperature dependency of the rate parameters was expressed in the Arrhenius form. The apparent rate parameters have been estimated by nonlinear least-squares methods, minimizing the sum of squares of the residual N20 conversion. Transport limitations could be neglected. 2.3. Characterization. In-situ diffuse reflectance FTIR (DR.Wr) experiments were carried out with undiluted samples of the zeolites in a Spectratech DRIFT cell and a Nicolet Magna 550 spectrometer. Most experiments were carried out in a flow mode, passing 0.84 ml/s of a gas mixture containing inert (He, Ar or N2) and N20, NO, CO or mixtures of these gases continuously through the cell at atmospheric pressure. Each spectrum was recorded by addition of 256 scans and a resolution of 8 cm ~. 3. RESULTS An impression of the activity of the different catalysts is given in figure 1. The activity order Cu>Co>Fe corresponds with literature [4]. The N20 pressure dependency for Co-ZSM-5 is given in figure 2. Due to the integral reactor behaviour the relation between conversion and partial pressure shows a curvature, but the reaction order equals 1 for Co, and slightly lower

644

1.0 t X(N20)

1.0

Cu-ZSM-5 .'"

0.8 f

0.8 X(N20)

,/

9

,, ""

0.6,,

"

Co-ZSM-5

,r

9

/ 650

0.2 L

5, 700

800

750

T/K

773 K

0.6 o4 9l

0.4

0.2'

Co-ZSM-5

o.~u . u .......

/~I--~...=..J"

.........".......

./" ..-"" o;s' o.o~

' '

~

.4................... . ..........

' o:, 0.1 ....

748 K 723 K

698 K o.~s 0.15 ....

0.2

P(N20 ) / kPa

Figure 1 Conversion as a function of Figure 2 Partial pressure dependency of temperature 0.1 kPa N20 and space time Co-ZSM-5 at a space time of 1.52"105 1.44" 10 s g.s/mol g.s/mol

Table 2

Apparent activation energies (kJ/mol) and reaction orders only N20 3%02 CO/N20=2 NO/N20=l.5 all data *) order n *) Co 106 110 115 134 106+15 1 Cu 132 170 187 140 138+ 17 0.88_-_-~.11 Fe 173 187 78 182+31 0.79!-0.15 *) 95% confidence limits values are found for Fe and Cu. The fitting results of apparent reaction orders and activation energies 08 ~ ~ ~-''~'743 K for the different experiments are given in table 2. The presence of 02 hardly affects the reaction over Fe- and Co-ZSM-5, but it inhibits slightly for the 0.4i / F~ZSM-5 / Cu system, although this effect seems to level off at higher oxygen concentrations (figure 3). The 0.21- . ~ 7 ~ - - : _ -. . . . . ---._ ~ apparent E,, for Cu increased by nearly 40 kJ/mol. 773 K In the presence of NO the reaction is enhanced for ~176 2 4 6 8 10 Fe, and not affected for Co and Cu (figure 4). The P(02) / kPa reaction over Fe-ZSM-5 was accelerated Figure 3 The effect of oxygen on the tremendously; at temperatures where the N20 conversion at 0.1 kPa N20 and decomposition does not noticeably take place high space time W/FN2o=2.87* 105 g.s/mol, conversions were obtained by addition of NO. In the product mixture of the Fe and Co samples NO2 was observed. With Cu-ZSM-5 no NO2 was found. Figure 5 gives a product composition for Co-ZSM-5. Nearly all the converted NO and N20 results in NO2 and N2. Similar results are obtained for Fe-ZSM-5. Addition of CO also enhances the N20 conversion, by about a factor of two for Co and tremendously for Fe (figure 6). For Cu a maximum in the N20 conversion appears as a function of the CO/N20 ratio in the feed. This maximum shifts to higher values with increasing temperature (figure 7). The apparent activation energy for Co is hardly altered, for Fe it decreased nearly 100 kJ/mol, while for Cu it is increased by 50 kJ/mol. The presence of H20 inhibits the reaction and gives rise to deactivation for Cu and Fe. 1.0

xl.o,

"-

645

Concentration / ppm 70o

1.01 k

0.8

/

~9

600

l X(N20) ( 0.6~-

Fe-ZSM-5 (673 K) ,~

0.4~

~

500

~'IF~''-'-'-'-'-~

?--.L-" ....... L

30O

"

2O0

Co-ZSM-5 (723 K) ............ A

100

"

....

......

015 . . . .

z

..........

1'.0 . . . .

NO

400

.. Cu-ZSM-5 (673 K)

.

T / o~1

Co-ZSM-5

11 . . . .

2.0

N2 NO 2

, .~~~_.,_~__/?

~io

O~ ,

700

molar NO/N 2~ ratio

750

800

T/K

Figure 4 Effect of NO on the N20 Figure 5 Product composition for a conversion at 0.1 kPa N~O and space NO/N20 feed mixture over Co-ZSM-5 at time W/FN2o= 1.52"105 g.s/mol. Included 0.1 kPa N20 and space time W/FN2o= 1.52" 105 g.s/mol. is the NO conversion (open symbols). 1.0

1.01 L

Cu-ZSM-S (673 K ~ _ 0.8 ~//I i "~'--...._ X(N20)

i

k

jr"

./

!

f

<673 K)

o~ ~ c o - z s M - s ' (; 6 9 3

~176 ....

o', .....

,!o

X(N20) 0.6

~"-~.

...~Fo-ZSM-S

. . . .

0.4 0.2

K) ;,

Cu-ZSM-5

0.8

. . . .

~.o

molar CO/N20 ratio

~

/

/x----~....~.....~..~.

//

648 K

!

" o'.s . . . .

10 ' ' ' ' ;.s . . . .

i.

2.0

molar CO/N 2~ ratio

Figure (i Effect of CO on the N20 Figure 7 Effect of CO on the N20 conversion at 0.1 kPa N20 and space conversion over Cu-ZSM-5 at 0.1 kPa time W/FN2o= 1.52"105 g.s/mol. N20 and space time W/F~2o= 1.52"105 g.s/mol. no H20

m

13.6kPa H20

1.o 0.8 X(N20) 0.6 0.4 0.2 o.o

Co-ZSM-5 793 K

m m

Fe-ZSM-5 833 K

Cu-ZSM-5 728 K

Figure $ The effect of water on the N20 conversion. Shown are levels before, during and after 13.5 kPa H20 (1 h) addition. Conditions 0.145 kPa N20 and space time 1.54" 105 g.s/mol

In figure 8 the activity of the catalyst before and after the addition of 13.5 kPa water is indicated for the three catalysts. After 10 h the activity of FeZSM-5 had declined by a factor of three and was still declining. Only the Co-ZSM-5 turned out to be quite stable, although it suffers the strongest from the inhibition.

SO2 nearly completely deactivated the Cu-ZSM-5, resulted in an inhibition for Co-ZSM-5 and an enlargement of the N 2 0 conversion over Fe-ZSM-5 (figure 9). Both the Fe and the Co systems returned to their original activity after removal of the SO2, this took several hours.

646 Characterization of the Cu-ZSM-5 catalyst 1.0 by in-situ diffuse reflectance FTIR 0.8 e-ZSM-5 spectroscopy after treatments in CO, air and X(N20), N 2 0 is presented in figure 10, the CO o.6! adsorption in figure 11. 0.4 The lattice vibration at 938 cm -], related to --co-z the oxidation state of the copper [17], indicates that after treatment in N20 or air the copper is in the 2-r oxidation state, %:0 . . . . o:s . . . . ,:o . . . . ,: . . . . zo while under CO it is reduced to +1 as is molar S ( : ~ O ratio indicated by the shift to 963 cm ]. This reduction could also be achieved by Figure 9 The effect of SO2 addition on the treatment in N2 at 770 K. CO (5 kPa) N20 conversion at 0.1 kPa N 2 0 and space strongly adsorbs at this catalyst at 450 K time W/FN2o= 1.52"105 g.s/mol. (figure 11 ). 2.0: A b

1.4

s

o r a n

1.2

0.5

1.0-

0.0 0.8" 1200

lT=450 KI

li00

1000

Waven~

900 (o'n1)

"~i 2137

1.0 84

NK)]

c

e

2157 .'.

1.5

800

Figure 10 DR]FT spectra of Cu-ZSM-5 at RT after reducing and oxidative treatments.

"tl

219s Y' 22oo

1V

\~. 2109

~. /

21"oo

Wavenumbers (cm")

26oo

Figure 11 DRIFT spectra of Cu-ZSM-5 at 450 K in 5 kPa CO in Ar. Dashed line is a Fourier self-deconvolution spectrum.

The deconvolution indicates bands at 2157 2132 1.2 2195(vw), 2176(m), 2157(vs), -2137(m) and 2109 (w) cm -l. At lower Cu loadings A b 1.0 the absorption bands are better resolved and s o at temperatures below 320 K the major r 0.8 bands are at 2176 (m), 2150 (s) and 2134 b (m) cm "1 , while above 500 K the only band a n 0.6 is at 2157 cm ]. These results correspond c excellently with work from Spoto et al. e 0.4 [18]: 2157 cm ] represents Cu*-CO, 2176 2170 2150 2130 2110 and 2151 cm 1 Cu+-(CO)2 and 2190 and wavetu'rt~s (cm1) 2164 cm ~ Cu*-(CO)3, and indicate that the di- and, possibly, tricarbonyls can be Figure 12 Change in CO vibration on Cupresent even at high temperatures. From ZSM-5 upon readsorption of water at RT in the temperature dependency of the CO 10kPa, after CO reduction at 770 K. absorption a heat of adsorption of 40-60 kJ/mol is estimated, which is larger than that for supported copper oxide [ 19], but corresponds to that for Co-ZSM-5 at low temperatures [20]. During the flow experiments it was noticed that below 400 K trace amounts of water, present in the gas mixture, adsorbed at the samples,

647 apparent from the rise of the broad band in the hydroxyl region, 3000-3800 cm ~, and a characteristic band of adsorbed water at 1608 cm "~. Simultaneously the 2150-2157 cm -l band decreased and a 2132 cm 1 band grew (figure 12), clearly due to an interconversion and not due to a shift. This band has also been observed in other studies [18,21,22] and most pronounced in CO adsorption at 77 K on HZSM-5 [23]. Therefore, it has been ascribed to condensed or 'liquid-like CO' [18], but no relation with water is indicated. Co-adsorption of charge donating ligands like nitrogen containing molecules like NH3 shift the carbonyl band considerably [24]. Due to this charge donation the copper ion can back-donate more to the CO resulting in a shift to lower values. Like ammonia, also water can coordinate, up to six molecules, to the copper ion in ZSM-5 [22]. So, we explain the observed shift to a charge backdonation to the CO under influence of coordinated water. The 2109 cm l band is ascribed to the presence of small copper oxide clusters in view of the overexchange level [ 19,25]. 1am

A

S O

0.4 84

1891

r b

/_ /

1874 ~1~9

1758

0.2,

n

0.2, 2000

0.4-

1900

1800

~~t~

(cm~)

,,,

1700

Figure 13 DRIFF difference spectra after NO (5 kPa) adsorption on Co-ZSM-5 pretreated in 10 kPa CO.

c e

20()0

1800

Waver~ll'i~B (ClT)"1)

Figure 14 DRIFF difference spectra at 575 K after NO adsorption (5kPa) on a reduced (by CO) and on an oxidised (by N20) Fe-ZSM-5 sample.

Since CO hardly adsorbs at the other catalysts NO has been used for their characterization. Figure 13 shows the spectra for Co-ZSM-5 and 14 for Fe-ZSM-5. In these figures the spectra of the samples in N2 at the same temperature have been subtracted for a better resolution. The Co spectra are not different whether the sample was pretreated in CO or in N20. Three clear bands are observed, corresponding to Co2+-NO at 1928 cm ~ and Co2+-(NO)2 at 1806 and 1891 cm ! [26]. The latter bands decrease with temperature but are still visible at 575 K. The adsorption of NO on reduced and oxidised samples gave different results. After reduction and addition of NO three clear bands arose ascribed to Fe2+-(NO)2 (1758 and 1849 cm ~ ) and FeE+-NO (1874 cm ~ ) [27]. The dinitrosyl bands decrease with time due to the oxidation of F e E+ t o F e 3+. Since the Fea+-NO absorption band is also present at 1874 cm 1 this band does hardly change. A band at 2133 cm l increases with time and represents NO2 formed during the oxidation reaction by NO. This has been observed more for Fe- and Cu-zeolites [28,29]. This band is also observed after oxidation in N20 (figure 14), but not in air. Apparently N20 deposits an oxygen that can react with NO to NO2 according to eq. (7), but molecular oxygen can not. Also after air pretreatment the NO adsorption gives rise to two overlapping bands at --1900 and 1878 cm "1, instead of one.

648 5. DISCUSSION The Cu-, Co- and Fe-ZSM-5 catalysts are active systems for the decomposition of N20, but their behaviour differs with respect to conditions and gas atmospheres. They all seem to obey a (nearly) first order dependency towards prq2o, which can he rationalised by the two step kinetic model given by eqs. (2) and (3). A step like eq. (3) is quite well feasible, since the TM ions in ZSM-5 can be coordinated by several ligands simultaneously [18,22]. The resulting rate expression is given by eq. (7). k~N r (7)

r = 1+ k I / k 2 "P~¢~o

where Nr represents the number of active sites (mol/g). This model predicts no inhibition by oxygen, as found for Fe and Co. The inhibition over Cu-ZSM-5 can be ascribed to the molecular [30] or dissociative adsorption of oxygen (eq. (4) backward). The latter is supported by the NO dissociation properties, e.g. [5,28]. Whether the presence of pairs or clusters of copper ions [17,30] is essential remians to be answered, although it is a tempting explanation. The ratio kl/k2 in eq. (7) equals [O*]/[*], so determines the state of the active sites. For kl/k2 >>1 the rate determining step is eq. (3) and the sites are oxidised, while for kdk2 <<1 the reverse holds. By addition of CO the removal of surface oxygen, according to eq. (6), is very fast and the active sites are virtually in the reduced form, *, so in fact the rate of eq. (2), kl NT p rq2o, is measured. The increase in N20 conversion indicates that for Co (2 times increase) -50% of the sites are oxidised, O*, while for Cu and Fe this is estimated to be more than 95%, in agreement with results for Fe-Y and Fe-Mor [7,8]. This means that for Fe and Cu the rate determining step is eq. (3) and the observed activation energy in only N20 concerns mainly Ea2. On the other hand in CO/N20=2 mixtures the observed activation energy represents the Eal for Co and Fe. Due to the strong adsorption of CO on reduced copper sites this is more complicated for Cu. At low CO pressures step (6) enhances the removal of oxygen, while at high CO pressures the adsorption blocks sites for N20 decomposition, eq. (8). The resulting rate expression and approximation at high CO pressure is given by eq. (9).

CO

+

*

~

*CO

kl N r P N2o r=

(8)

kl N r P N~o :=~

1 + klPN2° + KsPco k6Pco

KsPc°

for CO/N20 >>1

(9)

with: E °~' = E ~ l - a n ~

So, in the latter case the apparent activation energy is increased by the heat of adsorption of CO, amounting to about 40-60 kJ/mol as calculated from the IR experiments. Hence, for both the Co and the Cu samples Eal is slightly larger than Ea2 (table 2) while for iron Ea~ is considerably lower. All these values are compatible with values reported in the literature for Fe-zeolites [6,7,10,11 ] or dilute solid solutions of Co in MgO [31 ]. The kinetic and IR results with NO indicate that, like CO, it can remove the oxygen from the

649 surface of the Co and Fe catalyst, too, thereby forming NO2. The overall reaction is given by eq. (10). Although this occurs less efficient than with CO, it is thermodynamically quite feasible [3]. The NO does adsorb on the Co catalyst, evidenced by IR, yielding a slight increase in apparent activation energy. So, although at 723 K no enhancement is observed, but some is expected at higher temperatures. Kinetically, the blocking of sites by the observed NO adsorption on the Fe catalyst is negligible compared to the enhancement achieved by the oxygen removal effect. Over Cu-ZSM-5 steady state NO2 formation is not observed. Speculation offers two explanations; either it is not formed or it decomposes back to NO. IR measurements show by the presence of the 2134 cm -~ band that NO2 can be formed on CuZSM-5 [29,32], but from TPD experiments it appears that it decomposes rapidly in the range of 600-700 K [33]. So, in our case this could mean the occurrence of reaction (10) backwards or NO2 reacts with an oxidised site to NO and 02, acting as an oxygen carrier and hence competes with N20 in this respect. This does, however, neither result in an acceleration of the N20 decomposition nor in a changing temperature dependency. The NO2 formation offers the potential use of these catalysts in nitric acid plants off-gas treatment, where about equal amounts of NO and N20 are present. The produced NO2 can be reused in the nitric acid process [3]. SO2 also increases the N20 conversion in the case of Fe, probably according to eq. (11), although no product analysis is available. Co-ZSM-5 is inhibited by the SO2 or SO3, but not irreversibly. Copper is completely deactivated, as for other reactions over copper zeolites. NO

+

NzO

-->

SOE

+

O*

--~

NO2

+

Nz

(10)

SO3

+

*

( 11 )

l0

Water exerts both a deactivating and inhibiting influence on Cu and Fe samples, while the reaction over Co is only inhibited. The deactivation of Fe- and Cu-ZSM-5 is clearly due to migration and the sintering of the active component in H20 atmospheres [34]. The Co-ZSM-5 catalyst is much more hydrothermally stable in wet gas conditions [34,35]. The inhibition by water can be accounted for in a similar way as for CO via competitive adsorption on active sites, like in selective NO reduction studies [34]. For N20 decomposition this yields an expression like eq. (12). At 793 K Kl2 amounts to about 0.7 kPa ~ .

r=

klNrPu2o kl 1 + ~ + K12PH20 k2

(12)

Evaluating the results a clear kinetic picture of the catalysts has been obtained. In the steady state the active sites in Fe- and Cu-ZSM-5 are nearly fully oxidized, while for Co only -50% of the sites are oxidized. The former catalysts operate in an oxidation reduction cycle, Fe2+/Fe 3§ and Cu+/Cu 2§ Co a§ in zeolites is hardly oxidized or reduced, but ESR studies on diluted solid solutions of Co in MgO indicate that Co3§ formation is possible, rapidly followed by a migration of the deposited oxygen to lattice oxygen and reduction back to Co 2§ [36]. For Fe-ZSM-5 such a migration has been observed, so a similar model can be proposed for the zeolitic systems. Furthermore, it is obvious that application of these catalysts strongly depends on the composition of the gas that has to be treated.

650

ACKNOWLEDGEMENT These investigations have been supported by the European Union under contract no. JOU2CT92-0229 and J. R.-M. by a Human Capital & Mobilty grant.

REFERENCES P.L. Crutzen, J. Geophys. Res. 76 (1971) 7311. [1] G.G.d. Soete, Rev. Inst. Franc. Petr. 48 (1993) 413. [2] F. Kapteijn, J. Rodriguez-Mirasol and J.A. Moulijn, Appl. Catal. B: Env. accepted (1996) [3] Y. Li and J.N. Armor, Appl. Catal. B: Env. 1 (1992) L21. [4] J. Valyon, W.S. Millman and W.K. Hall, Catal. Lett. 24 (1994) 215. [5] V.I. Sobolev, G.I. Panov, A.S. Kharitonov, V.N. Romannikov, A.M. Volodin and K.G. Ione, [6] [7] [8] [9] [10]

[11] [12]

[13] [14]

[15] [16] [17]

[18] [19] [20] [21] [22] [23]

[24] [25]

[26] [271 [28]

[29] [30] [311 [32] [33] [341 [35] [36]

J. Catal. 139 (1993) 435. J. Leglise, J.O. Petunchi and W.K. Hall, J. Catal. 86 (1984) 392. C.M. Fu, V.N. Korchak and W.K. Hall, J. Catal. 68 (1981) 166. Y.-F.Chang, J.G.McCarty, E.D.Wachsman and V.L.Wong, Appl. Catal. B: Env. 4 (1994) 283. G.I. Panov, V.I. Sobolev and A.S. Kharitonov, J. Mol. Catal. 61 (1990) 85. Y.-F. Chang, J.G. McCarty and Y.L. Zhang, Catal. Lett. 34 (1995) 163. R.L. Garten, W.N. Delgass and M. Boudart, J. Catal. 18 (1970) 90. R.A. Dalla Betta, R.L. Garten and M. Boudart, J. Catal. 41 (1976) 40. A. Cimino, Chimica e Industria, 56 (1974) 27. F.S. Stone, J. Solid State Chem. 12 (1975) 271. J.O. Petunchi and W.K. Hall, J. Catal. 78 (1982) 327. G.D. Lei, B.J. Adeiman, J. Sarkany and W.M.H. Sachtler, Appl. Catal. B: Env. 5 (1995) 245. G. Spoto, A. Zecchina, S. Bordiga, G. Ricchardi and G. Martra, Appl. Catal. B: Env. 3 (1994) 151. M.A. Kohler, N.W. Cant, M.S. Wainwright and D.L. Trimm, J. Catal. 117 (1989) 188. E.E. Mir6 and J.O. Petunchi, J. Chem. Soc. Faraday Trans. 88 (1992) 1219. M. Anpo, M. Matsuoka, Y. Shioya and H. Yamashita, J. Phys. Chem. 98 (1994) 5744. M.W. Anderson and L. Kevan, J. Phys. Chem. 91 (1987) 4174. M.A. Makarova, E.A. Paukshtis, J.M. Thomas, C. Williams and K.I. Zamaraev, J. Catal. 149 (1994) 36. Y.-Y. Huang, J. Am. Chem. Soc. 95 (1973) 6636. Z. Chajar, M. Primet, H. Praliaud, M. Chevrier, C. Gauthier and F. Mathis, Appl. Catal. B: En v. 4 (1994) 199. Y. Li, T.L. Slager and J.N. Armor, J. Catal. 150 (1994) 388. K.-I. Segawa, Y. Chen, J.E. Kubsh, W.N. Delgass, J.A. Dumesic and W.K. Hall, J. Catal. 76 (1982) 112. M. Iwamoto, H. Yahiro, N. Mizuno, W.-X. Zhang, Y. Mine, H. Furukawa and S. Kagawa, J. Phys. Chem. 96 (1992) 9360. T.E. Hoost, K.A. Laframboise and K. Otto, Catal. Lett. 33 (1995) 105. P.A. Jacobs and H.K. Beyer, J. Phys. Chem. 83 (1979) 1174. A. Cimino and F. Pepe, J. Catal. 25 (1972) 362. J. Valyon and W.K. Hall, J. Phys. Chem. 97 (1993) 1204. Y. Li and J.N. Armor, Appl. Catal. 76 (1991) L 1. Y. Li, P.J. Battavio and J.N. Armor, J. Catal. 142 (1993) 561. J.N. Armor and T.S. Farris, Appl. Catal. B: Env. 4 (1994) L11. V. Indovina, D. Cordischi, M. Occhiuzzi and A. Arieti, J. Chem. Soc. Faraday Trans. 1, 75 (1979) 2177.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

651

O n t h e role o f free r a d i c a l s NO2 and 02 in t h e s e l e c t i v e c a t a l y t i c r e d u c t i o n (SCR) of NOx w i t h CH4 o v e r C o Z S M - 5 and H Z S M - 5 z e o l i t e s Dmitri B. Lukyanov ~ Julie L. d'ltri b, Gustave Sill c and W. Keith Hall c" "Department of Chemistry, UMIST, P.O. Box 88, Manchester M60 1QD, UK bChemical Engineering Department, and =Chemistry Department, University of Pittsburgh, Pittsburgh, PA 15260 USA The reactions of CH4 over CoZSM-5 and HZSM-5 zeolites with the mixtures of NO + O2 and NO2 + 02 and with three oxidizing components separately were studied. Differential reaction rates were determined. Comparison of the "light-off" temperatures as well as the activation energies of these reactions led to the conclusion that the SCR of NO. into N 2 and, consequently, CH 4 oxidation into COx are initiated by the reaction of CH 4 with NO 2. At low temperatures (300-400~ 02 does not compete with NOx for CH,, and its role is limited to the oxidation of NO into NO2. However, at higher temperatures a strong competition between NOx and 02 for CH4 results in a decrease in the selectivity of the SCR process. It is shown that this competition is stronger with CoZSM-5 catalyst than with HZSM-5, and this explains the higher selectivity observed with the latter catalyst. Based on these observations the formation of the CH3o free radical is postulated and possible pathways of the SCR of NOx into N 2 are discussed. 1. INTRODUCTION

Recently, CoZSM-5 and HZSM-5 zeolites were reported [1-4] to be effective catalysts for the SCR of NO with methane in the presence of oxygen. It was shown that 02 greatly enhances the NO conversion into N 2, and an important role of NO2 in the initiation of the reactions has been delineated [3-9]. In the present paper we wish to shift the focus from reduction of NO into N 2 to the competitive oxidation of the hydrocarbons [4] by NO. vis-a-vis 02. Evidence favoring NO2 as the dominant reagent has been presented [6-10], but the reaction of CH4 with 02 alone has not been considered. Hence it was of interest to better define the role of NO2 and 02 in the SCR of NO. with CH4. Consequently, the relative effectiveness of the several oxidizing agents has been determined for the reactions of CH 4 over CoZSM-5 and HZSM-5 catalysts. The same reactions were studied over the NaZSM-5 zeolite and in the homogeneous gas phase [10]. " To whom all correspondence should be addressed.

652

2. EXPERIMENTAL The reactions were carried out in the steady state flow mode as described previously [ 11 ]. Differential kinetics were determined from plots of conversion vs. W/F. Three catalysts CoZSM-5, HZSM-5 and NaZSM-5 (Si/AI = 11) were studied in this work. The catalyst preparation and the standard pretreatment used prior to reaction have been described previously [11]. It involved dehydration in flowing dried 02 as the temperature was raised slowly to 500~ The feed comprised CH 4 (0.28%), NO (0.21%) or NO2 (0.21%), and/or 02 (2.6%) in He. The flow rate was 75 ml/min and the gas hour space velocity (GHSV) was varied between 4 , 5 0 0 and 2 5 0 , 0 0 0 h 1 by changing the weight of catalyst samples.

3. RESULTS AND DISCUSSION 3.1. "Light-off" temperatures of CH 4 combustion in the presence of different oxidizing compounds Studies of CH4 reactions with NOx in the presence or absence of 02 have shown that the "light-off" temperature of methane combustion coincides in every case with the temperature at which N2 formation is initiated [10]. At a GHSV of 22,500 h 1 with CoZSM-5 catalyst this temperature was about 300~ (see Fig. 1) regardless of the nitrogen oxide used. Moreover, in the absence of 02 the "light-off" occurred with NO2 at the same temperature while with NO 450~ was required. In the absence of NO,,, oxidation of CH4 by 02 was observed at about 400~ (Fig. 1B). Taken together with the fact that over Co-containing zeolites [12] NO can be oxidized by 02 into NO2 at temperatures as low as 200~ our data suggest that the SCR of NO into N2 and, consequently, CH4 oxidation into CO2 are initiated by NO2. During NO oxidation to NO2, the catalyst is definitely involved [ 10]. Hence, O atoms held by the catalyst may interact with CH4 molecules, but as shown in Fig. 1 B, not as efficiently as when NO2 is present. Hence, it may be concluded that it is NO2 that initiates the SCR reaction. The same conclusion may be drawn on the basis of our data for the SCR reaction over HZSM-5 catalyst, although in this case the "light-off" temperature is higher. At a GHSV of 22,500 h ~ it is about 350~ for the reactions of methane with the mixtures of NO~ + 02 and with NO2 alone. With NO or 02 alone CH4 oxidation is observed only at temperatures higher than 550~ Evidently Brensted acidity is not essential.

3.2. Coupling of Combustion and SCR From the results discussed above as well as from the literature data [5-10,12-14] it follows that an important role of 02 in the SCR process is to convert NO into NO2. The latter then initiates methane oxidation into COx and is itself reduced into NO and N2. Both NOx and 02 may participate in CH4 oxidation (Fig. 1 B) and the ratio between the rates of these competitive oxidation reactions will be critical for the selectivity of the SCR process. Hence, the absolute rates of CH4 oxidation by 02 were compared with those occurring in the SCR process. The rates of these reactions were determined under different reaction conditions (using the

653

100 v

o~ 80 z o =: 6 0 m.. X

~100

A

0 CH4*NO*O 2

9CH4*NO CH4*N02*0 2 A CH4*N0 2

0

~

O z 40-

c

6O

0

40

0

O

:; 2 0 -

~

C:

C 0

0

o

0 260

860 460 Temperature (OG)

550

0

2O 0 2gO

460 T e m p e r a t u r e (OG) ,950

660

Figure 1. CH4 reactions with different oxidizing compounds over CoZSM-5 catalyst; conversion of NOx into N 2 (A) and of CH 4 into CO 2 (B) as a function of temperature. Catalyst weight was 100 mg, feed contained 0.28% CH4, 0.21% NO or NO2 (when used), and 2.6% 02 (when used) in He at a flow rate of 75 ml/min (GHSV = 22,500 h'l).

E..2o, ko.,,mo, I 4

~

Ea -14.5 kcal/mol

I

2

2. z

=.. o -

_j

E -14.2 kcal/mol 8

...I

0

1.2

-2[.

1.4

~.e

1000 IT (K -1)

.1

E=-29 kcallmol 1 .$

1.6 IO00/T (K -11

B

1.7

Figure 2. Arrhenius plots of differential rates of NO reduction (e) and of CH4 oxidation (o) during the SCR reaction, and for CH4 oxidation by 02 alone (A) over CoZSM-5 (A) and HZSM-5 (B) catalysts. Feed contained 0.28% CH4, 0.21% NO (when used) and 2.6% 02 in He.

654

linear portions of conversion vs. W/F plots) and are compared in Fig. 2. Several points n o w become clear. First, CH4 oxidation into COx and NO reduction into N 2 proceed with the same activation energy and, below 500~ at about the same rates (Fig. 2) indicating that these two reactions are coupled and have the same rate limiting step. Second, the Arrhenius dependence for NO reduction into N 2 is valid up to 500~ i.e., in the temperature range where thermodynamics favors NO2 over NO [6,13]. At higher temperatures the equilibrium constant becomes < 1 causing the observed deviation from linearity.Third, the activation energy of the SCR reaction is higher with CoZSM-5 catalysts than with HZSM-5, and reaction of CH4 with 02 has much the highest activation energy. Fig. 1B shows that at all temperatures the rate of CH4 oxidation by 02 alone is lower than the rate of CH4 oxidation during the SCR reaction, e.g., at 400~ with CoZSM-5 catalyst the difference between these rates is about 10 times. With increasing temperature this difference diminishes due to the different activation energies of these reactions (Fig. 2). At high temperatures these rates become comparable (in considering Figs. 1B and 2 recall that the rate of CH 4 oxidation during the SCR process includes a contribution from the rate of CH4 oxidation by 02 alone). These data suggest that below 5000C 02 does not compete effectively with NO,, for CH4, but that at high temperatures such a competition must exist. The data of Table I support this view. At 400~ an increase in 02 concentration results in an increase in conversions of both NO into N 2 and CH 4 into CO2. At the same time, variation of O2 concentration by a factor of 13 has practically no effect on the Table 1 Effecl; of 02 r

on the $CR reaction* over CoZSM-5""

At 400~ A B C

Conv. of NO into N 2 (%) Conv. of CH4into CO= (%) NO convCrl;ed into N~(mol) CH 4 converted into C02 (mol)

0.4

Concentration of 02 (%) 1.0 2.6 5.2

13.8 6.4 1.62

17.9 8.6 1.56

20.4 9.7 1.58

20.7 10.3 1.51

21.5 12.5 1.29

25.8 19.2 1.01

24.1 28.8 0.63

21.9 44.2 0.37

At 550 ~ D Cony. of NO into N 2 (%) E Cony. of CH4 into CO2 (%) F N O converted inl;o N~(mol) CH. converted into CD2 (mol)

*The reactant gas contained 0.28% CH4 (9.375 pmol/min), 0.21% NO (7.03 pmol/min) and x% 02 in He at a total flow rate of 75 ml/min. * *Steps A, B and C correspond to 400~ and GHSV of 45,000 hl; steps D, E and F correspond to 550~ and GHSV of 250,000 h~.

655

selectivity of the SCR process (defined as a ratio of the number of NO molecules converted into N2 to the number of CH4 molecules converted into CO2). This result can be easily Understood in terms of the higher rates of NO2 formation at the lower temperature (and, consequently, in the higher rates of the SCR process) at higher 02 concentrations. A different picture is observed at 550~ In this case an increase in 02 concentration results mainly in an increase in CH4 combustion (due to the higher activation energy for this reaction) and, consequently, in a decrease in the selectivity of the SCR process. Fig. 3 shows the effect of temperature on the selectivity of the SCR of NO with methane in the presence of oxygen. The data show that the highest selectivity of the SCR process is observed at low temperatures and low conversions of the reactants (NO and CH4) and does not differ greatly for CoZSM-5 and HZSM-5 catalysts. This fact can be easily understood since under these conditions CH4 should be oxidized, according to our data, solely by NOx. An increase in temperature or a decrease in NOx concentration should be followed by enhancement of the competition between 02 and NO2 for methane, and, consequently, by a decrease of the SCR selectivity as observed. Moreover, it follows from Fig. 2 that at any given temperature the difference between the rates of CH4 oxidation by the mixture of NOx + 02 and by 02 alone is larger for HZSM-5 catalyst than for CoZSM-5 . Thus, the competition between NOx and 02 for CH4 is stronger with CoZSM-5 than with HZSM-5. This explains the fact that the SCR selectivity (Fig. 3) is higher with HZSM-5 than with CoZSM-5.

r

8

"~" 8 [ 9400~ J O 450~

9400~ O 450~

.=..

E --~ 6

L/

O

E

o z

"

65ooc

4

q..

0

O

:; 2 r O

o

o

0

2

4

6

8

10

C o n y . of CH 4 ( ~ m o l l m l n )

g'o

0

2

4

6

8

10

C o n y . of CH 4 (~mollmln)

Figure 3. Effect of temperature on the selectivity of the SCR reaction over CoZSM-5 (A) and HZSM-5 (B) catalysts. Feed contained 0 . 2 8 % CH4, 0 . 2 1 % NO and 2.6% 02 in He at a flow rate of 75 ml/min ( flow rates of CH4 and NO were 9.375 and 7.03 pmol/min).

656 100

100

80

8O

v r

O z

6O

O 60 z

O

40

o 40

O

20

o 2O

q=-

o

O

0 260

360

0

450

660

Temperature (oc)

660

250

380

460

650

660

Temperature (~

Figure 4. Conversion of NO2 into different products by reaction w i t h CH4 over C o Z S M - 5 (A) and HZSM-5 (B) at various temperatures. Feed contained 0 . 2 8 % CH 4 and 0 . 2 1 % NO2 in He at a f l o w rate of 75 ml/min. GHSV was 2 2 , 5 0 0 h "1.

6p

4

A

ol

,

9NO 2 into N2 O NO 2 into NO A CH4 into COx

B 4

A

~ 3

_.52

kcal/mol

=,21 -

'

0

1.3

i

1.4

1.5

1.6

IO001T (K "1)

1.7

1.8

0

,

1.3

1.4

EI I -13.6 kcal/mol i

1.6

I

1.6

I

1.7

1.8

I O 0 0 / T (K "1)

Figure 5. Arrhenius plots of differential rates of NO2 reduction into NO and N 2, and for CH4 oxidation into COx over CoZSM-5 (A) and HZSM-5 (B). Feed contained 0 . 2 8 % CH4 and 0 . 2 1 % NO2 in He.

657

3.3. Reaction of CH4 with NOz

This reaction was studied in detail and some typical results are shown in Fig. 4. Previously, it was demonstrated [10] that in the absence of CH4, NO2 achieves equilibrium with NO + 89 02 over CoZSM-5 and HZSM-5 catalysts at temperatures higher than 350~ As shown in Fig. 4, in the presence of CH4 t w o additional reactions occur: (a) reduction into NO, and (b) reduction into N 2. The oxygen released during these two reactions is used to oxidize CH4 into CO2over CoZSM-5 catalyst or into CO2 + CO over HZSM-5. In the presence of CH4the conversion of NO 2 into NO + 89 decreases significantly and at 400~ and a GHSV of 22,500 h1 this decrease is about 9 times with CoZSM-5 (Fig. 4A) and 3 times with HZSM-5 (Fig. 4B). Invariably more of the NO 2 is reduced into NO than into N 2. At low temperatures (300-400~ the rate of this conversion, under differential conditions, was about 2 times higher than the rate of N2 formation. With increasing temperature the concentration of 02 in the tail gas decreases (probably, due to the reaction with CH4) and at temperatures higher than 500~ it is not observed in the reaction products. Simultaneously, the total conversion of NO 2 (into NO and N 2) reaches 100% at about 500~ Temperatures above 500~ resulted in a decrease in NO2 conversion into NO and in a corresponding increase in the conversion into N 2. Over HZSM-5 (Fig. 4B) this was not observed. These data clearly show that NO2 reduction with CH4 into NO and into N2 proceeds over CoZSM-5 and HZSM-5 catalysts via processes that have at least one common reaction step, viz., the activation of CH4. It is of interest to compare the activation energy data for the reactions of NO2 (Fig. 5) with those for the SCR reaction (Fig. 2). Within experimental error, they are the same; 22 kcal/mol and 21 kcal/mol, respectively for CoZSM-5 and 14 kcal/mol for HZSM-5 processes. The conclusion that combustion of CH4 and reduction of NOx processes are coupled and have the same rate limiting step is strongly supported by these facts. We therefore conclude that the SCR of NO with CH 4 in the presence of 02 is initiated by the reaction of NO2 with CH 4. Additional information on the reaction system (NO2 + CH,) follows from Fig. 4A. These data were obtained over CoZSM-5 catalyst at a GHSV of 22,500 h 1 . At this space velocity and at temperatures higher than 450~ the direct reaction between NO and CH 4 occurs resulting in N 2formation. Obviously, this reaction must also occur in the system of NO2 + CH4 and explains a decrease in NO2 conversion into NO and a corresponding increase in conversion into N 2 at temperatures higher than 550~ However, in the SCR of NOx in the presence of 02 the reaction between NO and CH, does not play an appreciable role, since it proceeds only at high temperatures and with a rate much lower than the rate of CH4 oxidation by 02. Interestingly, Shelef et al. [7] fed approximately equilibrium mixtures of NO and NO2 with and without 02 over their catalyst (CuZSM-5) together with C3H8 and found most of the NO 2 fed had been converted into NO or N2; the system moved away from equilibrium by the selective reaction of NO2 with hydrocarbon faster than it could be produced by NO reacting with 02.

658

3.4. Pathways of the SCR of NO with CH4 in the presence of O z We have already established that oxidation of CH 4 and consequently the catalytic reduction of NOx are initiated by interaction of the former with NO2. At a GHSV of 2 2 , 5 0 0 h 1 "light-off" occurs at 300 ~ and 350~ with CoZSM-5 or HZSM-5 catalysts, respectively. In the empty tube the "light-off" of CH4 combustion with NO 2 occurs about 450~ [10]. Under these conditions the homogeneous oxidation of CH4 and the conversion of NO2 into NO (formation of N 2 was not observed) increase sharply with temperature and at 600~ reaches 100%. With NaZSM-5 catalysts this picture does not change much, the sole difference from the homogeneous reaction being the difference in the products of CH4 oxidation. In the presence of the catalyst, CH 4 is oxidized into CO and CO2 (CO/CO2 ratio is about 2) and the carbon balance closed at all temperatures. In the empty tube the carbon balance did not close at 500 ~ and 550~ possibly due to the formation of formaldehyde [ 10]. Thus, we suggest that over NaZSM-5 catalyst a rapid oxidation of intermediate oxidation products occurs. Yokoyama and Misono [8] reported that N 2 was formed in the reaction of NO2 with C3He over NaZSM-5. Many years ago Wojciechowski and Laidler [15] studied the homogeneous decomposition of CH4 and C2He in the presence of NO and concluded that the stable free radical NO could abstract an H atom from a paraffin molecule. Since NO2 is also a stable free radical and a much stronger oxidizing agent than NO, the reaction CH4 + NO2 --" C H 3 9+ "OH + NO (or HONO)

(1)

seems very likely. This finds support in the work of Cant and co-workers [16,17] who found a first order isotope effect on the rate of N 2 formation when CD4 was substituted for CH4. Coupling of CH3 9radicals would not be expected because of much higher concentration of NO,, playing the role of a radical trap, and was not observed. The formation of CH3NO2 (and/or CH3NO) seems much more likely. Yokoyama and Misono [8] reported that the former, when introduced as a reactant, reacts with NO2 forming CO2 and H20. Similarly, formaldehyde may be produced by (2),

CH3 9 + ONO --- CH20 + HNO

and may react further with NO2 producing CO that, in turn, may be oxidized to CO2. i.e. CH20

+

NO 2 ~ CO +

CO + NO2 ~

CO2 + NO

H20 +

NO

(3), (4).

Equations (2) - (4), considered together with reactions like eOH + HNO ~ H20 + NO and CH3 9+ HNO -~ CH4 + NO, explain the homogeneous reduction of NO2 into NO and the corresponding oxidation of CH4 into CH20, H20, CO and CO2. The same

659

reactions may occur over NaZSM-5, but in this case, the rates of the reactions (3) and (4) are much higher than in the empty reactor. Moreover, it seems likely that the reactions (2)-(4) occur also over CoZSM-5 and HZSM-5 catalysts forming the pathway for NO2 reduction into NO. In contrast to NaZSM-5 zeolite, introduction of CoZSM-5 or HZSM-5 zeolite in the reaction system shifts the "light-off" temperature and modifies the chemistry; n o w not only NO but N 2 is formed. Hence, some intermediate species required for N 2 formation must be stabilized on the catalyst surface. The "light-off"temperature shifts observed with CoZSM-5 and HZSM-5 catalysts may result from the enhanced redox capacity provided by these catalysts or from the NOJNO equilibrium achieved more readily than with NaZSM-5. Moreover, equilibrium is approached at a somewhat lower temperature over CoZSM-5 than HZSM-5, and much lower than with the empty reactor (see Fig. 1 of Ref. lO).The decomposition reaction of NO2 into NO + 89 occurs readily on these catalysts and the "light-off" temperature of both combustion and SCR is lower in comparison with that of the homogeneous reaction. Summarizing, our findings permit us to speculate on possible mechanisms occurring in the SCR reaction. Over these catalysts NO 2 is reduced into NO and N 2. These t w o reactions are coupled and have the same rate limiting step, i.e., activation of CH4 by NO2. The results suggest that this reaction results in formation of the CH3e radical and this idea finds support in the isotope effects for SCR [17] vis-a-vis CH4 coupling [ 16] and in the fact that CH4 coupling catalysts also catalyze SCR[18,19]. Finally, recently reported [20] calculations on CH4 activation with an O atom adsorbed on a surface demonstrate that the abstraction of an H atom from a CH4 molecule with formation of CH3o free radical should be energetically preferable compared with the other possible mechanisms of CH 4 activation. Thus, the chemistry of Eq. (1) seems plausible. It is not unreasonable to think that free radical chemistry can occur within the zeolite pore system, just as it can in solution. However, the lack of molecular sieving effects with larger hydrocarbons [4] leads us to believe that part of the reaction occurs homogeneously. The NO formed as product may desorb from the catalyst, react with 02 to produce NO2, or with the CH3 9radical to produce CH3NO molecule. The adsorbed CH3NO molecule may be transformed into HCN via dehydration, i.e., Z-CH3NO---Z-HCN + H20

(5).

The formation of HCN and adsorbed CN species during the SCR process over CeZSM-5 and CuZSM-5 zeolites have been reported previously by several research groups [8,21,22]. HCN may undergo various transformations in the presence of NO2 and 02 [23,24]. Some of these produce eCN and eNCO radicals. The latter radical, as has been shown by Cooper et al. [25] reacts with NO producing the ON-NCO molecule which, in turn, decomposes thermally into N2 and CO2. The transient experiments performed by Yogo and Kikuchi [9] over GaZSM-5 and InZSM-5 together with the results of Hayes et al. [22] suggest that eCN may play a role in forming the N-N bond.

660 Another possible path for the formation of eNCO radical may include formation of CH3NO2 which may undergo transformation on the active sites of the catalyst into H20 and HC(N)O species. The latter may be converted, according to the literature data [24], into oNCO radical. In conclusion, although at present time we cannot unambiguously define the reaction steps leading to N2 formation, we have attempted to suggest steps that seem to be reasonable and hope that our proposals will be useful for further investigation of the detailed mechanism of the SCR process. Acknowledgement Thanks are due to the Department of Energy, Division of Basic Energy Sciences, under Grant No. DE-FGO2-95ER14539. REFERENCES o

2. 3. 4. 5. .

7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25.

Y. Li and J.N. Armor, Appl. Catal. B, 1 (1992) L31. Y. Li and J.N. Armor, Appl. Catal. B, 2 (1993) 239. K. Yogo, M. Umeno, H. Watanabe and E. Kikuchi, Catal. Lett., 19 (1993) F. Witzel, G.A. Sill and W.K. Hall, J. Catal., 149 (1994) 229. J.N. Armor and Y. Li, Preprints of Symposium on NOx Reduction. ACS, Division of Petroleum Chemistry, Inc., 39 (1994) 141. J.O. Petunchi and W.K. Hall, Appl. Catal. B, 2 (1993) L17. M. Shelef, C.N. Montreuil and H.W. Jen, Catal. Lett., 26 (1994) 277. C. Yokoyama and M. Misono, J. Catal., 150 (1994) 9. K. Yogo and E. Kikuchi, Stud. Surf. Sci. Catal., 84 (1994) 1547. D.B. Lukyanov, G. Sill, J.L. d'ltri and W.K. Hall, J. Catal., 153 (1995) 265. J.O. Petunchi, G. Sill and W.K. Hall, Appl. Catal. B, 2 (1993) 303. Y. Li, T.L. Slager and J.N. Armor, J. Catal., 150 (1994) 388. K.A. Betkhe, C. Li, M.C. Kung and H.H. Kung, Catal. Lett. 31 (1995) 287. H. Yasuda, T. Miyamoto and M. Misono, in "ACS Symposium Series 587", Chapter 9 (1994). B.W. Wojciechowski and K.J. Laidler, Can. J. Chem., 38 (1960) 1027. N.W. Cant, E.M. Kennedy and P.F. Nelson, J. Phys. Chem., 97 (1993) 1445. A.D. Cowan, R. Dumpelmann and N.W. Cant, J. Catal., 151, (1995) 356. X. Zhang, A.B. Waiters and M.A. Vannice, J. Catal., 146 (1994) 568. X. Zhang, A.B. Waiters and M.A. Vannice, Appl. Catal. B, 4 (1994) 237. M.Yu. Sinev, L.Ya. Margolis and V.N. Korchak, Russian Chem. Rev., 64 (1995) 349 (English translation). F. Radtke, R.A. Koeppel and A. Baiker, Appl. Catal. A, 107 (1994) L125. N.W. Hayes, W. Grunert, G.J .Hutchings, R.W. Joyner and E.S. Shpiro, J. Chem. Soc., Chem. Commun., (1994) 531. J.A. Miller and C.T. Bowman, Prog. Energy Combust. Sci., 15 (1989) 287. M.C. Lin, u He and C.F. Melius, Intern. J. Chem. Kinetics, 24 (1992) 1103. W.F. Cooper, J. Park and J.F. Hershberger, J. Phys. Chem., 97 (1993) 3283.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

661

An infrared study of N O reduction by CH4 o v e r C o - Z S M - 5 A. W. Aylor, L. J. Lobree, J. A. Reimer, and A. T. Bell Center for Advanced Materials, Lawrence Berkeley National Laboratory and Department of Chemical Engineering, University of California, Berkeley, CA 94720, USA An in situ infrared investigation has been conducted of the reduction of NO by CH4 over Co-ZSM-5. In the presence of 02, NO2 is formed via the oxidation of NO. Adsorbed NO2 then reacts with CH4. Nitrile species are observed and found to react very rapidly with NO2, and at a somewhat slower rate with NO and 02. The dynamics of the disappearance of CN species suggests that they are reactive intermediates, and that N2 and CO2 are produced by the reaction of CN species with NO2. While isocyanate species are also observed, these species are associated with Al atoms in the zeolite lattice and do not act as reaction intermediates. A mechanism for NO reduction is proposed that explains why O 2 facilitates the reduction of NO by CH4, and why NO facilitates the oxidation of CH4 by 02. 1. INTRODUCTION There has been considerable interest recently in the use of metal-exchanged zeolites for the selective reduction (SCR) of NO by methane [ 1-20]. Amongst the various catalysts tested, CoZSM-5 and Co-ferrierite have shown particularly high activity. Attempts to explain the mechanism by which CH4 reduces NO over these catalysts have largely been based on studies of reaction kinetics. Several studies have shown that the rate of NO reduction is considerably more rapid when 02 is added to the feed [1, 2, 10, 13, 14, 18], and that under such circumstances NO2 is formed rapidly by the reaction of NO with 02. It has also been found that the reaction of NO2 with CH4 occurs at a comparable rate to that observed for the reaction of NO with CH4 in the presence of 02 [ 13, 14, 18]. Based on this evidence it has been proposed that in the presence of 02 the reduction of NO is initiated by the oxidation of NO to NO2, then subsequent reaction of CH4 with either gas-phase or adsorbed NO2 [11, 13-16, 18]. Isotopic labeling studies have demonstrated that the reduction of NO in the presence of 02 occurs more rapidly with CH4 than with CD4, suggesting that C-H bond scission is very likely the rate-limiting step in SCR [17]. Infrared studies of NO adsorption on Co-ZSM-5 and Coferrierite have identified the presence of both mono- and di-nitrosyls at room temperature [ 15]. Over Co-ferrierite both of these species can be converted to adsorbed NO2 in the presence of oxygen at elevated temperatures. The adsorbed NO2 is thought to react with CH4 thereby producing CH3 radicals and adsorbed HNO2. More recently, studies involving infrared spectrsocopy and mass spectrometry have shown that CH4 reacts with adsorbed NO 2 to form N2 and CO2 [20]. The goal of this work was to determine surface species present under reaction conditions, and to investigate the interactions of adsorbed NO and NO2 with CH4. Infrared spectra were collected under reaction conditions, and in various mixtures of NO, Oz, NO2, and CH4. It was of particular interest to make direct observations of the factors affecting the formation of NO2 and its reaction with CH4, since NO2 has been suggested as an intermediate in the reduction of NO by CH4. A further objective was to elucidate the pathway by which N2 and CO2 are formed.

662 2. EXPER/MENTAL Na-ZSM-5 was obtained from UOP. About 15 g of the zeolite was added to a 3.5 L solution of 0.01 M cobalt acetate. This mixture was stirred at 25 ~ for 21 hr, then at 60 ~ for 33 hr, and finally at 70 ~ for 19 h [1]. The zeolite was faltered, washed, and dried overnight in a vacuum oven at 120 ~ Elemental analysis of the catalyst determined the Si/A1 ratio to be 18.4, and the Co/A1 ratio to be 0.28. For infrared spectroscopy, 20-50 mg of the cobalt-exchanged zeolite was pressed into a self-supporting wafer and placed into an infrared cell similar to that described by Joly et al. [21]. Spectra were recorded on a Digilab FTS-50 Fourier-transform infrared spectrometer at a resolution of 4 cm-l. Typically, 64 or 256 scans were coadded to obtain a good signal-to-noise ratio. A reference spectrum of Co-ZSM-5 in He taken at the same temperature was subtracted from each spectrum. Gases were supplied to the infrared cell from a gas manifold. 4.99% NO in He and 2.14% CH4 in He were obtained from Matheson. Oxygen and helium were obtained on-site. The He, NO, and CH4 cylinders were passed through an oxysorb trap, an ascarite trap, and a molecular sieve trap, in that order, for additional purification. The 02 was passed through an ascarite and a molecular sieve trap. Prior to each experiment the catalyst was: (1) heated at 500 ~ in 02 for about 30 min., (2) heated at 500 ~ in He for about 3 h, and (3) cooled to the desired temperature in helium. The temperature ramp experiments were run at about 0.6 ~ To determine its activity, the catalyst was placed in a quartz microreactor. Reactants were supplied through mass flow controllers and the product composition was determined by mass spectrometry. A typical reaction mixture contained 3,600 ppm NO, 1.06% CH4, and 6.0% 02, with the balance He. A 0.05 g sample of the catalyst was used with a total flow rate of 100 cm3/min, resulting in a GHSV = 60,000 (based on an apparent bulk density of the zeolite of 0.5 g/cm3). The conversion of NO was based on the amount of N 2 formed and the conversion of CH4 was based on the amount of CO2 formed. Carbon and nitrogen mass balances were closed to within 10%. 3. RESULTS Figure 1 shows the effects of temperature on the conversion of NO to N2. In the absence of O2 significant conversion of NO occurs above 500 ~ When 02 is present the light off temperature decreases to 350 ~ Similar light off temperatures are observed for the conversion of CH4 to CO2. The ratio of NO to CH4 consumption increases from 0.7 to 1.04 as the temperature rises from 550~ to 700~ for the reaction of NO with CH4 in the absence of O2. This ratio is significantly lower than 4.0, the ratio for the reaction CH4 + 4 NO -~ 2 N2 + CO2 + H20, suggesting that additional CO2 may be produced by steam reforming of CH4. In the presence of 02 in the feed, the ratio of NO to CH4 consumption decreases from 0.84 to 0.18 as the temperature rises from 400~ to 600~ Similar ratios of NO to CH4 consumption have been reported previously [2, 13], and the decline in this ratio with increasing temperature has been attributed to the production of CO2 as a consequence of CH4 combustion by O2. Figure 2 shows a series of infrared spectra taken during NO desorption from the catalyst in He after it had been exposed to 6,600 ppm of NO for 42 min at room temperature. At 25 ~ sharp bands are observed at 2132, 1941, 1894, and 1815 cm-1, together with shoulders at 1874 and 1799 cm-l, and a doublet centered at about 1400 cm-~. Weak bands are also seen at 1633, 1599 and 1528 cm-~. Purging the sample in He for 4 min at 25 ~ causes the disappearance of the band at 1633 cm-l, and an increase in intensity of the features at 1599 and 1528 cm-l.

663

N O + I:::I't4 + O 2

- - . - - N O + C~-I,, + 02 ---,.-NO + C~FI,,

40

~ 4o

o -~ 30

~9 30

O Z 2o

~ 2o

et

~NO

50

+

o

i

IQ

C

L

200

300

4OO

500

60O

700

200

300

400

500

600

700

Temperature (~

Temperature (~ Figure I a. N O conversion versus temperature. [NO] = 3600 p p m . [CH 4] = 1.06%, [02.] = 6.0%.

Figure I b. CH~ oonvea~ion versus temperature. [NO] = 3600 ppm, [CH,] = 1.06%, [02] = 6.0%.

Total flow rate = 100.0 cm3/min, GHSV = 60,000.

Total flow rate = 100.0 cm3/min, GHSV = 60,000.

4

-

-

-

i

-

-

.

i

.

-

-

v

-

.

-

i

-

-

-

3.5

3.5

3

3

2.5

2.

0.5 !~5o'c~ ~400~ 0

l

2350

.

.

.

-2

O.5

. .

.

2150

~'~2

.

4.00~ .

.

.

.

1950

.

.

.

.

.

1750

.

.

' . . . i

1550

1350

wavenumber (cm-!) Figure 2. 6600 ppm NO was preadsorbed for 40 min. at room t e m p e r a r ~ and desorbed in He, during temperatu~ ramp.

0 2350

_ 2150

1950

__ 1750

1550

1350

wavenumber (cm-1) Figure 3. 6400 ppm NO was pxeadsorbed for 40 min at room t e m p e ~ u r e , and desorbed in 2.14% CH,/He during temperature ramp.

664 Based on previous studies [15, 22-25], the band at 1941 cm-1 is assigned to Co2*(NO), and the pair of bands at 1894 and 1815 cm-~, to Co2+(NO)2. The shoulders at 1874 and 1799 cm-~ may be due to a second dinitrosyl species. While little is known about the location and coordination of the Co2§ in ZSM-5, it is likely that cobalt ions are associated with both [Si-OA1]- and [A1-O-Si-O-A1]2- structures in the zeolite. In the former case, the cobalt cations are assumed to be present as Co2+(OH-) cations and in the latter case as Co2§ cations. The presence of cobalt cations in different environments could account for the appearance of two sets of dinitrosyl bands. The band at 2132 cm-1 is present not only on Co-ZSM-5 but also on H-ZSM-5 and Na-ZSM-5, and has been observed by several authors on Cu-ZSM-5 [26-28]. A careful investigation of this feature suggests that it is attributable to NO2 ~ associated with both Bronsted acid and M§ cations [29]. The doublet at 1400 cm-~ is identical in position and appearance to that observed in a sample of Na-Y doped with NaNO3 [30]. We therefore assign this peak to a NO3- species affiliated with the residual sodium in the catalyst. The position of the band at 1528 cm-l is very similar to that for nitrito species in Co-A and Co-Y [22, 23] and is, therefore, assigned to Co-ONO. The features at 1599, and 1574 cm-I are best assigned to Co-O2NO [30]. The band at 1633 cm-1 is similar to that observed on H-, Na-, and Cu-ZSM-5. We believe that this feature is best assigned to nitrito (NO2) or nitrate (NO3-) species. Temperatures in excess of 150 ~ are required to desorb NO from Co-ZSM-5 (see Fig. 2). The ratio of the intensity of either of the dinitrosyl peaks to the mononitrosyl peak decreases with increasing temperature suggesting that NO in dinitrosyls is less strongly bound than that in mononitrosyls. The NO2 a§ band decreases in intensity slowly and disappears completely at temperatures above 100 ~ NaNO3 is not very stable, and decomposes completely by the time the temperature reaches 100 ~ Switching from NO to He at room temperature leads to a sudden increase in the intensity of the bands associated with NO2 and NO3. These features increase in intensity up to 150 ~ whereafter they decrease, disappearing completely above 350 ~ Analysis of the effluent gas during TPD reveals that NO desorbs from Co-ZSM-5 without decomposition. A series of spectra taken during the TPD of NO into a stream containing 2.14% CH4 in He are shown in Figure 3. At temperatures up to 300 ~ these spectra are identical to those presented in Figure 2. However, at 300 ~ the nitrito peak at 1515 cm-I [shifted from its position at 1528 cm-1 at room temperature] is notably absent, presumably due to the reaction of Co-ONO with CH4. Spectra obtained during TPR of NO are shown in Figure 4. Nitrosyls are the principal species observed as the temperature is elevated, and in the presence of gas phase NO there is no adsorbed NO2 at temperatures above 150 ~ The absence of adsorbed NO2 is very likely due to its displacement by NO, since it was found that NO readily displaces adsorbed NO2 at 300 ~ During exposure of the catalyst to NO at 300 ~ the intensity of the band at 1528 cm-1 goes through a maximum with time. This suggests that a small amount of NO2 is probably being formed throughout the temperature ramp and is simultaneously being displaced by NO. As the temperature increases the ratio of mononitrosyl to dinitrosyl species increases. Figure 5 shows a series of infrared spectra taken during the TPR of NO with CH4. At temperatures less than 350 ~ the spectra in Figure 5 are virtually identical to those for NO TPR seen in Figure 4. Above 350 ~ the nitrosyl bands are more intense in the presence of CH4. A new peak appears at 2270 cm-I when the temperature is raised to 400 ~ and above, and another one appears at 2173 cm-~ at 450 ~ Neither of these bands was observed when the reaction mixture was passed over Na-ZSM-5. When 15NO was substituted for 14NO, the two bands appeared at 2256 cm-1 and 2144 cm-l, and when 13CH4 was substituted for 12CH4, the bands shifted to 2237 cm-1 and 2132 cm-l. Based on previous studies [32-34], the band at 2270 cm-1 is best attributed to NCO species adsorbed at AI3§ sites. The observed shift in the

665

3.5

3.5

3

1

2.5

2 25oc

1.5

.

lsooc

___...-~/~ ~ _

1 25o*c

0.5

0.5 ~

0

2350

2150

1950

1750

1550

1350

.

,

2350

,

.

,

I

_

,

2150

,

.

i

.

1950

.

1750

1550

1350

wavenumber (cm" 1)

wavenumber (cm-t)

Figure 5. 5100 ppm NO + 9800 ppm CH 4 was passed over the catalyst for 20 min at room temperature, before beginmng temperatur~ ramp.

Figure 4. 6000 ppm NO was pxeadsorbed for 20 min at room temperature, and maintained over the catalyst during Icmperatmr ramp.

3.5 - , - , . , . , . , . , . , - , - , - , - , . , . , . , . , . , - , . , .

3

1.6

2.5

!so~ 8

1.2

2 ~5

200* 250~

1.5

'400~

0.5

0.8

0.4

. l . l . l . l

. l l l . l l i . l . l l t , l . l . i . l . t . i . l .

2300 2200 2100 2000 1900 1800 1700 1600 1500 1400

wavenumber (cm" l)

Figure 6.

3800 ppm NO + 5.4% 02 was passed over the c.atalyst for 20 min at room temperature, before beginmng temperature ramp.

2300 2200 211111 211131)1900 1800 1711(I 1600 151113141713 wavenumber (cm-t) Figure 7. 3600 ppm NO + 1.1% CH,t + 5.6% 0 2 was passed over the catalyst for 20 min at room temperature, before beginning temperature ramp.

666 frequency of the of this feature upon isotopic substitution is consistent with the assignment to the band to NCO species. The band at 2173 cm-1 is more difficult to assign. Based on its position, this feature might be attributed to NCO species adsorbed on Co2+ [31]. This would also be consistent with the observation of a band at 2185 cm-I for NCO species adsorbed on Cu2+ cations in Cu-ZSM-5 [34]. However, the observed frequency shifts upon isotopic substitution are inconsistent with this assignment, and are in much better agreement with the assignment of the 2173 cm-1 band to CN species on Co2§ While the position of this band is about 30 cm-1 lower than that reported for Co complexes containing CN ligands, the presence of electronegative ligands such as NO and NO2 within the coordination sphere of cyano complexes is known to shift the frequency of the C-N vibration upscale by as much as 30 cm-l [31]. Since the Co2§ cations in Co-ZSM-5 are assumed to occur as Co2+(OH-), the presence of OH- may be responsible for the upscale shift in the vibrational frequency of the adsorbed CN species. A series of spectra taken during TPR of a mixture of NO and 02 are presented in Figure 6. Bands are observed for both mon- and dinitrosyls, together with bands characteristic of NO2 and NO3- species. As the temperature rises, the ratio of nitrosyl to NO2/NO3 bands increases, consistent with what is expected on the basis of equilibrium considerations for the reaction NO + 1/2 02 = NO2 [35]. Figure 7 shows spectra recorded during a TPR experiment in which a mixture of NO, 02, and CH4 are passed over the catalyst. At room temperature several new bands are present. These are located at 2189, 1878, and 1747 cm-~. The peak at 2189 cm-~ is most likely due to NO2r~ [36, 37], since this band is observed upon adsorption of NO2 at room temperature (see Figure 8). The band at 1747 cm-I is assigned to N204 [38], and the feature at 1878 cm-I is probably due to N203 [30, 39]. Elevating the temperature removes all three of these bands. The NO2/NO3 bands are quite intense at room temperature relative to the mono- and dinitrosyl nitrosyl bands. As the temperature rises, the ratio of nitrosyl to NO2/NO3 band intensities increases in a manner similar to that seen in Figure 6. Above 350 ~ the intensities of the NO2 and NO3 bands are smaller than those observed in the absence of CH4, a pattern identical to that already noted in the comparison of Figures 2 and 3. When the temperature is raised to 450 ~ the only features remaining are weak bands located at 2264, 1934, and 1635 cm-1. The first two bands are attributed to A13+-NCO and Co2§ respectively, and the third is due to adsorbed H20. Since the formation of NO2 can occur homogeneously, it was of interest to establish whether adsorbed NO could be oxidized. NO was adsorbed at 225 ~ after which the infrared cell was purged with He and subsequently a stream of 10.1% 02 in He was allowed to flow over the catalyst. Prior to the introduction of the O2-containing stream, the only features evident were those for mono- and dinitrosyls. In the presence of 02 at 225 ~ the intensities of the bands for both mono- and dinitrosyl species attenuated and new features appeared at 1628 and 1518 cm-1, corresponding to nitrate and nitrito species, respectively. A similar experiment carried out in the absence of 02, showed only a small decrease in the intensity of the nitrosyl bands due to NO desorption and the absence of bands for nitrate and nitrito species during a 30 min purge in He at 225 ~ To investigate the behavior of adsorbed NO2 further, TPR experiments with both NO2 and an NO2/CH4 mixture were performed. These results are shown in Figs. 8 and 9. Figure 8 shows that upon exposure of the catalyst to NO2, bands appear at 2189, 2133, 1747, 1672, and 1528 cm-l. Of these features, only the band located at 1672 cm-I represents a feature not seen previously in the spectra of adsorbed NO. While this band can be attributed to some form of adsorbed NO2, the exact structure is not known. Upon raising the temperature to 50 ~ all of the bands below 1700 cm-1 disappear, the band at 1528 cm-l decreases in intensity, and the

667

4.5

.

, . , . , . , ' . , ,

d,.,,

4.5

~ . , . , . , - , - , - , . , . , - , - , -

3.5

3.5

3

3

~ 2.5

~

-

,-

,

-',

-

,

-

,

-,

-

,'--v':,

-

,-',~"-

,'-

,

-

,-

,

-

,

-

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-

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250~

f ,,x __

41~C 0.5

0.5

_,.,.t.t.t.t. ). J . t . l . , . l . l . , . t.t .1., . 2300 2200 2100 2000 1900 1800 1700 1600 1500 1400

2300 2200 2100 2000 1900 1800 1700 1600 1500 1400

wavenumlxa" (cm-t)

wavenumber (cm" t)

Figure 8. 5100 ppm NO2 was passed over the catalyst for 20 rain at room tempetatmc, before bcginmng t c m l x a ~ m ramp.

Figure 9. 4900 ppm NO2 + 1.1% CH 4 was passed over the catalyst for 20 rain at room tcmpcmun=, before bcginning temperature ramp.

I

Co ~

+ NO

~

Co 2+ (NO)

2

Co 2+('NO) + NO

~

Co2§

3

Co2+(NO) 2 + 0 2

~

C~

4

Co2+(NO2) + CH4

~

C~

5

Co2+(CH3NO) + 02 (OH-)

6

Co2+(CH2 NO)

7

Co2+(CN) + NO 2

~

Co 2+ + N 2 + CO 2

$

Co2*(CN) +

NO

~

Co 2. + N 2 + CO

9

Co2§

02

~

C~

+

2 )

+ NO2

+ OH.

Co2+(CH2NO) + HO2- (H20) Co2+(CN) + H 2 0

+ NO + CO

Figure 10. Proposed reaction mechanism for NO reduction by CH 4 in the presence of O 2 .

668 band at 1672 cm-I increases in intensity. A new feature also appears at 1633 cm-1. At temperatures above 50 ~ the band at 1672 cm-1 disappears, the band at 1633 cm-1 decreases in intensity, and the band at 1528 cm-1 first increases and then decreases in intensity. Comparison of Figures 8 and 9 shows that the spectra in Figure 9, taken in the presence of CH4, are identical to those shown in Figure 8 below 400 ~ To determine whether the isocyanate and/or nitrile species are reaction intermediates, experiments were performed to compare the reactivity of these species in He, NO, 02, and NO2. Isocyanate and nitrile species were first produced at 450 ~ by reaction of NO with CH4. At t = 0 the reactant flow was replaced by either He or a mixture containing 1.0% NO, 0 2, o r NO2 in He. Upon cessation of the flow of NO and CH4, there was an immediate rise in the intensity of the NCO band (2270 cm-1) and a decrease in the intensity of the CN band (2173 cm-1) caused by partial oxidation of CN to NCO species by residual NO or NO2 in the reactor. The intensity of the CN band then decreased with an apparent first-order rate coefficient of 2.1x10-3 s-1 in He, 1.0x10-2 s-1 in NO, 9x10-3 s-l in 02, and > 10-1 s-1 in NO2. The intensity of the NCO band decayed with an apparent fin'st-order rate coefficient of 3.9x 10 .3 s-l, in He, NO, and 02, and lx10-2 s-1 in NO2. Based on the initial intensity of the CN band, it is estimated that about 1% of the C02* cations are occupied CN species at the start of a transient experiment. Using this figure, the turnover frequency for the disappearance of CN species is estimated to be about 10-4 s-1 in NO and > lx10-3 s-1 in NO2. The latter figure is close to the turnover frequency measured during steady-state reaction of NO with CH4 at 450~ (8.4x10-4 s-l), lending support to the suggestion that CN species are reaction intermediates. NCO species adsorbed on A13§ react with NO2 at a rate that is an order of magnitude lower than the rate of reaction of Co2+CN species, and at an even lower rate when the reaction of NCO is with NO or 02. For this reason, it seems unlikely that NCO species play a primary role in the formation of N2 and CO2, the final products of NO reduction. 4. DISCUSSION The infrared observations presented in Figures 3 and 9 clearly demonstrate that CH4 reacts with adsorbed NO2 rather than adsorbed NO. Infrared spectroscopy also reveals that NO2 is readily formed when NO and 02 are present in the feed stream and that NO2 is more strongly adsorbed than NO. The temperature at which adsorbed NO2 begins to react with CH4, 350~ is identical to that at which significant NO conversion is observed during the reaction of NO with CH4 in the presence of 02. The higher lightoff temperature for NO reduction by CH4 in the absence of 02 (see Figure 1) is, thus, attributable to very low concentrations of adsorbed NO2, rather than to a higher barrier for the activation of CH4. Figure 10 illustrates a possible mechanism for the reduction of NO by CH4 in the presence of 02 based upon an amalgamation of elementary steps previously suggested in the literature [11, 13-16, 18, 20] and those deduced from the experiments presented here. The sequence begins with the adsorption of NO adsorption to form both mono- and dinitrosyl species (reactions 1 and 2). The later species undergo oxidation in 02 (reaction 3) to form adsorbed and gas phase NO2. The reaction of CH4 with adsorbed NO2 (reaction 4) is assumed to form an hydroxyl radical and adsorbed CH3NO. Weiner and Bergman [39] have reported the formation of CH 3NO and other nitrosoalkanes by migratory insertion of coordinated NO into Co-C bonds of cobalt alkyl species. Subsequent reaction of adsorbed nitrosomethane with either OH radicals or 02 (reactions 5) followed by the elimination of water (reaction 6) leads to the formation of adsorbed CN species. The nitrile species are hypothesized to react with NO2 to form N2 and CO2 (reaction 7), or with NO to form N2 and CO (reaction 8). The reaction of nitrile species with O2 (reaction 9) could result in the formation of NO and CO. While not

669 indicated, the CO released in reactions 8 and 9 is envisioned to undergo further oxidation to

CO2. The reaction sequence presented in Figure 10 is consistent with the mechanistic arguments given previously by Li et al. [ 15], who proposed that the first step in the reduction of NO by CH4 over Co-ferrierite is the reaction of gas-phase CH4 with adsorbed NO2, but differs in regard to the sequence of reactions leading to N2 and CO2. It is noted that the possibility of CN serving as a precursor to N2 and CO2 was suggested recently by Li et al. [32], based on studies conducted with Cu/ZrO2 and by Hayes et al. [41] based on studies conducted with CuZSM-5. The proposed mechanism is attractive in that it explains not only the manner in which NO2 initiates the reaction of CH4, but also the pathway to CO2 and N2. This mechanism would also explain why NO facilitates the combustion of CH4 by 02 [13, 18]. TPD experiments conducted in our laboratory have shown that Co-ZSM-5 will not adsorb 02, whereas it will adsorb NO2. If the product of the reaction of CH4 with NO2 is retained as an adsorbed species, then it is easy to see how NO2 (derived from the oxidation of NO) could facilitate the oxidation of CH4 by 02. 5. CONCLUSIONS In situ infrared observations show that the primary species present during the reduction of NO by CH4 over Co-ZSM-5 are adsorbed NO2 and CN. When 02 is present in the feed NO2 is formed by the homogeneous and catalyzed oxidation of NO. In the absence of 02, NO2 is presumed to be formed via the reaction 3 NO = NO2 + NzO. The CN species observed are produced via the reaction of methane with adsorbed NO2, and transient response studies suggest that CN species are precursors to N2 and CO2. A mechanism for the SCR of NO is proposed (see Figure 10). This mechanism explains the means by which NO2 is formed from adsorbed NO and the subsequent reaction sequence by which adsorbed NO2 reacts with CH4 and O2 to form CN species. N2 and CO or CO2 are believed to form via the reaction of CN with NO or NO2. CH3NO is presumed to be formed as a product of the reaction of CH4 with adsorbed NO2. The proposed mechanism explains the role of O 2 in facilitating the reduction of NO by CH4 and the role of NO in facilitating the oxidation of CH4 by 02. 6. ACKNOWLEDGMENT This work was supported by a grant from the Gas Research Institute. REFERENCES 1. Y. Li and J. N. Armor, Catalytic Reduction of NOx Using Methane in the Presence of Oxygen, U. S. Patent No. 5 149 512 (1992). 2. Y. Li, and J. N. Armor, Appl. Catal. B, 1 (1992) L31. 3. Y. Nishizaka and M. Misono, Chem. Lett., (1993) 1295. 4. Y. Li, J. Battavio and J. N. Armor, J. Catal., 142 (1993) 561. 5. R. Burch and S. Scire, Appl. Catal. B, 3 (1994) 295. 6. T. Tabata, M. Kokitsu and O. Osamu, Catal. Letc, 25 (1994) 393. 7. J.N. Armor and T. S. Farris, Appl. Catal. B, 4 (1994) L11. 8. R. Gopalakrishnan, P. R. Stafford, J. E. Davidson, W. C. Hecker and C. H. Bartholomew, Appl. Catal. B, 2 (1993) 165. 9. J.L. d'Itri and W. M. H. Sachtler, Appl. Catal. B, 2 (1993) L7. 10. Y. Li and J. N. Armor, Appl. Catal. B, 2 (1993) 239. 11.J.O. Petunchi and W. K. Hall, Appl Catal. B, 2 (1993) L17. 12. R. Burch and S. Scire, Appl. Catal. B, 3 (1994) 295.

670 13. F. Witzcll,G. A. Silland W. K. Hail, J. CataI., 149 (1994) 229. 14. Y. Li and J. N. Armor, J. Catai.,150 (1994) 376. 15. Y. Li, T. L. Slager and J. N. Armor, J. CataI., 150 (1994) 388. 16. E. Kikuchi and K. Yogo, Catal.Today, 22 (1994) 73. 17.A.D. Cowan, R. Dumplemann and N. W. Cant, J. Catal., 151 (1995) 356. 18. D. B. Lukyanov, G. Sill,J. L. d'Itriand W. K. Hall, J. Cat~., 153 (1995) 265. 19. Y. Li and J. N. Armor, Appl. Catal. B, 5 (1995) L257. 20. B. J. Adelman, T. Beutcl,G.-D. Lci, and W. H. M. Sachtler,J. Catal., 158 (1995) 327. 21.J.F. Joly, N Zanier-Szyldowski, S. Colin, F. Raatz, J. Suaussy and J. C. LavaIIey, Catal. Today, 9 (1991) 31. 22. K. A. Windhorst and J. H. Lunsford, J. Am. Chem. Soc. Farad. Trans.,97 (1975) 1407. 23.J.H. Lunsford, P. J. Hutta, M. J. Lin and K. A. Windhorst, Inorg. Chem.,17 (1978) 606. 24.M.C. Kung and H. H. Kung, Catal. Rcv. Sci. Eng., 27 (1985) 425. 25. W. Zhang, H. Yahiro, M. Iwamoto and J. Izumi, J. Chem. Soc. Farad. Trans., 91 (1995) 797. 26.M. Iwamoto, Y. Hidenori, N. Mizuno, W.-X. Zhang, Y. Mine, H. Furukawa and S. Kagawa, J. Phys. Chem., 96 (1992) 9360. 27. A. W. Aylor, S. C. Larsen, J. A. Reimer and A. T. Bell, J. Catai., 157 (1995) 592. 28. J. Valyon and W. K. Hall, J. Phys. Chem., 97 (1993) 1204. 29. T. E. Hoost, K. A. Laframboise and K. Otto, Catal. Lett., 33 (1995) 105. 30. C. -C. Chao and J. H. Lunsford, J. Am. Chem. Soc., 93 (1971) 71. 31. K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination Compounds, John Wiley and Sons, New York, 1986. 32. V. A. Bell, J. S. Feeley, M. Deeba and R. J. Farrauto, Catal. Lett., 29 (1994) 15. 33. C. Li, K. Bethke, H. H. Kung and M. C. Kung, J. Chem. Soc. Chem. Commun., (1995) 813. 34. F. Solymosi and T. Bansagi, J. Catal., 156 (1995) 75. 35. Y. Li and W. K. Hall, J. Phys. Chem., 94 (1990) 6145. 36. J. C. Evans, H. W. Rinn, S. J. Kuhn and G. A. Olah, Inorg. Chem., 3 (1964) 857. 37.J.W. Nebgen, A. D. McElroy and H. F. Klodowsky, Inorg. Chem., 4 (1965) 1796. 38. G. M. Begun and W. H. Fletcher, J. Molec. Spec., 4 (1960) 388. 39. I. C. Hisatsune and J. P. Devlin, Spectrochimica Acta, 16 (1960) 40 I. 40. W. P. Weiner and R. G. Bergman, J. Am. Chem. Soc., 105 (1983) 3922. 41. N. W. Hayes, W. Grunert, G. J. Hutchings, R. W. Joyner, and E. S. Shpiro, J. Chem. Soc. Commun. (1994) 531.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

671

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

Precious metal loaded In/H-ZSM-5 for reduction of nitric oxide with methane in the presence of water vapor M. Ogura and E. Kikuchi* Department of Applied Chemistry, School of Science and Engineering, Waseda University, 3-4-10kubo, Shinjuku-ku, Tokyo, Japan The catalytic activity of In/H-ZSM-5 for the selective reduction of nitric oxide (NO) with methane was improved by the addition of Pt and Ir which catalyzed NO oxidation, even in the presence of water vapor. It was also found that the precious metal, particularly Ir loaded In/H-ZSM-5 gave a low reaction order with respect to NO, and then showed a high catalytic activity for the reduction of NO at low concentrations, if compared with In/H-ZSM-5. The latter effect of the precious metal is attributed to the enhancement of the chemisorption of NO and also to the increase in the amount of NO2 adsorbed on In sites.

1. Introduction Emission of nitrogen oxides (NOx) in combustion exhausts is a serious problem that should be solved in order to protect the earth from acid rain, or other environmental concerns [1]. Recent progress in catalytic removal of NOx is the emergence of an impressive SCR process, which takes an advantage of using unburned hydrocarbons in exhausts as the reductant. Since the discoveries by Iwamoto et al. [2] and Held et al. [3] that the reduction of NO with hydrocarbons could be catalyzed by Cu-ZSM-5 zeolite, interest in this field has been increasing and various catalysts have been proposed for this reaction [4-10]. However, no catalysts having enough activity and selectivity in a practical use have yet been developed, particularly the catalyst having enough durability against the inhibition by water vapor, which is inevitably contained in practical combustion exhausts. In our previous work [11], it has been shown that the reduction of NO with CH, on GaJ and In/H-ZSM-5 catalysts selectively proceeds in the following two stages: NO + 1/2 02 --> NO2 NO2 + CH, + NOx --> N2 + COx + H20

[ on zeolite acid sites ] [ on Ga or In sites ]

(1) (2)

The catalytic activity for NO oxidation [reaction(1 )] was strongly inhibited by water vapor, because this reaction occurs on Lewis acid sites of zeolite as

672

proposed [12]. Although the reduction of NO2 with CH4 [reaction (2)] was also inhibited by water vapor, the catalytic activity of In/H-ZSM-5 for this reaction was more durable against H20 than that of Ga/H-ZSM-5 [11]. Furthermore, the catalytic activity of In/H-ZSM-5 for the reaction of NO-CH4-O2 even in the presence of water vapor was found to be improved by the addition of precious metals such as Pt, Rh, and Ir, because these metals catalyzed NO oxidation instead of the Lewis acid sites of the zeolite [13]. The effects of precious metals on In/H-ZSM-5 was found not only to simply catalyze NO oxidation but also to enhance NOx chemisorption. It is noted that NO conversion on the Ir/In/H-ZSM-5 exceeded NO2 conversion in NO2-CH402 reaction on In/H-ZSM-5, when the concentration of NOx was decreased [14]. This study shows the catalytic activities of In/H-ZSM-5 promoted by precious metals for the removal of low concentration NOx and the promotive effects of the precious metal will be discussed.

2. Experimental Na-ZSM-5(a molar SIO2/AI203 ratio=23.8) provided by Tosoh Corp. was used. In(4wt%)/H-ZSM-5 and Ir(lwt%)/H-ZSM-5 catalysts were prepared by the ion exchange method using NH4-ZSM-5 derived from the Na-ZSM-5 with aqueous solutions of In(NO3)3 at 368 K for 8 h and IrCI(NH3)5CI2 at room temperature for 24 h, respectively. Addition of precious metals, lwt% platinum and iridium to In/H-ZSM-5 was carried out by impregnating the In/NH4-ZSM-5 in aqueous solutions of Pt(NH3)4CI2 and IrCI(NH3)5CI2, respectively. The catalysts were calcined at 813 K for 3 h. Reaction was carried out in a fixed-bed flow reactor mainly by passing a reactant gas mixture of 100-1000 ppm NOx (NO or NO2), 1000 ppm CH4, 10% 02 and 0 or 5% H20 in He at a rate of 100 cm3(STP)omin -1 over 0.1 g of catalysts (GHSV = 36000 h-l). Kinetics studies employed higher GHSV to obtain low levels of NOx conversion below 30%. Water vapor was admitted by passing He through a saturator heated in water bath controlled at 313 K, and the wet He was mixed with other reactants. The reactant stream-line was heated at a higher temperature than the temperature of the saturator to avoid H20 condensation. Reaction products were analyzed by means of on-line gas chromatography with a TCD detector and chemiluminescence NOx analysis. The catalytic activity was evaluated by the conversion of NOx into N2. Chemisorption of NOx was measured by the pressure swing adsorption method reported by Zhang et al. [15]. A gas mixture containing 100 - 1000 ppm NOx was admitted on to 0.1 or 0.2 g catalyst at 673 K, and the concentration of NOx passing through the catalyst bed was detected using a NOx analyzer. Typical response curve and breakthrough point are shown in Fig. 1. The amounts of adsorbed NOx were calculated from these breakthrough curves: the areas shown by a and b correspond to the amounts of totally adsorbed and reversibly adsorbed NOx, respectively.

673

He

I_ F .

(1) u) E 0 Q.

iI

ii

_1 -I

NOx+O2 m

i

ii

ii

al

ii

i

i

m

ii

ii

ii

m

m

am

lille

I

He

I

i

.

e e

e e e e ! I |

n,-

e o | ! o e e | | i | | |

Time Figure 1. Measurement of NOx adsorption" a, corresponds to the total amount of adsorbed NOx; b, the amount of reversibly adsorbed NOx.

3. Results Figure 2 shows the effect of NOx concentration on the conversion of NOx reduced by CH4 in the presence of 5% H20. In the NO-CH4-O2 system, In/HZSM-5 showed low catalytic activity in the whole range of NO concentration. On the other hand, this catalyst was active for the NO2-CH4-O2 reaction, while the conversion of NO2 decreased with decreasing concentration of NO2. The catalytic activity of In/H-ZSM-5 for the reduction of 1000 ppm NO was enhanced by the addition of Ir and Pt almost to the level of NO2 reduction on In/H-ZSM-5, indicating that these precious metals worked as the catalytic sites for NO oxidation, which is a necessary step for NO reduction with CH4. With decreasing NO concentration to 100 ppm, however, the increase in NO conversion was observed on Ir/In/H-ZSM-5 and the conversion of NO exceeded that of NO2 on In/H-ZSM-5. This can not simply be explained by the catalytic activity of Ir for NO oxidation. Kinetic parameters for NOx reduction are summarized in Table 1. It is obvious that the addition of Ir to In/H-ZSM-5 led to the decrease in reaction orders with respect to NO, CH4, and 02 in the NO-CH4-O2 reaction. The decrease in the order for NO can explain that Ir/In/H-ZSM-5 was effective for the reduction of NO at low concentrations. On the contrary, the reaction orders with respect to NO2, OH4, and 02 in the NO2-CH4-O2 reaction were not significantly changed by the addition of Ir. The retarding effect of CH4

674

100 o~

80

OJ

z O *-

60

._o L_

>

40

0 z

20

co o x

f 0

A

A

I

I

I

i

I

200

400

600

800

1000

NOx concentration / ppm Figure 2. Catalytic activities of In/H-ZSM-5 (e), Pt (4) and Ir ( 0 ) l o a d e d In/H-ZSM-5 for NO reduction and that of In/H-ZSM-5 (n) for NO2 reduction in the presence of 5% H20 as a function of NOx concentration. Catalyst weight, 0.1 g. Reaction temperature, 773 K. Table 1. Summary of kinetic data for NOx-CH4-O2 reaction on In/H-ZSM-5 and Ir/In/H-ZSM-5 catalysts. Reaction

NO-CH4-O2

NO2-CH4-O2

Catalyst

/ m~

ro

19

Reaction orders ,2 with respect to 1

NO

NO2

02

OH4

In/H-ZSM-5

1.5 x 10-6

0.80

-

0.80

0

Ir/In/H-ZSM-5

6.1 x 10-6

0.65

-

0.13

-0.18

In/H-ZSM-5

1.0 x 10-5

0.44

0

0.47

Ir/In/H-ZSM-5

1.5 x 105

0.47

0

0.44

I

19 r0, the initial rate reaction under the standard conditions: NOx, 100 ppm; CH4, 1000ppm; and 02, 10%. 29 Concentrations: NOx, 100 - 1000 ppm; CH4, 500 - 2000 ppm; 02, 5 - 12 %. Reaction temperature, 673 K. Catalyst weight: 5.0 - 40 mg for In/H-ZSM-5; 3.0 - 7.0 mg for Ir/In/H-ZSM-5.

675

observed as the negative order of reaction in the reduction of NO on Ir/In/HZSM-5 might be due to the competitive oxidation of NO and CH4 on Ir sites. This effect was not found in the reduction of NO2. Chemisorption of NOx on In/H-ZSM-5 and Ir/In/H-ZSM-5 was measured as a function of NOx concentration. The isotherms of reversible adsorption are shown in Fig. 3. A larger amount of NO2 was adsorbed on In/H-ZSM-5 than NO. Chemisorption of NO was remarkably enhanced by the addition of Ir leading to a larger amount of NO being adsorbed on Ir/In/H-ZSM-5 than NO2 on In/H-ZSM-5. Figure 4 shows the amounts of adsorbed NO from the feed of NO, NO-O2, NO2, or NO2-O2 on In/H-ZSM-5, Ir/H-ZSM-5, and Ir/In/H-ZSM-5. It is apparent from these results that Ir/H-ZSM-5 adsorbed little NOx. NO could hardly be adsorbed on every catalyst in the absence of 02. It is interesting to note that Ir/In/H-ZSM-5 adsorbed larger amounts of NOx from the mixture of NO and 02 than from NO2, and it was also larger than the amount of NOx adsorbed on In/H-ZSM-5 from NO2.

40 ! "7 O)

o

E

30

0 x

~ x 0 z

20

.~

10

L_

0

0

=F,'-

0

I,

I

A

I

I

200 400 600 800 10001200 NOx concentration / ppm

Figure 3. Isotherms of reversible adsorption at 673 K for NO on In/H-ZSM-5 (e) and Ir/In/H-ZSM-5 (0) and for NO2 on In/H-ZSM-5 (m) and Ir/In/H-ZSM-5

(n).

Catalyst weight: 0.2 g for In/H-ZSM-5; 0.1 g for Ir/In/H-ZSM-5.

676

Catalysts" Ir/H-ZSM-5

NO N O + 02 NO2 NO2 + 02 NO

In/H-ZSM-5

Ir/In/H-ZSM-5

~['-]NO + 02

"I////A rl///i//~

IN 02 IN 02 + 02

7] NO Z/////I////IA ~//I IN02 ~////////,/A

0

5

NO + 02

INO~2+ 02 10

15

Amount of adsorbed NOx / 10"6 mol.(g-cat)1 Figure 4. Comparison of chemisorption of NOx at 673 K from various kinds of NOx mixtures admitted on to In/H-ZSM-5, Ir/H-ZSM-5, and Ir/In/H-ZSM-5 catalysts: E] parts correspond to the reversible adsorption (b in Fig. 1). NOx, 100 ppm; 02, 0 or 10%. Catalyst weight: Ir/In/H-ZSM-5, 0.1 g; In/H-ZSM-5, 0.2 g; Ir/H-ZSM-5, 0.1 g.

4. Discussion As shown in Table 1, the reaction order with respect to NO2 on In/H-ZSM-5 was smaller than that of NO. This is in accordance with the proposed reaction sequence that NO is firstly oxidized to NO2 and the NO2 reacts with CH4. Coincidence in the order of reaction for NO2 between In/H-ZSM-5 and Ir/In/HZSM-5 catalysts means that NO2 react on a common active site which should be In species. Chemisorption data shown in Fig. 4 show that either NO or NO2 was hardly adsorbed on Ir site. Furthermore, chemisorption of pure NO was negligibly small on In/H-ZSM-5 and Ir/In/H-ZSM-5 at 673 K, while NO in the presence of 02, as well as NO2, could significantly be adsorbed. These results also support our supposition that chemisorption of NO2 is important and that of NO is less important on these catalysts. In the presence of 02, chemisorption of NOx, both NO and NO2, was enhanced by the addition of Ir on to In/H-ZSM-5.

677

It has been reported that GaJ and In/H-ZSM-5 had low activities for the dissociative adsorption of oxygen, and this result can explain why these catalysts show high selectivities for NO reduction with hydrocarbons [16]. Reaction order with respect to 02 was lowered by the addition of Ir to In/HZSM-5 in the NO system. It is noteworthy that the added precious metals promote the adsorption of oxygen, which is an indispensable component for HC-SCR. 02 is activated on Ir to adsorb dissociatively. On the other hand, the similar reaction orders were obtained on In/H-ZSM-5 and Ir/In/H-ZSM-5 in the NO2 system. NO2 adsorption on In is not inhibited by the adsorbed oxygen on Ir. It is previously reported that the CH4 selectivity to NO reduction was lowered by added precious metals, Ir and Rh, which were active for CH4 oxidation with oxygen [13]. The selectivity of Ir/In/H-ZSM-5 decreased when the concentration of CH4 increased, while it increased with increasing NO concentration. Therefore, 02 can easily be adsorbed on Ir and the adsorbed oxygen can activate not only NO but also CH4. These results suggest that NO

"7,

5

4

x

3

o;2 -s o

1

0

0

2

4

6

8

10

12

1/(NOx concentration) x 10 3 / p p m 1 Figure 5. 1/P - 1/V plot for NO adsorption on In/H-ZSM-5 (o) and Ir/In/H-ZSM-5 (0) and for NO2 adsorption on In/H-ZSM-5 (ll) and Ir/In/H-ZSM-5 (El). Reaction conditions and symbols are the same with those in Fig. 3.

678

and CH4 competitively react with adsorbed oxygen on Ir, resulting in the inhibition of the NO oxidation and in the negative order with respect to CH4 in the NO system on Ir/In/H-ZSM-5. It is noted that the amount of NO adsorption on Ir/In/H-ZSM-5 in the presence of 02 was remarkably larger than that of NO2 on In/H-ZSM-5. To understand the chemisorptive property of Ir/In/H-ZSM-5 in a comparison with that of In/H-ZSM-5, the chemisorption data were analyzed in detail. The data shown in Fig. 3 are relatively well fitted to the Langmuir isotherm, as shown in Fig. 5. From these relations, the equilibrium constant (K) and the amount of NO2 adsorbed at saturation (Vo) were determined according to the following equations: V = KPVo 1 +KP

(3)

1 = 1+1 x 1 V Vo KVo P

(4)

Here, 0 and V represent the surface coverage and the amount of adsorbed NOx, respectively. Calculated K and V0 are summarized in Table 2. If the important adsorption site on these catalysts consists of In species, and if NO2

Table 2. Equilibrium constant of adsorption (K) and adsorbed amount at saturation (Vo) for NOx at 673 K on In/H-ZSM-5 and Ir/In/H-ZSM-5. []

I

NOx in the reactant feed

V0 / 10"6 mol (g-cat)-1

K / 103

NO

9.3

1.4

NO2

21

2.0

i

In/H-ZSM-5

i

Ir/In/H-ZSM-5

NO

57

1.6

NO2

63

1.4

Reactant feed: NOx, 100 - 1000 ppm; 02, 10%. Catalyst weight: 0.2 g for In/H-ZSM-5; 0.1 g for Ir/In/H-ZSM-5.

679

is a sole common adsorbed species, then the adsorption equilibrium constant would be unique. The observed values lie in the range of (1.6 4- 0.4)x10 -3. Taking into consideration the scatter in the observed data probably due to interconversion between NO and NO2 during chemisorption measurements, these values could be regarded almost within experimental errors. The value of Vo sounds more meaningful. The adsorbed amount of NOx at saturation should correspond to the number of adsorption sites. These are in the range of 1 x 10 -5 to 2 x 10 -5 mol-g-catalyst -1 for In/H-ZSM-5. They are extremely small compared with the amount of In in the catalyst (4.8 x 10 -2 mol-g -1). By the addition of Ir, Vo became 3 to 6 times greater. This seems to explain why Ir/In/H-ZSM-5 gave a greater rate of NOx reduction than In/HZSM-5, as shown by ro in Table 1. Although the exact reason of the larger Vo for Ir/In/H-ZSM-5 is not certain at present, a possible explanation is that NO can diffuse into zeolite pores more easily than NO2 and diffused NO can be oxidized on Ir in the pore to NO2 which is then reduced on In sites also existing in the pore. In our previous work [11], it has been shown that the reduction of NO with CH4 on Ga and In/H-ZSM-5 catalysts proceeds through the reactions (1) and (2), and that CH4 was hardly activated by NO in the absence of oxygen on these catalysts. Therefore, NO2 plays an important role and the formation of NO2 is a necessary step for the reduction of NO with CH4. In the works of Li and Armor [17] and Cowan et al. [18], the rate-determining step in NO reduction with CH4 on Co-ferrierite and Co-ZSM-5 catalysts is involved in the dissociative adsorption of CH4, and the adsorbed NO2 facilitates the step to break the carbon-hydrogen bond in CH4. It is suggested that NO reduction by use of CH4 needs the formation of the adsorbed NO2, which can activate CH4. From the result that the conversion of NO at a low concentration (100 ppm) on Ir/ln/H-ZSM-5 exceeded that of NO2 on In/H-ZSM-5, it can be concluded that there are some correlations between the increase in the quantity of adsorbed NO2 and the enhancement of the catalytic activity for the reduction of NO with CH4.

4. Conclusion The catalytic activity for the selective reduction of NO with CH4 was significantly enhanced by the addition of precious metals, particularly Ir, to In/H-ZSM-5. The role of added Ir was not only to promote NO oxidation which is a necessary step for NO reduction, but also to enhance the abilities of the catalyst to adsorb NO and to increase the amount of NO2 adsorbed on in sites.

680

5. Acknowledgment This work was supported by the Grant-in-Aid for Scientific Research on Priority Areas from the Ministry of Education, Science, and Culture of Japan.

References [1] H. Bosch and F. Janssen, CataL Today, 2, 369 (1988). [2] M. Iwamoto, H. Yahiro, Y. Yuu, S. Shundo, and N. Mizuno, Shokubai(Catalyst), 32, 430 (1990). [3] W. Held, A. KSnig, T. Richter, and L. Puppe, SAE Paper, 1990, 900496. [4] H. Hamada, Y. Kintaichi, M. Sasaki, and T. Itoh, App/. Cata/., 64, L1 (1990). [5] M. Misono and K. Kondo, Chem. Lett., 1991, 1001. [6] K. Yogo, M. Ihara, I. Terasaki, and E. Kikuchi, Catal. Lett., 17, 303 (1993). [7] Y. Li, P. J. Battavio, and J. N. Armor, J. Catal., 142, 561 (1993). [8] J.O. Petunchi, G. Sill, and W. K. Hall, App/. Catal. B2, 303 (1993). [9] Y. Nishizaka and M. Misono, Chem. Lett., 1993, 1295 [10] K. Yogo, M. Umeno, H. Watanabe, and E. Kikuchi, Cata/. Lett., 19, 131 (1993). [11] E. Kikuchi and K. Yogo, CataL Today, 22, 73 (1994). [12] J. G. M. Brandin, L. A. H. Anderson, and C. U. I. Odenbrand, CataL Today, 4, 187 (1989). [13] E. Kikuchi, M. Ogura, N. Aratani, Y. Sugiura, S. Hiromoto, and K. Yogo, Catal. Today, in press. [14] M. Ogura, S. Hiromoto, and E. Kikuchi, Chem. Lett., 1995, 1135. [15] W.X. Zhang, H. Yahiro, N. Mizuno, J. Izumi, and M. Iwamoto, Langmuir, 9, 2337 (1993). [16] T. Tabata, M. Kokitsu, and O. Okada, Catal. Lett., 25, 393 (1994). [17] Y. Li and J. N. Armor, J. Catal., 150, 376 (1994). [18] A.D. Cowan, R. D0mpelmann, and N. W. Cant, J. CataL, 151,356 (1995).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Atmiversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

681

Interfacial RhOx/CeO 2 sites as locations for low temperature N20 dissociation

J. Cunningham%J.N. Hickey%R. Cataluna~'%J-C Conesab,J. Sofia b and A.Martinez-Arias a Physical Chemistry Laboratories, University College Cork Ireland b CSIC Institute for Catalysis, Universidad Autonoma, Cantoblanco, Madrid, Spain.

SUMMARY Temperatures required for extensive N20 dissociation to N 2, or to N 2 plus 02, over 0.5% RhOx/CeO 2 materials, and over polycrystalline Rh203 or CeO2, are compared for preoxidised and for prereduced samples on the basis of conversions achieved in pulsed-reactant, continuous-flow and recirculatory microcatalytic reactors. Influences of sample prereduction or preoxidation upon those measurements and upon results from parallel ESR and FTIR studies of N20 interactions with such materials are presented and compared. Over partially reduced 0.5% RhOx/CeO 2 materials complete dissociation of N20 pulses to N 2 plus 02 is obtained at temperatures 50-100 ~ lower than those required for extensive dissociation over prereduced Rh203. Furthermore, N: was the sole product from the latter. Higher ongoing N20 conversions to N 2 plus 02 at 623 K over 0.5% Rh/CeO2 in pulsedreactant than in continuous-flow mode point to regeneration of active sites under helium flushing between pulses. The TPD profile for dioxygen release from Rhodia containing samples at temperatures 350-550 K is presented. ESR measurements reveal complementary effects of outgassings at temperatures, Tv, > 573 K upon the availability at RhOx/CeO 2 surfaces of electron-excess sites reactive towards N20. Differences from observations over Rh203 and CeO2 can be understood by attributing the low-temperature activity of RhOx/CeO 2 to electron excess sites at microinterfaces between the dispersed Rhodia component and the Ceria support.

1. INTRODUCTION Since nitrous oxide, N20, is a designated "greenhouse" gas, and may contribute to depletion of the ozone layer, its removal from emissions to atmosphere is desirable [1]. However, there are several reports that N20 can be formed at low selectivity as an undesirable by-product of NO+CO conversions during the initial warm-up-from-cold periods in three-way-catalytic (TWC) converters or components thereof [1-3]. TWC's commonly contain Rhodium and Ceria and although N20 dissociation over Rh203 has been extensively studied [4], the following are among mechanistic possibilities as yet

682 incompletely resolved concerning N20 formation and its eventual catalytic removal over binary Rhodimn+Ceria materials: (a) relative importances of the reduced and oxidised forms of Rhodium and Ceria components and of microinterfaces between them in providing active sites [5-7]; (b) relationship of rates for N20 adsorption and dissociation on clean Rh/CeO2 surfaces to corresponding rates on surfaces having some active sites blocked by adsorbed O2 [4]; and (c) direct or indirect influences of anion-vacancies/lattice-oxygens from the ceria support in promoting/inhibiting catalytic dissociation [8]. The present work seeks insights into those possibilities, mainly through comparative studies of N20 interactions with 0.5% Rh/CeO 2, CeO2 and Rh203 samples in both preoxidised and prereduced forms. Comparison between their "initial" and "steady-state" activities have been obtained with a microcatalytic reactor system designed to switch N20 reactant from pulsed to continuous flow. Complementary information concerning the nature of species formed on surfaces of the materials upon contacting N20 and/or 02 with samples activated by prior vacuum-outgassing at increasing temperatures (Tv) has been obtained from measurements by electron-spin-resonance, (ESR) and FTIR, and compared with previous observations of electron localization by 02 at such surfaces. [9,10].

2. EXPERIMENTAL Materials and Characterizations: Rh203 powders commercially available from Aldrich, Rh203(A), and Johnson Matthey, Rh203 (JM) with BET=I6 m2g-~, were used after calcinations/reductions indicated in the text. Wet impregnation of Rhodium acetylacetonate onto CeO 2 powder (BET---IIO m2g~ from Rhone Poulenc) from solution in tetrahydrofuran or methanol yielded 0.5% RhOx/~CeO2 and 4% RhOx/wCeO2 after calcination in 02 at 823 K. [8]. Identical procedures were followed to achieve dispersions of oxidised Rhodium species upon 27% CeO2-AI203. Analagous procedures using Rhodium nitrate as precursor [9] were employed in preparing samples used in ESR and FTIR studies. Coprecipitation by NH4OH from solutions containing nitrate(s) of Rhodium and/or Cerium was the first step of an alternative preparation yielding powders having rhodium ions initially dispersed throughout Ceria e.g. 0.5% RhOxPCeO2. TPR profiles of aliquots precalcined at 823 K were obtained under 3% H2/Argon, with particular attention to position and magnitude of any rhodia-related features at 373 ~ 473 K. Use of 0.5% RhOx/sCeO2 material prepared by wet impregnation of Rhodium acetylacetonate onto CeO2 (rp) which had been sintered at 1273 K overnight, was found necessary in studies of oxygen temperature programmed desorption from 0.5% RhOx/CeO2, otherwise 02 desorption from surface oxygens of h.s.a. CeO2 in the range 473-773 K obscured a rhodia-related O2-desorption feature in that range. X-band ESR spectra were taken at 77 K on a Bruker system (ER 200D), g-values being calibrated with DPPH (g = 2.0036). Ca. 20 mg of sample were placed in a quartz cell with double greaseless stopcocks, and were subjected to outgassing or gas adsorption; prior to taking the spectra, any excess 02 was pumped out at 77 K where necessary to avoid magnetic dipolar broadening of ESR lines. FT-IR spectra were taken with a Nicolet 5ZDX spectrometer, accumulating 200 scans at 4 cm "~ resolution. Self-supported wafers of ca. 25 mg/cm 2 were placed in a high vacuuna cell, provided with NaCI windows, for outgasing and adsorption treatments. All IR data were obtained with the cell at room temperature, and the sample at the temperature induced by the IR beam (estimated to be .~ 320 K). Outgassing and gas treatments of samples for spectroscopic experiments were made in a

683 conventional high vacuum line capable of keeping a dynamic vacuum of I O -3 Nm 2. Research grade 02 and N20 (SEO) were purified by freeze-thaw cycles before use.

Microcatalytic Reactor Systems: Most catalytic results were obtained with a flow reactor system which operated at I atm total pressure, comprised mainly of helium or argon carrier gas with only low content (usually 3%) of N20 therein. With the aid of a 10-port valve the reactant flow could rapidly be switched between: Configuration A in which individual N20 pulses passed over the catalyst and thence straight to a pair of GC columns for on-line analysis, upon completion of which a second pulse could be delivered etc. During each 1015 min. interval between pulses the sample remained at reaction temperature under a flow of pure helium; Configuration B in which a continuous flow of 3% N20/He became established over the same sample, but with gases exitting continuously therefrom being sampled only periodically to the GC columns for analysis. A separate, evacuable, all-glass recirculating reactor system operating at low pressures 1-5 mbarr and equipped for mass spectrometric analysis was used to investigate possibilities of oxygen isotope exchange between lattice- ~60 and N2~80, in parallel with temperature programmed catalytic dissociation (TPCD) of the nitrous oxide.

3. RESULTS AND INTERPRETATION

3.1 Microcatalytic Reactor Studies Rh203: Parts (A) and (B) of figure 1 illustrate results obtained from delivery of the first I0 pulses (A), and then a very large number of N20 pulses (B), at 673 K to a sample of Rh203 powder prereduced in H2 at 473 K. The latter treatment was chosen on the basis of a TPR profile indicating reduction of Rh203 to be complete at 473 K and with a view to facilitating later comparison of N20 pulse interactions over prereduced Rh203 with those over prereduced 0.5% Rh/PCeO2 and PCeO:. In Fig. 1A, titration-like trends are evident in the level of N20 conversion achieved upon delivery of the first ten N20 pulses over prereduced Rh203 at 673 K: conversion to N 2 being complete in the first pulse, but then declining rapidly towards zero for pulses "4 --> ~10, and conversely for survival of nondissociated N20. This behaviour could be understood in terms of the development of an inert Rhodium plus Oxygen overlayer on the metallic Rhodium, thereby passivating it. Such inert overlayers have been reported in other systems [I1], sometimes as precursors of identifiable Rh203 overlayers [12]. Data in Fig. 1B for the N:O conversions achieved in later N20 pulses of the long pulse sequence at 673 K over prereduced Rh203 demonstrate how, following decline of N20dissociation activity almost to zero for pulses ~8 to ~12, such activity increased progressively until reaching 80% after 85 pulses. Peak area comparisons indicated that the product was predominantly N 2 and certainly not N2+l&O2, thereby pointing to an autocatalytic oxidation of Rhodium by N20 rather than catalytic N20 dissociation. Bearing in mind that the sample was exposed to flows of Helium at 673 K during each 10 min interval between pulses, and during two overnight interruptions in delivery of the long train of pulses, ample opportunity existed for progression of structural reorganization of any initial inert oxygen overlayer a Rhodia surface layer [12] at 673 K and for thickening of such rhodia overlayer in an autocatalytic reaction with N20. Results similar to Fig. 1B did not develop in N,O pulse sequences at 423 or 523 K over prereduced Rh203. Furthermore, periodic analyses of exit

684

Figure 1. Relative GC peak areas for N2-only product and any residual N20 reactant from passage of N20 pulses at 673 K over prereduced Rh203: (1A) declining conversion in pulses 1 -~ 10; (1 B) reversal of initial decline for pulses 12 - 85.

gases from a continuous flow of N20 over Rh203 prereduced at 473 K showed no measurable conversion to N2 or O 2 during 2 h on-stream at 623 or 673 K. The following results were obtained from passage of N20 pulse sequences over Rh203 aliquots which had not been prereduced: (i) full ongoing conversion to N2+1/~O2 was achieved at 623 K over Rh203 previously flushed for 1 hr with helium at that temperature, whereas ca. 30% ongoing conversion resulted at 573 K; (ii) ongoing conversion at 28% level to N2+1/~O2 was observed from N20 pulses at 573 K over an aliquot which had been preoxidised for 1 hr at 773 K under a flow of IO% OJHelium. 0.5% RhOx/CeO2: At reaction temperatures ca. 50 ~ lower than those indicated above as necessary for Rh203, the following features were observed over aliquots of coprecipitated 0.5% RhOx/CeO 2 materials pretreated and tested in conditions closely similar to those used for Rh203 aliquots: (i) analyses of exit gases from N20

pulses delivered at 623 K over prereduced aliquots showed complete dissociation to N2 as sole product from the first ten pulses, whereupon a rapid switch-over at pulses ~10-12 was observed in composition of gases emitting from complete N20-pulse dissociation to yield N2+~AO2 from pulses ~13-~30 (cf. Fig. 2A). This switch-over could be understood in terms of a predominance of surface-site reoxidation process, N20+S%ed ~ N2+O"/Sox, durmg the first ten pulses [13] after which came an onset of N20-dissociation to N2+~A02 proceeding via redox cycling at such sites via N20+0n/Sox ~ N2+02+S"~d etc; (ii) Figure 2B illustrates that analyses of exit gases after passage of N20 pulses at 623 K over preoxidised aliquots provided evidence for presence of N 2 plus O2 product already after the second pulse and continuing to be present in subsequent pulses thereafter with undissociated N20 at a level indicating only 50% ongoing conversion. (iii) Analyses of exit gases from N20pulses delivered at T~x=523 or 423 K over prereduced aliquots yielded only an initial dissociation to N2 as sole product for the first 7 or 4 pulses respectively i.e. similar to Fig. 1A. Analyses of exit gases after higher pulse numbers at these relatively low temperatures over the prereduced aliquots showed only N20 with zero N2 or O 2 product, thereby

685 establishing an absence of ongoing dissociation which contrasted with results from N20pulses over the same material at 623 K. Tests made in the pulsedN20 mode at 423, 523 and 623 K upon a "reference" CeO2 material prepared by NH4OH precipitation from cerium nitrate and pretreated in similar fashion to the 0.5% RhOx/CeO 2 material did not discover any activity for ongoing N20 dissociation, but only a small initial dissociation to N 2 similar to Fig. IA. Following determination of the % conversions of N20 pulses to N 2 plus 02 at 623 K just noted over prereduced or preoxidised aliquots of 0.5% RhO]CeO 2, the reaction system was instantly reconfigured to flow N20 reactant continuously over the samples at that temperature. In each case the result was an approximate halving of the N20 conversion relative to that achieved in pulsedN20 configuration, despite similar flow rates in both configurations and continuation of flow for up to 2 h. It seemed probable from this halving, and from an abovementioned enhancing effect of prior helium flush at 623 K upon activity of Rh203 for N20 dissociation to N2 plus 02, that the periodic 10 min flushed with pure helium at 623 K experienced by 0.5% RhOx/CeO ~ between Figure 2. Relative GC peak areas of exit gases from N20 pulses over 0.5% RhOx/CeO2 at 623 K: (2A) Complete conversion to N2-only from pulses 1 ~ 10 over prereduced coprecipitated material followed by a switch to N 2 plus 02 products for pulses 12 ~ 27; (2B) partial conversion at 623 K over the same material when preoxidised; (2C) complete conversion to N 2 plus 02 products for pulses 1 ~ 21 over material prepared by wet impregnation.

686 N20 pulses, contributed to formation/regeneration of Rhodium-containing sites having activity for dissociation of each incoming N20 pulse to N 2 plus 02 at 623 K. Possibilities for such formation/regeneration of active sites include the diffusion/desorption of blocking species away from Rhodium-containing sites, which are further considered below in respect of (N202) "n and 02. Part c of Fig. 2 demonstrates that ongoing complete dissociation of N20 pulses at 623 K to N2+1/202 occurred over 0.5% RhOx/CeO 2 prepared by wet impregnation rather than the coprecipitation method used for the material used in part a of the figure. The extensive dissociation over material prepared by wet impregnation is consistent with greater localization of the Rhodium component in surface and near-surface regions. Further support for this emerged from differences in extent of dissociation at pulse numbers greater than ~15 in N20 pulse sequences at 523 K: no ongoing dissociation then being detectable over the coprecipitated material, whereas dissociation to N2 plus 02 continued at pulse numbers "16 -~ '27 over the material originated by wet impregnation. Evidence was obtained for significant Oxygen isotope exchange between N2~80 and ~60" species on the surface of a preoxidised sample of 4% Rh/wCeO2-Al203. Data were obtained by introducing N2~80 at 3 mbarr pressure into a recirculatory reactor system containing the preoxidised catalyst and applying a temperature ramp at 10 ~ min ~ to the reactor segment whilst analysing by mass spectra at 30 sec intervals the composition of gases recirculating in the system. Results showed onsets of a decrease in PCN2~80) @ m/e=46, together with an increase in P(N2) @ m/e=28 at ca. 520 K, confirming dissociation of N2~80 at this and higher temperatures. An interesting effect, not revealed by other studies, was indicated by parallel onset at 520 K of another process producing a gaseous species, having rn/e=44 and identified as N2~60. (The possibility that this might be due to CO2 could be rejected on the basis that no equivalent rise in signal at m/e=44 was observed to onset at this temperature from a CeO2-AI203 sample subjected to identical in-situ preoxidation at 823 K prior to testing isotopically with N2J80 in conditions similar to those for the 4% Rh/CeO2-A1203 material). Comparison of the rates of rise @ m/e=28 and m/e=44 with temperature in the range 520-773 K showed that the process yielding N2160 proceeded with ca. 50% of the efficiency of that yielding N2~g~ product formation, possibly via (N2~sO~60)" intermediates at Rhodia-containing sites. The observed increase in P(N2160) up tO ca. 725 K implies presence of some such surface precursor up to those temperatures and makes it a candidate for blockage of Rhodia-containing active sites. 3.2. Oxygen Desorption Profiles for temperature programmed desorption of O2(TPD) were compared for 0.5% RhOx/CeO2, 4% RhO~/sCeO2, Rh203 and sCeO 2 materials pretreated in-situ to a flow of 02 to 823 K, cooled down in O2 to 300 K and flushed overnight in helium at 300 K to remove physically adsorbed oxygen. Results are summarised in Fig. 3, which shows that all the Rhodia-containing materials yielded a detectable TPD feature in the range 350 -~ 550 K. Facile detection of that Oxygen TPD feature from Rh203 (plot c) contrasted with greater difficulty for its detection from 4% RhO• (plot d) because of a large background 02 desorption from that high surface area Ceria support. Attainment of the indicated clear differentiation between zero desorption from (plot a) and a small,but observable feature from 0.5~ RhO• (plot b) only became possible by dispersing 0.5% Rhodium upon CeO 2 presintered overnight at 1250 K and by using the sintered CeO2 as reference (plot a).

687

3

t

J

Aldrich RhOx

Release of ~602 from preoxidised 4% RhOx/CeO 2 at temperatures > 473 K into a low pressure of Oxygen-18 enriched 1802 was reported previously on the basis of measurements made with the recirculatory reactor system [8]. An analagous experiment in the present study, featuring a low pressure of N2160 in contact with preoxidised 4% RhOx/CeO2 while temperature was raised 300-500 K, likewise showed a small release of 02 (together with N2). However, 02 release was not observed when that sample was mildly prereduced nor when CeO2 was similarly tested. Preoxidised rhodia dispersed upon Ceria is thus seen to enhance the ease of release of 02 at temperatures < 550 K.

1

)

4.50E-09

OOOO- 0.5% Rh on i CeO2 calcined [ at 1273K

O-O

4.00E-09

4 0 RhCeO2

Ia

!

1 I

CeO2 calcined i at 1273K. I

]

3.50E-09

J ~3.00E-09 t..

~2.50E-09 ,~

~2.ooF~-o9

3.3. Spectroscopic Studies of NzO --Surface Interactions: Previous work has

1.50E-09

1.00E-09

5.00E-lO

0.00E+00 [ 300

i 500

: 700

\ Temperature/-K. Figure 3. Oxygen TPD profiles from the indicated preoxidised materials.

900

demonstrated the value of adsorbing 02 at 77 or 300 K upon Ce02 surfaces activated by prior vacuum outgassing at increasing temperatures (Tv) to serve as an electron-accepting surface probe yielding information, via ESR measurements on the number and type of O~ radicals detected at 77 K concerning relative densities of reduced surface centres, such as electron-vacancy complexes (Ce3"-Vo), [9,10]. However, the application of similar procedures, using N_,O rather than 02 as adsorbed probe, to samples of CeO2 outgassed up to Tv=773

K did not produce any new paramagnetic species, despite FTIR observations confirming appearance of IR features attributable to adsorbed N20 (2234 and 1256 cm" in Fig. 4a) upon contact with N20 at 300 K. Stepwise decreases in magnitude of those IR features were, however, observed in each of a sequence of FTIR spectra taken after separate N20 adsorptions at increasing adsorption temperatures (TO up to 573 K (Fig. 4b-d). From these FTIR observations it could be inferred that increased T a for contact between N20 and vacuum-outgassed CeO 2 resulted in increased fractional decomposition of the N20 introduced. FTIR spectra did not show bands due to peroxide species after N20 adsorption.

688

~

d ~

-

-

~

. . . . . . . . . . . . . . . . . . . . . . . . . . .

__

A

2234

..

.

.

.

.

.

.

_

-

.

_

-

a__., k _ _ _ _ i

II

iiii

ii ii~

2OO0

~

I

[

Ii

iiiiI

II I

I

I _

i

1500

Wavenumbers (cm- 1)

Figure 4: FTIR spectra after N20 adsorption on CeO 2 preoutgassed at T~ = 773 K; T~ = RT (a) and Ta = 473 K (b). Adsorption of N20 at RT on RhO• 2 preoutgassed at T,, = 473 K (c) and T v = 773 K (d). Parallel ESR measurements did not reveal the accompanying formation of O or 0 2 which was to be expected if Ce3*-Vo or other one-electron reducing sites initiated sequences known over other oxides [13] viz N20 --~ N 2 0 --> N 2 + O N__2,O N2 + 02") This leaves N20 --> N 2 + Os2" as the likely alternative pathway for the FTIR-observed increase in decomposition of adsorbed N20 at the higher Ta and also helps to account for dissociation of the first few pulses in profiles similar to Fig. 1A. Information on the fraction of reduced surface sites not oxidised by contacts with N20 was sought using a procedure involving the following steps: (i) vacuum outgassings at Tv = 373-773 K to produce (Ce 3*Vo) centres, such outgassing pretreatment being performed before each N20 adsorption in order to regenerate a desired initial surface condition; (it) N2O adsorption at a selected Ta, followed by desorption at 300 K (verified by FTIR) in order to avoid coverage of adsorption sites by residual N20; (iii) oxygen adsorption at 300 K. Application of this strategy to CeO 2 activated by Tv < 473 K showed that availability of (Ce3*-Vo) or other oneelectron reducing sites for reaction in step iii to yield O 2 was no___]significantly affected by N20 adsorptions even at T~ = 473 K. For CeO 2 more severely outgassed at T~ > 573 K, however, intensity of the 0 2 ESR signal achieved by step iii did progressively decrease with increasing T a for N20 until only a small residual 0 2 signal was detected after T a = 473 K (Fig. 5A,c). Much greater capability to react with N20 at T~ = 473 K and destroy the sites having potential to produce 02 - from 02 was thus established for sites generated on the CeO2 surfaces by T~ > 573, than was the case for reducing sites generated by T, _< 473 K. No new ESR signals were observed upon adsorption of N20 onto RhOx/CeO 2 outgassed

689 at Tv --- 473 K at 300 K, although a decrease was observed in a broad signal with < g > =2.16, previously assigned to Rh 2§ cations in Rhodium oxide clusters [9,10]. Unlike the CeO2 results, N20 adsorption at 300 K onto RhOx/CeO2 after T~ _> 573 K did produce new ESR signals termed 01 (g2 = 2.035, gx 2.016 and g y = 2.011) and RO (g~ = 2.19, g2 = 2.00, g3 = 1.986), accompanied by decreases in broad signals assigned to Rh 2§ cations in rhodium oxide clusters - Figure 5B. ESR signals very similar to these were previously observed to form upon contacting 02 with RhOx/CeO2 > 573 K, and were assigned respectively to Ce 4§ 02-centers stabilised at isolated vacancies and to weak [Rh-OO] 2§ adducts. The contributions by signals 01 and RO changed not only with T~ but also with N20 contact time. For example, during 3 h contact of N20 at 300 K with RhOx/CeO2 after Tv = 573 K intensity of signal RO experienced a large increase although signal 01 remained almost constant (cf. Fig. 5B,plots b and c). This indicated that formation of 02 and related surface complexes is measurable (with ESR), but slow at 300 K. The slower build-up of signal RO may be due to weaker character of its adsorption bond, as evidenced by its disappearance upon outgassing at 300 K while signal 01 remains with substantial amplitude (Figure 5B,d). Neither of these ESR signals was detected in the present study whenever Rh203 powder, after vacuum outgassing at T~ > 573 K, contacted N20 at room temperature. Upon such interaction with N20 the spectra did, however, show a decrease of the broad Rh 2§ signals present in such outgassed rhodia specimens, indicating that electron transfer to the adsorbed molecule did occur, probably with formation of O2-type diamagnetic species as explained above for CeO2.

z.d,, m

_

xl X

O~ o.,

1.5

Z,ltq C

I", o

XIE

I

e,l

o 11

3275

I

|

I

5A

3375

MAGNETIC FIELD (G) Figure 5A. ESR spectra after 02 adsorption at RT on CeO2 preoutgassed at Tv = 573 K. Without N20 preadsorbed (a). With N20 preadsorbed at T N = 373 K (b) and T~ = 473 K.

__J

2.oH 27CX)

32(X~

MAGNETIC FIELD (G) Figure 5B. ESR spectra of RhOx/CeO2. After outgassing at Tv = 573 K (a). Subsequent N20 adsorption at RT during 1/2 h (b) and 3 h (c). The latter outgassed at RT (d).

690 4. CONCLUSIONS The failures of CeO2 alone or Rh203 alone to yield RO or 01-type paramagnetic centers, when subjected to the same experimental strategy as that which did yield such centers over RhOz/Ce02 after Tv >- 573 K, pointed to synergism between the RhOx and CeO2 components of the latter in producing at 300 K the dioxygen species required for 01 and for RO-type signals, possibly through the intermediacy of Rh-O-n-Ce and Rh-O-O-Ce interfacial sites.

Evidence emerges from both the microcatalytic and spectroscopic studies for enhanced dissociative interaction of N20 with preoxidised RhO]CeO2 materials exposed to heliumflush or vacuum-outgassing at T > 573 K. This is consistent with generation of reduced active-sites, such as electron-vacancy complexes, by such treatments.

Prereduced 0.5% RhOx/CeO2 was not passivated under N20 pulses at 623 K in analagous fashion to the passivation of rhodia powder prereduced to metal, thereby again implying synergy between CeO2.x and Rhodia dispersed thereon.

Acknowledgements. Financial support for research at both laboratories under EC contract SCI-CT91-0704 is gratefully acknowledged by all co-authors. The coordinator (JC) also thanks M. Jauch, F. Farrell, J.A. Sullivan and M. O'Neill for valuable assistance.

REFERENCES

1. 2. 3. 4. 5. 6. 7. 8. .

10. 11. 12. 13.

K.C. Taylor, Catal.Rev.Sci.Eng., 35(1993)457. C.H.F. Deden, D.N. Belton and S.J. Schmeig, J.Catal.,155(1995)204. D.N. Belton and S.J. Schmeig, J.Catal.144(1993)9. E.R.S. Winter, J.Catal.34(1974)431; 19(1970)32; 15(1969)144. C. Lamonier, G. Wrobel and J.P. Bennell, J.Mater.Chem.,4(1994)1927. S.E. Golunski, H.A. Hatcher, R.J. Rajaram and T.J. Truex, Appl.Catal.B.Envir.,5(1995) 367. J. El-Fellah, S. Boujara, H. Dexpert, A. Kiennemanov, J. Magerus, O. Tourtet, F. Villain and F. Le Normand, J.Phys.Chem.,98(1994)5522. J. Cunningham, D. Cullinane, F. Farrell, J.P. O'Driscoll and M.A. Morris, J.Mater.Chem.,5(1995)1027. J. Soria, A. Martinez-Arias and J.C. Conesa, J.Chem.Soc.Farad.Yrans.,91(1995). J. Soria, A. Martinez-Arias and J.C. Conesa, Vacuum(1992)437. A.D. Logan and A.K. Datye, Surf.Sci.,245(1991)280. D.G. Castner and G.A. Somorjai, Applic.Surf.Sci.,6(1980)29. J. Cunningham, J.J. Kelly and A.L. Penny, J.Phys.Chem.,74(1970)1992.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

691

The activity of VOx/ZrO2 for the selective catalytic reduction of NO V. Indovinaa, M. Occhiuzzi a, P. Ciambelli b, D. Sannino b, G. Ghiotti c and F. Prinettoc, aChem. Dept., "La Sapienza" University, Roma, P.le A. Moro 5, 00185, Roma, Italy bChem and Food Engineering Dept., Salemo University cChem. Inorg., Chem. Phys. and Materials Dept., Torino University Abstract

Samples VOx/ZrO2, prepared by i) adsorption from aqueous NH4VO3 solutions at pH=l-4, ii) dry impregnation with the same solution, or iii) adsorption from vanadyl acetylacetonate solutions in toluene, were characterized by means of ESR, XPS and IR spectroscopies. In the selective catalytic reduction of NO with NH3 in the presence of 02 (SCR), VOx/ZrO2 catalysts were active and stable. In the NO+NH3 reaction, they had much lower catalytic activity. Their activity depended only on the vanadium content, not on the method used for preparing the catalysts. Catalytic activity (molecules nm -2 s -1) markedly increased with the vanadium concentration up to 3 atoms nm -2 and changed little thereafter, paralleling the increased concentration of specific polyoxovanadates, detected by IR. The surface concentration of NH~ also paralleled the SCR activity. The results suggest a possible role in SCR for NH; ions and adjacent chelating nitrates, also identified by IR.

1. INTRODUCTION The selective catalytic reduction of NO with NH3 in the presence of 02 (SCR) has been extensively studied mainly on VOx supported on TiO2 [1-4]. The commercial catalysts for the SCR of flue gases from stationary sources are V205-TiO2 and V205(-WO3)/TiO2. Many studies have investigated the dispersion, the nuclearity and the oxidation state of vanadium supported on TiO2 [5-14]. All these properties might depend on the support and it was therefore of interest to extend the study to other supports and particularly ZrO2. Szakacs et al. [15] have studied the SCR activity of VOx/ZrO2 catalysts prepared by adsorption on ZrO2 of VO(acetylacetonate)2 from toluene solutions. The SCR mechanism has been investigated by isotopic tracers [16-18], and the surface species on VOx/TiO2 by spectroscopy [19-22]. Takag! et al. [23] proposed a mechanism involving the reaction of adsorbed NO2 with NH~. Although Tops~e et al. [22] evidenced the participation of NH~ species, they are in favour of a mechanism involving reaction with gas phase NO or weakly adsorbed NO. Several investigators have proposed a redox mechanism involving vIv/v v species [1, 3, 6-8, 24] and have pointed out the need for two adjacent V sites [2, 4, 9, 15]. In this paper we report (i) the catalytic activity for SCR of VOx/ZrO2 samples prepared by various methods (adsorption from aqueous metavanadate solutions at different pH values, dry impregnation, and adsorption from VO(acetylacetonate)2 in toluene), (ii) sample characterization (nuclearity, dispersion and oxidation state) by means of XPS, ESR and FTIR and (iii) the nature and reactivity of the surface species observed in the presence of the reactant mixture. Catal .ytic results are here reported in full. Characterization data relevant to the discussion of the catalytic activity will be given, whereas details on the catalysts preparation and

692

characterization will be reported elsewhere.

2. E X P E R I M E N T A L 2.1. Sample preparation

The zirconia support was prepared by hydrolysis of zirconium oxychloride with ammonia, as already described [25]. Before its use as support, the material was calcined in air at 823 K. VOx/ZrO2 samples were prepared by three methods: (i) adsorption from a solution of ammonium metavanadate (AV) at pH values from 1 to 4, adjusted by nitric acid, (ii) dry impregnation with AV solutions and (iii) adsorption from a solution of VO(acetylacetonate)2 in toluene. VOx/ZrO2 catalysts were designated as ZVx(y)pHz, where x gives the analytical vanadium content (weight percent), y specifies the preparation method (a, adsorption, i, impregnation or acac, acetylacetonate) and z the AV solution pH. The V-content was determined by atomic absorption (Varian Spectra AA-30) after the sample had been dissolved in a concentrated (40%) HF solution. BET surface areas (SNm2g -1) were measured by N2 adsorption at 77 K. The SA of ZrO2 was 49 m2g -1. The SA of some ZV samples was determined after the various treatments. All these samples had SA values ranging from 45 to 49 m2g -1, slightly smaller than those of zirconia.

2.2. Procedure and characterization techniques

Specimens were placed in a silica reactor that was equipped with two side tubes for XPS and ESR measurements and connected to a circulation apparatus, described elsewhere [25, 26]. The catalysts, dried at 383 K, were characterized as prepared (a.p.), after heating in dry oxygen at 773 K (s.o.), or after reduction with CO. In some experiments, as specified, samples were exposed to NO, NH3, or various mixtures NO-O2-NH3. Electrons per V atom (eN) were determined from the CO consumed. The average oxidation number of vanadium was calculated as 5 - eN. FT-IR spectra were recorded at RT on a Perkin-Elmer 1760-X spectrophotometer equipped with a cryodetector, at a resolution of 2 cm -1 (number of scans -100). In the 1070-960 cm -1 region, band integration and c u r v e fitting were carried out by "Curve fit, in Spectra Calc." (Galactic Industries Co.). Powdered materials were pelleted in self-supporting discs of 25-50 mg cm -z and 0.1-0.2 mm thick, placed in an IR cell allowing thermal treatments in vacuo or in a controlled atmosphere. The ESR measurements were made at RT or 77 K on a Varian E-9 spectrometer (X-band), equipped with an on-line computer for data analysis. Spin-Hamiltonian parameters (g and A values) were obtained from calculated spectra using the program SIM14 A [26]. The absolute concentration of the paramagnetic species was determined from the integrated area of the spectra. Values of g were determined using as reference the sharp peak at g = 2.0008 of the E'I center (marked with an asterisk in Fig. 3); the center was formed by UV irradiation of the silica dewar used as sample holder. XPS measurements were obtained with a Leybold Heraeus LHS 10 spectrometer operating in FAT mode and interfaced to a 2113 HP computer. Mgk(~ (1253.6 eV) radiation (12 kV and 20 mA) was used. The a.p. sample was pressed onto a golddecorated tantalum plate attached to the sample holder. After the various treatments (s.o., or reduction with CO), the specimen was transferred into the above mentioned XPS tube without exposure to the atmosphere. The spectra were collected by the computer in a sequential manner (figure in parenthesis gives the kinetic energy): Ols (719.0 eV), V2pl/2 (724.5 eV), V2P3/2 (732.0 eV), Zr3d3/2 (1062.5 eV) and Zr3d5/2

693

(1066.0 eV). The binding energy (BE) of Ols (530.0 eV) was taken as reference. The spectrum analysis involved (i) satellite subtraction of Mgk(z components; (ii) inelastic background removal by a linear integral profile; (iii) curve-fitting by a least-squares method, using a mixed Gaussian-Lorentzian function and (iv) determination of the peak area by integration. Satellite subtraction of ec3 and (~4 oxygen components allowed V2pl/2 to be partially resolved. 2.3. Catalytic experiments Catalytic experiments were done in an apparatus consisting of a flow measuring and control system (mass flow controllers, Hitech), fixed-bed flow microreactor, electrically heated and equipped with a temperature programmer-controller (Ascon), two on-line IR analyzers, one for NO (Radas 1G, Hartmann & Braun) and the other for NH3 (Siemens, Ultramat 5E), and an on-line gas chromatograph (Dani 86.10 HT), equipped with a 2 m length column (AIItech CTR) for the analysis of 02, N2 and N20. Typical experiments were conducted in the temperature range 473-723 K, feeding a gas mixture containing 700 ppm NO, 700 ppm NH3 and 3.6 % 02 in helium. The effect of 02 partial pressure was also tested. The flow rate of the reactant gas was 60 L/h (W/F = 5x10 -6 g h cm-3). NH3 was oxidyzed by feeding a gas mixture containing 700 ppm NH3 and 3.6 % 02 in helium. The nitrogen mass balance was better than 90%. Catalytic data were expressed as NO or NH3 conversion percent, or calculated as apparent kinetic constants (k/NO molecules nm -2 s-l), assuming the occurrence of a single reaction (4NO + 4NH3 +02 = 4N2 + 6H20), first order with respect to NO and zero order with respect to NH3.

3. RESULTS AND DISCUSSION 3.1. Vanadium uptake For samples ZV(a)pH1, up to about 2.5 mmol L -1, the vanadium uptake (atoms nrn -2) was proportional to the AV concentration, but as the concentration increased further, the uptake tended to level off (Fig. 1, a). o,I

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a

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r

E o

15

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10

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2

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20

30

NH4VO 3 concentration/mmol L1

40

0

5

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V available/atoms n m "2

Figure 1. Section a V-uptake (atoms nm -2) vs. AV concentration (retool L-l) at (z~) pH=l, (El) pH=2, (O) pH=3 and (o) pH=4. Section b: V-uptake (atoms nm -2) vs. the vanadium available in the AV solution (referred to the ZrO2 surface area) at (z~) pH=l, (El) pH=2, (O) pH=3 and (o) pH=4. For samples ZV(a)pH2-4, up to about 2.5 mmol L-1 the uptake increased linearly

604 and thereafter remained about constant up to 6.5 mmol L-1 . In contrast to ZV(a)pH1, in ZV(a)pH2-4 samples, at AV concentration > 6.5 mmol L -1, the vanadium uptake increased sharply. The increase depended on the precipitation of a vanadium phase on the zirconia support. Accordingly, in these samples, the X-ray analysis showed the presence of a segregated phase, which transformed into V205 after calcination in air at 773 K or s.o. treatment. Because V-uptake is a surface phenomenon, we plotted (Fig. 1, b) the V adsorbed per unit area of zirconia_(atoms nm -z) as a function of V atoms available in the AV solution (V atoms nm-z, referred to the ZrO2 surface area). For ZV(a)pH2+4 samples, up to 2.5 atoms nm -z, all the available V was adsorbed. As the available V increased further, uptake reached an extended plateau, corresponding to about 3 V atoms nm -2. By contrast, for ZV(a)pH1 samples, V-uptake progressively increased throughout the region of the available V. The maximum V-uptake was about 2.4 atoms nm -2. 3.2. XP$ characterization For all samples, both a.p. and s.o., irrespective of the preparation method, the experimental intensity ratios, V2p/Zr3d, increased proportionally to the V-content up to 3 atoms nm -2 (Fig. 2). The ratio approaches those calculated with the "spherical model" proposed recently by Cimino et al. [27] (full line in Fig. 2). For ZV samples with V-content _< 3 atoms nm -2, this finding shows that vanadium species are uniformly spread on the ZrO2 surface. On ZV catalysts with a larger V content (not shown in Fig. 2), the intensity ratios were markedly larger than the corresponding values yielded by the spherical model. The results obtained on samples with V-content > 3 atoms nm -z point therefore to a V surface enrichment.

Figure 2. Intensity ratio, V2p/Zr3d, vs. V-content. Samples: ZV(a)pH1 (o) a.p. and (l) s.o.; ZV(a)pH2 (z~) a.p. and (&) s.o.; ZV(a)pH3 (El) a.p. and (11) s.o.; ZV(a)pH4 (<>) a.p. and (O) s.o.; ZV(i) ( v ) a.p. and ( v ) s.o.; ZV(acac) (+) a.p. and (x) s.o.

Figure 3. ESR spectra at RT of reduced VOx/ZrO2 samples (CO at 623 K). Samples: (a)ZV0.05(a)pH1; (b) ZV0.33(a)pH4; (c) ZV0.58(a)pH1; (d) ZV1.09(a)pH4.

Because of the intensity of V2ppeaks, which is much lower than that of the nearby Ols peak, the vanadium oxidation state could be reliably ascertained only for ZV(a) and ZV(i) specimens with a V loading > 1% (2.5 atom nm-2). The binding energy value of the V2p3/2 component, obtained by curve fitting of the region Ols-V2p,

695

showed Vv only (517.1 eV), in a.p. and s.o. samples, and complete reduction to V Iv (516.6 eV), after reduction with CO at 500 K. On the same sample the redox cycle showed eN=l, corresponding to an average vanadium oxidation state of 4. 3.3. E SR characterization In a. p. and s.o. ZV(a) and ZV(i) samples, no ESR signals were detected. In a.p. ZV(acac), a weak ESR signal of vanadyl species was detected (5% of total V), absent after the s.o. treatment. The spectra of samples reduced with CO at 400 to 623 K consisted of a signal showing a resolved hyperfine structure (Vh), overlapping a broad (,~Hpp = 300 Gauss) and nearly-isotropic band (Vb, giso = 1.97) (Fig. 3). When recorded at 77 K, both Vh and Vb maintained the same shape as at RT, and their intensity as a function of temperature followed the Curie law. The spectroscopic features of Vh (gll=1.913, gL=1.983, and All=204 Gauss, Aj.= 76 Gauss, line-width, dependence on recording temperature, and number of lines), allow the signal to be assigned to mononuclear isolated V TM in a square pyramidal configuration (vanadyl species). The absence of a hyperfine structure and the large value of the line-width of Vb, both features arising from dipolar and exchange interactions among paramagnetic species, suggest its assignment to magnetical!y interacting V IV, formed by the reduction of polyoxoanions anchored to the zlrcon=a surface. The sequence of spectra, referring to ZV samples with increasing V-content, shows that in the low-loading ZV samples up to 0.2 atoms nm -2 isolated mononuclear V Iv species prevailed, whereas with increasing V-loading interacting VIV became prevalent (Fig. 3). Exposure of s.o. samples to NH3-NO at 623 K caused the formation of a weak Vh signal. A subsequent treatment with NO-NH3-O2 mixtures, containing increasing amounts of 02, caused a progressive decrease in Vh and its disappearance with 1-2% 02. Exposure to NO-NH3 of reduced samples (CO at 623 K), therefore containing Vh and Vb, caused a decrease in the ESR detected V IV by 65%. After this treatment Vb species were nearly absent. Exposure of reduced samples to NO-NH3-O2 caused the complete oxidation of V Iv species. 3.4. FTIR characterization in a.p. ZV(a) and ZV(i) samples, broad bands arising from hydrated vanadates were detected in the 800-1100 cm -1 region. Metavanadate-like species (band at 920 cm -1) prevailed on ZV samples with V-content < 1.5 atoms nm -z and decavanadates (bands at 850-880 cm -1 and 960-990 cm -1) in the range 1.5-3 atoms nm -2. A.p. ZV(acac) samples showed bands from CH3 and C=O, suggesting the adsorption of VO(acac)2 as such (spectra not reported). Spectra of s.o. samples differed markedly from those of a.p. samples and were unaffected by a subsequent evacuation up to 673 K (Fig. 4, a). Spectra consisted of a composite envelope of heavily overlapping bands at 980-1070 cm -~, with two weak bands at 874 and 894 cm -1. Irrespective of the preparation method, the integrated area (cm -1) of the composite band at 980-1070 cm "1 was proportional to the V-content up to 3 atoms nm -z. An analysis of spectra by the curve-fitting procedure showed the presence of several V=O modes. The relative intensity of the various peaks contributing to the composite band depended only on the V-content and did not depend on the method used for preparing the catalysts. Samples with V > 3 atoms nm -~" had IR-spectra features similar to those of pure V205 (spectrum 8 in Fig. 4, a). According to the dependence of the intensity of the various peaks on the V-content, we distinguished vanadates with different nuclearities (roughly three types). The first, corresponding to peak 3 and prevailing in most dilute samples (spectra 2 and 3 Fig. 4, b), is a low nuclearity species possibly mononuclear (type-I). Type-II vanadates had an increasing concentration in the vanadium range 0.4-0.8

696

atoms nm "2, peaks 1 2 and 4 (Fig. 4, b and r Type-Ill vanadates had a markedly increasing concentration in the vanadium range 1.5-3 atoms nm -2 and increased little thereafter, peaks 5, 6 and 7 (Fig. 4, b and r

Figure 4. IR spectra of s.o. samples. Section a: ZrO2, curve 1; Z V0.18(a)pH4, curve 2; ZV0.30(acac), curve 3; ZV0.58(a)pH4, curve 4; ZV0.83(a)pH4, curve 5; ZVl.05(i), curve 6; ZV1.21 (a)pH4, curve 7; ZV4.65(a)pH4, curve 8. Section b: curve fitting of the band at 980-1070 cm-1; 990-1000 cm -1, peak 1, 1007-1008 cm -1, peak 2, 1017-1020 cm -1, peak 3, 1025-1029 cm -1, peak 4, 1034-1038 cm -1, peak 5, 1042-1045 cm -1, peak 6 and 1050-1052 cm -1, peak 7. Sample s.o. ZV0.58(a)pH4. Section c: as in section b. Sample s.o. ZV1.21(a)pH4.

Figure 5. IR spectra of s.o. samples after various treatments. Section a: after adsorption of NH3 (1 mbar) at RT: ZrO2 (curve 1), ZV0.58(a)pH4 (curve 2), ZV1.21(a)pH4 (curve 3). Section b: ZV0.58(a)pH4 sample after adsorption of NO+O2 at 623 K (curve 1), after subsequent adsorption of NH3 (1 mbar) at RT (curve 2) and after subsequent heating at 623 K (curve 3). Bands assigned to bridged bidentate nitrates (*) and to chelating nitrates (**). Section c: the same treatments as in section b on s.o. ZV 1.21 (a)pH4. At RT, NH3 adsorbed on Lewis acid sites, Zr Iv and V v. Accordingly, the intensity of bands from NH3 decreased little with the V-content, by 15% at most, as expected on account of the similar Lewis acid strengths of Zr Iv and Vv. The symmetric bending

697

of NH3 was 1157 cm -1 on pure ZrO2, and shifted to higher frequency on ZV. In particular, on most dilute ZV the frequency band was 1195 cm -1 and with increasing V-content it progressively increased up to 1208 cm -1 (Fig. 5, a). NH~ did not form on ZrO2 and ZV samples with V-content < 1.5 atoms nm -2, whereas it did form on more concentrated samples and markedly increased with V-content up to 3 V/atoms nm -2 (Fig. 5, a). At RT, NO adsorption on s.o. ZV samples gave weak bands from N20 and nitrites. The same species formed on pure ZrO2. Adsorption of NO+O2 gave strong bands from bridged bidentate-nitrates (spectra 1 in Fig. 5, b and r On both ZV0.58(a)pH4 and ZV1.21(a)pH4 after evacuation at RT, the subsequent addition of NH3 at RT gave more intense NH~ bands than those on s.o. samples and caused the concomitant transformation of bridged nitrates into chelating nitrates (spectra 2 in Fig. 5, b and r NH~ species were much more intense in ZV1.21(a)pH4 than in ZV0.58(a)pH4. A subsequent heating at 623 K caused the disappearance of chelating nitrates in ZV1.21(a)pH4 (spectrum 3 in Fig. 5, r and their decrease (to 50%) in ZV0.58(a)pH4 (spectrum 3 in Fig. 5, b), while surface-OH formed and H20 and N2 were detected by analysis of the gas phase. Exposure of s.o. samples to NH3-NO at 623 K, caused a slight reduction of V v to V iv, whereas exposure to NO-NH3-O2, did not affect the vanadium oxidation state. Exposure of reduced samples (CO at 623 K) to NH3-NO caused slight oxidation, whereas exposure to NO-NH3-O2 oxidized all V IV. Catalytic activity On all catalysts, the activity for SCR was stable as a function of the time on stream and the ratio NO/NH3 remained very close to unity. 3.5.

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Figure 6. NO conversion (%) vs. temperature. Section a: (n) ZV0.18(a)pH4, (<>) ZV0.34(a)pH4, (m) ZV0.58(a)pH1, ( v ) ZV0.60(a)pH4, (~) ZV0.83(a)pH4, (o) ZV1.17(a)pH4, (zx) ZV4.79(a)pH4. Section b: (<>)ZV0.17(i), (v)ZV0.32(i), (~)ZV0.64(i), (m)ZVl.05(i), (z~)ZV0.30(acac), (o)ZV0.96(acac), (D)ZV1.36(acac)and (x)ZrO2. In samples ZV(a) (Fig. 6, a), ZV(i) and ZV(acac) (Fig. 6, b), NO conversion increased with V-loading at all temperatures. In the whole temperature range, the selectivity to N2 was ve~ high. A small amounts of N20 (< 3%) were detected only above 573 K. Pure zirconla showed some SCR activity at T > 573 K, comparable with

698

that of ZV0.18(a)pH4. The best performance was obtained with the catalyst ZV1.17(a)pH4, containing 2.8 V atoms nm -2, namely a V-content close to that of the adsorption plateau in ZV(a) samples (3 atoms nm-Z). On the sample ZV4.79(a)pH4 (12.6 V atoms nm-2), containing segregated V205, NO conversion reached a maximum at 623 K and decreased thereafter. Higher reaction temperatures resulted in the formation of very large amounts of N20 (>10 %) arising from the oxidation of NH3; NH3 conversion still monotonically increased with the temperature. In all ZV catalysts, the apparent activation energy (Ea/kJ mol "1) was nearly independent of the V-content (42 + 4 kJ mol-1). Therefore, the dependence of catalytic activity on the V-content can be conveniently inspected through k ~ values, the pre-exponential factor of the Arrhenius equation. The finding that k ~ values, irrespective of the method used for preparing the catalysts, stay on the same curve, shows that the SCR activity is mainly controlled by the vanadium content (Fig. 7, a). The marked and non-linear increase of k ~ with the V-content clarifies that the concentration of the active vanadium is not proportional to the V-loading. Namely, only specific configurations are active. To identify the active vanadium configuration, we divided k ~ values by the intensity of i) the composite band at 980-1070 cm -1, ii) peak corresponding to type-I vanadates, iii) peaks of type-II polyvanadates, and iv) peaks of type-Ill polyvanadates. Normalized k ~ values monotonically and markedly increase by a factor 8 to 11, but normalized k ~ values for type-Ill polyoxovanadates remain about constant, well within a factor of two (Fig. 7, b).

A

/

E120 =

[]

IZ 2s

/

90

~ 6o

/

c/ Ii

[3

/

0 z

DO

J

-~ ao

"----

[3

J o o

O-'-'--

O

-

T

0

1

2

V/atoms nm "2

3

0

V/atoms n m "2

Figure 7. The dependence of catalytic activity on the V-content (a) and its correlation with type-Ill polyoxovanadates (b). Section a: 10-3 k ~ vs. V-content on ZrO2 (0), ZV(a) (El), ZV(i) (z~), and ZV(acac) (O). Section b: 10.2 k ~ divided by the total integrated area (cm -'1) of bands in the region 980-1070 cm -1 (El), and 10-~k ~ divided by the sum of components 5, 6 and 7 (0) areas (cm-1). In the absence of 02, NO reduction continued, however at a rate about ten times lower than that in the presence of 02. During 20 h experiments NO conversion remained constant. On 02 addition, the catalytic activity increased with 02 content in the mixture up to about 1000 ppm, and changed little thereafter. We noticed that increasing the 02 concentration caused NO conversion to become lower than that of NH3, probably due to changes in the stoichiometry of the overall reaction (the NO/NH3 ratio passed from 1.5 to 1). Catalytic tests of NH3 oxidation with 02 yielded high selectivity to N2 (66-90%), which decreased with the higher loading catalysts. in special experiments we tested the activity of ZV samples prereduced in an NH3

699

flow (700 ppm in He) for 1 h at 623 K. In the SCR and NO + NH3 reaction the prereduced ZV1.04(a)pH4 sample showed the same activity as the s.o. sample. 4. C O N C L U S I O N S The heating in 02 at 773 K of VOx/ZrO2 stabilizes surface vanadates of various nuclearities. Provided that the V-content in the samples is < 3 atoms nm -2, XPS shows effective spreading of vanadates on the ZrO2 surface. As the V-content increases, the concentration of monomers increases little, whereas that of polyoxovanadates increases markedly, particularly that of type-Ill polyoxovanadates. The relative abundance of the various vanadates depends only on the V-content, not on the method used for catalyst preparation (impregnation, or adsorption from both aqueous and toluene solutions) and pH of the solution used for adsorption (pH =1-4 for AV solutions). This finding strengthens our earlier proposal [28, 29], based on the results obtained on the related MoOx/ZrO2 system, that the ZrO2 surface has a buffer effect. When samples are heated, the buffer effect causes the condensationdecondensation of vanadium species, therefore making the surface composition independent of the nature of the precursor-adsorbate. In agreement with the results from the characterization, the SCR activity of VOx/ZrO2 also depends only on the V-content, not on the method used for catalyst preparation. The marked increase in SCR activity with the V-content shows that only specific vanadium configurations are active. Although we assess the V=O modes associated with these active configurations, IR analysis did not specify the structure of active polyoxoanions. The presence of V Iv on the surface before catalysis is unessential for ca.tai~ic activity. We cannot however rule out an SCR redox mechanism involving Vv-v Iv. ESR and IR results show that the oxidation state of surface vanadium at the reaction temperature is controlled mainly by the composition of the reactant mixture. "7,

U}

E r

150

|

.~ lO0o

E

O z

5o-

0

4

8 12 16 (NH,=)* concentration

Figure 8. The catalyt!c activity for SCR (10 -3 k~ vs. concentration of NH~ (from the IR area of <5asym/Cmq). (n) ZrO2, (o)ZV(a), and (A) ZV(acac) samples. For the abatement of NO with NH3 in the absence of 02, ZV catalysts give low, but stable activity. The activity is strongly enhanced in the presence of 02. This is a common feature to all SCR catalysts, including the ZSM-5 based system [30]. In the presence of 02, our results show the formation of bidentate nitrates and, on NH3 addition, their transformation into chelating nitrates. Our results also show that the

700 concentration of NH~ ions correlates with the SCR activity of VOx/ZrO2 catalysts (Fig. 8), as already observed by Topsoe et al. [22] for VOx/TiO2 catalysts. The concomitant formation at RT of NH~ species and chelating nitrates and their reactivity at higher temperature suggests that they participate in the SCR reaction.

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.

29. 30.

H. Bosch and F. Janssen, CataL Today, 2 (1988) 369. G. T. Went, L. J. Leu, R. R. Rosin and A. T. Bell, J. CataL, 134 (1992) 492. N.Y. Topsoe, J. A. Dumesic and H. Tops~e, J. CataL, 151 (1995) 241. P. Ciambelli, L. Lisi, G. Russo and J. C. Volta, Appl. CataL, in press. M. Rusieka, B. Grzybowska and M. Gasior, AppL CataL, 10 (1984) 101 H. Eckert and I.E. Wachs, J. Phys. Chem., 93 (1989) 6796. G.C. Bond and S. Flamerz Tahir, Appl. CataL, 71 (1991) 1. G. Centi, E. Giamello, D. Pinelli and F. Trifir6, J. Catal., 130 (1991) 220. J.A. Odriozola, J. Soria, G.A. Somorjai, H. Heinemann, J.F. Garcia de la Banda, M. Lopez Granados and J.C. Conesa, J. Phys. Chem., 95 (1991) 240. G.T. Went, L.J. Leu and A.T. Bell, J. CataL, 134 (1992) 479. P. Ciambelli, G. Bagnasco, L. Lisi, M. Turco, G. Chiarello, M. Musci, M. Notaro and D. Robba, Appl. CataL B:E nviromental, 1 (1992) 61. M.M. Kantcheva, K.I. Hadjiivanov and D. Klissurski, J. CataL, 134 (1992) 299. M.M. Kantcheva, A. Davydov and K.I. Hadjiivanov, J. Molec. CataL, 81 (1993) L25 M.M. Kantcheva, V. Bushev and D. Klissurski, J. CataL, 145 (1994) 96. S. Szakacs, G.J. Altena, T. Fransen, J.G. Van Ommen and J.R.H. Ross, CataL Today, 16 (1993) 237. U.S. Ozkan, Y. Cai and M.W. Kumthekar, J. CataL, 149 (1994) 390. B.L. Duffy, H.E. Curry-Hide, N.W. Cant and P.F. Nelson, J. CataL, 149 (1994) 11. B.L. Duffy, H.E. Curry-Hide, N.W. Cant and P.F. Nelson, J. CataL, 154 (1995) 107. G. Ramis, G. Busca, V. Lorenzelli and P. Forzatti, AppL CataL, 64 (1990) 243. N.Y. Topsoe, J. CataL, 128 (1991) 499. L. Lietti, G. Busca, P. Forzatti, G. Ramis and F. Bregani, AppL CataL B: Enviromental, 3 (1993) 13. N.Y. Topsoe, H. Topsoe and J.A. Dumesic, J. CataL, 151 (1995) 226. M. Takagi, T. Kawai, M. Soma, T. Onishi and K. Tamaru, J. CataL, 50 (1977) 441. H. Bosh, F.J.J.G. Janssen, F.M.G. Van den Kerkhof, J. Oldenziel, J.G. Van Ommen and J.R.H. Ross, Appl. CataL, 25 (1986) 239. A. Cimino, D. Cordischi, So De Rossi, G. Ferraris, D. Gazzoli, V. Indovina, G.Minelli, M. Occhiuzzi and M. Valigi, J. CataL, 127 (1991) 744. A. Cimino, D. Cordischi, S. De Rossi, G. Ferraris, D. Gazzoli, V. Indovina, M. Occhiuzzi and M. Valigi, J. CataL, 127 (1991) 761. A. Cimino, D. Gazzoli and M. Valigi, J. Electron Spectrosc. Relat. Phenom., 67 (1994) 429. V. Indovina, A. Cimino, D. Cordischi, S. Della Bella, S. De Rossi, G. Ferraris, D. Gazzoli, M. Occhiuzzi and M. Valigi, Proceedings of the l Oth International Congress on Catalysis; L. Guczi, F. Solymosi and P. Tetenyi, Eds.; Elsevier: Budapest, A (1993) 875. F. Prinetto, G. Cerrato, G. Ghiotti, A. Chiorino, M.C. Campa, D. Gazzoli and V. Indovina, J. Phys. Chem., 99 (1995) 5556 M. Shelef, Chem. Rev., 95 (1995) 209.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

701

Selective Reduction of NOx by P r o p e n e over A u / y - m l 2 0 3 Catalysts M.C. Kung, J.-H. Lee, A. Chu-Kung, and H.H. Kung* IpatieffLaboratory, Center for Catalysis and Surface Science, Northwestern University, Evanston, IL 60208. U.S.A.

The activity of A u / y - m l 2 0 3 catalysts for NO reduction by propene was studied in the presence of 1.5 % H20 and 4.7 % 02. The behavior of the catalysts prepared using different procedure could be classified into three groups. Group A catalysts were active in N2 production below 673 K. They contained aggregates of small grains of Au. Group B catalysts were active above 673 K. They contained large Au crystallites. On these catalysts, the AI203 support contributed substantially to the observed catalytic activity. Group C catalysts were not very active for N2 production, but were the most active for propene combustion. They contained small grains of Au highly dispersed on AI203. It was established that the oxidation of NO to NO2 was not important in NO reduction over catalysts effective for N2 formation.

I. INTRODUCTION A major obstacle in the development of lean-burn, gasoline engines is the lack of a practical catalyst that can reduce NOx emission in the oxidizing atmosphere of the exhaust from such engines. One of the best catalysts to-date for lean NOx reduction is Cu-ZSM-5 [ 1]. However, this catalyst is known to deactivate in use, due to changes in the Cu species and degradation of the zeolite structure [2-4]. On the other hand, the common support, AI203, is generally very stable, and can be further stabilized by modification with other oxides, such as lanthanum oxide. Recently, Maeda, et al. reported the synthesis of an AI203 that is thermally stable, even in the absence of additives [5], that can stabilize Pt and Pd against sintering [6]. Therefore, it is worth exploring Al203-supported catalyst for lean NOx reduction applications. Al203-supported transition metal catalysts have been studied for lean NOx reduction [7-10]. In general, noble metal (Pt, Pd) catalysts are active at low temperatures (<523 K). In particular, Pt is the most active for NO conversion, but the majority of the nitrogen product is the undesirable N20. The first row transition metals show an activity pattern across the period similar to ionexchanged zeolite catalysts, although the activities are lower [ 11 ]. Among these examples, Au/AI203 is one of the more active ones, showing 40% reduction of NO to N2 by propene in the presence of 1.8% H20 and 5% 02 at 573K [ 10]. Since the unusually high CO oxidation activities at low temperatures exhibited by Au/AI203 depend on the details of the preparation method [ 12], it appears possible that the NO reduction activities over these catalysts can be influenced by the preparation conditions also. The objectives of this study are to investigate the possible dependence of lean NOx reduction activities on preparation variables of Au/AI20 3 and to propose a possible explanation for such a phenomenon.

702 2. EXPERIMENTAL Thermally stable A1203 was synthesized as in ref. 5, by hydrolysis of A1 isopropoxide (99.99+% Aldrich Chemicals) dissolved in 2-methylpentane-2,4-diol. The resulting solid was filtered, washed in 2-propano[, and dried in air at 3 73 K. Then, it was calcined in flowing dry air, while the temperature was raised at 1 K/min to 733 K, when 2.4% H20 was introduced to the flowing air. Afterwards, the temperature ramp was continued to 973 K. The sample was kept at 973 K for 2 h in 7% water. The isoelectric point of the resulting Y-Ai203 was pH 8. The BET surface areas were 205 to 235 m2/g, and the average pore size radius was around 8.3 nm. Au/y.,Al203 catalysts were prepared by the deposition-precipitation method similar to the procedure of ref. 13. HAuCI4 (99.999% Aldrich Chemicals) was handled with teflon coated spatula, and dissolved in 50 ml ofdeionized distilled H20 at the desired temperature to make a 6.6 mM Au solution. With vigorous stirring, the Au solution was added to 2.5 g of Al2Os. The initial pH of this mixture was 4, which climbed steadily to 4.4 after 0.5 h. In some cases, the pH was adjusted to the desired value with a Na2COs solution immediately after mixing. The temperature, pH, and the time of addition of Mg citrate were varied for different preparations. If Mg citrate was used, the molar ratio of Mg/Au was 2.5. The solid was suction filtered 2 h alter the initial mixing, resuspended in HzO of the same temperature as during synthesis. The mixture was twice stirred vigorously for l0 min and suction filtered. Then the catalyst was stirred in 323 K H20 for l0 min and suction filtered again. The resulting solid was dried in air at 373 K for 0.5 h, and then calcined at 623 K for 4 h. Before reaction, the catalysts were activated in a reaction mixture of 02, H20 , NO, and propene at 723 K for 2 h and then at 773 K for 1 h. The HCl-treated Al20s was prepared at room temperature by adding 50 ml of aqueous HCI to Al20s such that the Cl- concentration was the same as in the catalyst preparation. After stirring the mixture vigorously for 1 h, Mg citrate was added, the mixture was stirred for another h, and then the solid was filtered, washed, dried, calcined, and activated in the same way as in the preparation of Au/Al203. NOx reduction was conducted in a quartz U-tube microreactor. Either 0.1 g of catalyst and a total flow rate of 136 ml/min, or 0.5 g and a total flow rate of 104 rnl/min were used. The feed was 4.7% 02, 1.5% H20, 0.1% NOx (NO or NO2), and 0.1% CsH6. The hydrocarbon products, CO2, and NzO were analyzed by GC with a 10 tt Porapak Q column, and N 2, 02, and CO were analyzed with a column of 1 ft carbosphere and 2 ft molecular sieve 5A column linked in series. NO oxidation reaction was conducted using a feed of 4.7% 02, 0.1% NO and 1.5% H20. The reaction was monitored with a Beckman model 951 NO/NO x analyzer. The Au contents were analyzed by ICP. UV-vis diffuse reflectance spectroscopy (DRS) was performed on a Varian Cary 1E UV-VIS Spectrometer. X-ray diffraction patterns were collected using a Rigaku dit~actometer with Cu K~ radiation. The surface area and the pore size distribution was obtained using nitrogen adsorption. A term, competitiveness factor (C.F.), was defined as the percent of CsI~ reacted to reduce N~, assuming the following reactions: For NO reduction: 9 NO + CsI-I6 = 9/2 N 2 + 3 CO 2 + 3 H20 For NO2 reduction: 9/2 NO 2 + CsH6 = 9/4 N 2 + 3CO 2 + 3 H20 It was calculated by: For NO reduction: C.F. = 2 x rate o f N 2 f o r m a t i o n / 9 x rate of propene reaction For NO2 reduction: C.F. = 4 x rate o f N 2 formation/9 x rate of propene reaction

703 3. RESULTS 3.1 N O reduction on Au]AI20 3

The NO reduction activities of the catalysts were found to depend on the preparation procedure. In general, the catalysts could be classified into three groups, depending on their competitiveness factors (C.F.) and NO conversion activities. Table 1 shows the properties of representative catalysts from each group. Group A catalysts were active, had high C.F.'s, and showed maximum conversions of NO to N 2 (>50%) at around 643-673 K. The major combustion products was CO2. N2 was the only detectable reduction product of NO. No N20 was observed. Group B catalysts had relatively low activities because, despite the 2 to 4 times higher Au loadings, their activities per unit weight of catalyst were comparable to or lower than those of group A. The maximum NO conversions on these catalysts were attained at temperatures >683 K. These catalysts showed high C.F.'s. Furthermore, even at the maximum NO conversion, combustion of propene to CO2 was incomplete, and substantial amounts of CO were formed. N2 was the only detectable reduction product of NO. Table 1 Activites for NO Reduction by Propene over Au/ml20 3 Catalysts. Group

Catalyst (wt. %

Temp. K

NO conv. %

Au)

to N2 ,,

A

1-7 (0.6 %) .

.

~L

.

636 653 688 .

.

.

.

.

C3H 6 conv.

to N20 .

.

.

.

.

%

to CO2

to CO

a

C.F. %b

Au Au particle plasmon size, nmc peak widthd

.

46 56 45

0 0 0

70 97 100

2 tr 0

6.9 6.2 4.8

3

Narrow

9 11 tr

7.2 6.7 5.5

4.5

Narrow

.

1-11 (0.8 %)

633 673 723

27 59 54

0 0 0

29 78 100

i-5 (1.9 %)

673 698 723

41 52 68

0 0 0

29 50 82

18 22 18

9.8 8.0 7.6

22

Narrow

1-13 (2.4 %)

648 683 703

37 60 57

0 0 0

28 58 87

13 19 13

9.3 8.1 6.0

17

Narrow

B- 10B (1.2 %)

613 623 643

33 36 32

0 4 10

100 100 100

0 0 0

3.8 3.8 3.0

3

Broad

B- 15 (1.0 %)

628 638

0 0

1.9 2.8

3.5

Broad

,,,

B

C

.

.

.

22 28

.

.

0 0

.

.

100 100

.

a) Reaction conditions" 0.5 g catalyst; 0.1% NO, 0.1% propene, 4.7 % 02, 1.5 ~/o H20, balance He; 104 ml/mm total flow. b) Competitveness factor, as defined in text. c) For catalysts after reaction, by XRD line broadening of 311 peak of Au. d) For catalysts after reaction.

704 Group C are catalysts that had high activity but low C.F.'s. The NO conversions reached the maximum values at much lower temperatures of 623-643 K. The propene conversion was 100% at this point, and CO2 was the only detectable oxidation product. In addition, significant amounts of N20 were detected. For some catalysts (such as catalyst B-10B), N20 was observed under the conditions used in Table 1. For others (such as catalyst B-15), it was observed at higher space velocities (see Table 2). Three of the catalysts in Table 1 were also tested using 0.1 g of catalyst (i.e. five times the space velocity as Table 1). The results are shown in Table 2. The trends shown in Table 1 were even more apparent at the higher space velocity. The values of C.F. did not depend much on the space velocity and the decrease in maximum NO conversion was slight relative to the increase in the space velocity, but the temperature where maximum NO conversion occurred increased significantly. This observation is similar to those reported for other lean NOx catalysts [ 1,14]. Figures 1 and 2 show the NO and propene conversions of these three groups of catalysts as a function of temperature. For comparison, the NO conversion of a 3.2 wt.% Cu-ZSM-5 (Si/AI=70) catalyst is also shown, which was obtained at twice the space velocity as the Au catalysts. It can be seen that the NO conversions on Au/AI203 of high C.F.'s were comparable to those on Cu-ZSM-5 under these conditions. 3.2 Characterization of A u / A I 2 0 3 To better understand the differences in their catalytic behavior, the catalysts were characterized by XRD and UV-vis DRS. Unfortunately, except for the peak at 77.6 ~ 20 (311 diffraction), the other Au diffraction peaks overlapped with those of y-Al203. The size of the coherent domains of Au, listed in Table 1, were estimated using the width of this diffraction peak and the Debye-Sherrer equation. They showed that catalysts of both groups A and C had small coherent domains, whereas those of group B had large domains. Figure 3 shows the UV-vis DRS spectra of the three groups of catalysts. In all cases, a prominent Au plasmon peak around 525 nm was observed. This peak was sharper for catalysts of both groups A and B, and broader for catalysts of group C. That is, catalysts of lower C.F.'s had broader peaks. In addition, there were three peaks at 270, 230, and 200 nm These bands were related to the hydroxyls on AI203, since they were observed on pure AI203 also, and their intensities changed with the moisture content of the sample. 3.3 Oxidation of NO to NO2 and Reduction of NO2 on Au/A1203 The activities of the catalysts for the oxidation of NO to NO2 and reduction of N O 2 by propene were measured, and compared with NO reduction under similar conditions. The results are shown in Table 2. For all catalysts, the activities of the catalysts for NO oxidation were low, compared with NO or N O 2 reduction. For catalysts in groups A and B, N O 2 reduction was much faster than NO reduction of NO, and the C.F.'s were much higher. Under the conditions used in this study, the catalytic activities were stable for NO reduction for all catalysts. However, in N O 2 reduction, deactivation was observed. For catalyst 1-7, there was a rapid, reversible deactivation that was more noticeable at lower temperatures. The activity could be restored by removing propene from the feed. Therefore, it was likely due to carbonaceous deposits on the catalyst. In addition, there was slow deactivation. For example, after the experiment in Table 2 and cleaning in a flow of NO/O2/TI20 (0.1%/4.7%/1.5%, balance He) at 500~ the catalyst showed an NO conversion of 33% and propene conversion of 42% at 450~ versus 53 and 99%, respectively, before deactivation. For catalyst 1-5, only slow deactivation was observed.

120 B-1 OB 100 t

.z 0

80

9 0 0

a?

60 40

20 01 550

I I

600

I

650

I

700

I

750

800

850

Temperature (K)

0

5 1

600

650

700

750

800

Temperature (K)

Fig. I : NO conversion in NO reduction over Au/y-Al,O, and CU-ZSM-5 catalysts. Reaction conditions as in Table I

Fig. 2: C,H, conversion in NO reduction over Au/y-Al,O, and Cu-ZSM-5 catalysts. Reaction conditions as in Table 1 .

Fig. 3: UV-vis DRS spectra of Ady-Al,O, after reaction 1.

200

I

300 400 500 600 700 800 Wavelength, nm

706

Table 2" N O Oxidation and N O 2 Reduction Activities o f Au/Al203. NO: reductionb

Group Catalyst NO oxidation to NO:a T,K

NO conv. % T,K

NO reductionc

NO: conv. CsH~ conv. T,K toNs, % %

NO conv. toN:, %

CsH6 conv. %

C.F.

A

1-7

723 748 773

4.8 6.3 9.5

573 673 723

35 44 57

30 28 63

673 698 723 773

30 44 53 42

44 70 99 100

6.8 6.3 5.3 4.2

B

1-5

698 723 748 773

1.9 4.6 6.9 8.9

673 698 723 773

30 37 44 48

30 42 62 100

698 723 773

12 19 34

14 21 56

9.5 9.9 6.9

C

B-15

698 748

4.3 9.8

723

13d

100

1.4

a. Reaction conditions: 0.1 g catalyst; 0.1% NO, 4.7 % 02, 1.5 % HsO , balance He; 136 ml/mm total flow rate. b. Reaction conditions: same as a, except that 0.1% NO2 and 0.1% propene added instead of NO. c. Reaction conditions: same as b, except 0.1% NO instead of NOs. d. In addition, 8% conversion to NsO was detected.

a

Table 3

N O and N O 2 reduction over y -A1203.

Catalyst (wt.)

NOx

Temp. K

NOx conv. to N2,~/o

CHH 6 c o n v .

toCO: HCl-treated (0.1 g)

NO

773 813

10 21

HCl-treated (0.5 g)

NO

700 723 773 813

32 50 69 82

non-HCltreated

NO 2

623 673 723 773

NO

698 723 813

(o.1 g) non-HCltreated (0.5g)

a: Reaction conditions same as Table 2.

C.F.

% to CO 8 23

11.1 8.8

15 24 29 54

13 17 33 39

11.4 12.1 11.1 8.9

34 59 80 90

13 13 13 13

5 12 22 38

46 54 52 40

21 31 57

20 38 85

5 4 15

8.1 8.1 6.6

707 3.4 Catalytic activity of't-Al203 The NO reduction activity of the HCl-treated ml203 was measured in order to evaluate its contribution to the observed activity of Au/AI203. The results are shown in Table 3. As observed before [26], the C.F. on AI203 was high. CO was a very significant product of propene oxidation. Also shown in the table are the results of NO2 reduction over a sample without HCI treatment. AI203 was very active in reducing NO2 to N2, and was able to utilize propene with a high C.F. It was interesting that the NO conversion over a 0.1 g sample was much less than 1/5 of that over 0.5 g. This indicated that a reactive species was formed on AI~O3 that could greatly facilitate NO reduction to N2. A possibility was NO2, since NO2 could be reduced to N 2 much more rapidly on AI203 than NO. In these experiments, the catalytic activities were stable. That is, there was no evidence of deactivation. There was no evidence of oscillation either, unlike the case when the feed did not contain H20, then oscillation was observed between about 723 to 773 K [ 15].

4. DISCUSSION One of the most interesting results of this work is that properly prepared Au/y-ml203 are effective lean NOx reduction catalysts in the presence of 1.5 % H20 and 4.7 % 02. Their activities are stable, and comparable or higher than a Cu-ZSM-5 catalyst under similar reaction conditions. Another interesting result is the observation that the activity depends strongly on the preparation procedure, which must be related to the detailed structure of the catalyst and the nature of the active sites.

4.1 Structure of Au particles in Au/AI203 It has been suggested that the width of the 525 nm Au plasmon peak is broader for smaller Au particles in Au colloids [ 16,17]. Using this correlation, the data in Fig. 3 would suggest that the Au particles are smaller for catalysts in group C than in groups A and B. However, the XRD line broadening results show that the coherent domains of the Au particles of group B is much larger than those of groups A and C. These results could be reconciled by the following model. The Au particles in the group C catalysts are small and well dispersed. This is consistent with both the XRD and UV-vis RDS results, and the fact that the Au loadings of these catalysts are relatively low. The Au particles in group B catalysts are large crystallites, as suggested by both the XRD and UV-vis RDS results. This is also consistent with the fact that these samples have high Au loadings. The Au particles in the group A catalysts are aggregates of small grains of Au of about 3 nm in size. However, these grains are in close contact with each other, such that the aggregates produce an UV-vis RDS spectrum that resembles that of a much larger particle. 4.2 Nature of active site There are at least three types of sites in Au/AI203: those on AI203, those at the Au-AI203 interface, and those on the Au surface. The data in Tables 1, 2 and 3 show that these sites contribute differently on different catalysts. Because of the fact that the NO conversion on AI203 increases faster than linearly with sample weight, the contribution of AI203 to the observed activity of Au/AI203 increases when a larger amount of catalyst is used, and at higher temperatures. When 0.1 g of catalyst is used, the data in Tables 2 and 3 show that AI203 does not contribute much to the observed NO reduction

708 activities of catalysts in groups A or B. When 0.5 g of catalyst is used, from the data in Tables 1 and 3, it can be estimated that up to 1/5 of the observed N 2 production on group A catalysts, and up to 1/3 on group B catalysts could be due to reaction on AI203. The higher contribution of AI203 to the group B catalyst is consistent with the higher concentration of CO produced and higher C.F.'s, which are characteristic for reactions on Al203. In this model of active sites, the difference between catalysts in groups A and B can be attributed primarily to different surface areas of Au. In spite of their lower Au loadings, group A catalysts have higher surface areasof Au because the Au particles exist as aggregates of small grains. In contrast, the Au in group B catalysts exists as large Au crystallites. Catalysts in group C are very active below 673 K. Thus, it can be concluded that AI203 does not contribute much to the activity of these catalysts. The Au in these samples are small grains of about 3 nm in size. But, unlike those in group A catalysts, they are highly dispersed on Al203. Thus, small Au particles, when isolated on the Al203 surface, are active for propene combustion, but not effective for NO reduction. That dispersed small grains of Au on an oxide is active for combustion agrees with literature results showing that highly dispersed Au in TiO2, Fe,zO3, and Co304 are much more active for CO oxidation at low temperatures than poorly dispersed Au [ 18]. They are also active for hydrocarbon combustion [ 10]. Their high activity was explained by the model that the active sites for CO oxidation were at the Au-oxide interface. Although somewhat less active than these catalysts, highly dispersed Au/AI203 are also quite active for CO and H 2 oxidation below 3 73 K [ 19]. Thus, it seems likely that the poor performance of the group C catalysts is related to the high density ofAu-Al203 interface associated with the high dispersion of Au. In contrast, in group A catalysts, the formation of aggregates of small Au grains substantially reduces the density of AuAI~O3 interface, thus minimizing their propene combustion activity. 4.3 Dependence of properties of Au/AI203 on the preparation variables The catalysts used in this study were prepared using different conditions to result in different dispersions of Au particles. Among the variables in the deposition-precipitation method, pH, temperature of the solution, and time of addition of Mg citrate were studied systematically. The values of these variables in the preparation of the catalysts of Table 1 are shown in Table 4. Separately, we have determined that the AuCI4" species is rapidly hydrolyzed to AuCl4.x(OH)x" (x is about 2.8 at 298 K, and 2.2 at 278 K). The isoelectric point of the AI203 used was pH 8. Thus, during the preparation of group C catalysts, the surface of A1203 has a low density of positive charge for the adsorption of AuCl4.x(OH)x'. This could explain the high dispersions and low loadings of Au particles on these catalysts. For groups A and B, the AI203 surface is more positively charged, and can adsorb more Au species. For group A, citrate competes for adsorption onto the surface positively charged sites, and chelates the Au species to minimize the condensation of Au species into a large cluster. Thus, the final Au crystallites are smaller and the total Au loading is also lower than group B catalysts. 4.4 Role of NO~ in NO reduction It has been shown that on Cu-ZSM-5 and C u - Z r O 2 catalysts, reduction of NO and N O 2 in the presence of large excess of 02 p r ~ at about the same rate [20,21 ]. This is because over these catalysts, N O 2 is rapidly reduced to NO (and not NO being rapidly oxidized to NO2) [20,22,23]. On the other hand, on catalysts that do not contain transition metal ions, such as NaZSM-5 [24], Ga~O3 [25], AI203 [26], and H-ZSM-5 [26], NO2 reduction to N2 proceeds much

709 faster than NO reduction. The behavior of the Al203in this study is similar to those reported in the literature. Because of the substantial contribution of Al~O3 to the observed activity of Au/A1203, NO 2 reduction also proceeds faster than NO reduction on these catalysts. On the other hand, the data in Table 2 also show that the rate of NO oxidation to NO 2 is much slower than NO reduction. Thus, NO reduction does not proceed via first the formation of gaseous NO2, similar to the behavior over Cu-containing catalysts [22].

5. C O N C L U S I O N It was observed that properly prepared Au/'f-A1203 catalysts are active catalysts for lean NOx reduction to N2 by propene. The active sites are Au surfaces. The activity of the catalyst is higher if the catalyst contains aggregates of small Au grains than large crystallites. The formation of aggregates minimizes the Au-AI203 interface, which is active for combustion of propene. AI203 also contributes noticeably to the observed activity. Because of this requirement of specific arrangement of Au species, the variables in the preparation have strong influence on the properties of the final catalyst. The production of gaseous NO2 is also not important in NO reduction on these Au/AI203 catalysts, similar to the much studied Cu-containing catalysts.

ACKNOWLEDGEMENT This work was supported by the U.S. DOE, Basic Energy Science, Chemical Sciences, and GM Corporation. The NOx analyzer was donated by Ford Motor Company, and Rika Suzuki and Jason Brooks contributed to the preparation of catalysts.

Table 4: Preparation Variables in the Preparation of Group

a. b. c. d. e.

Au/AI203Catalysts.

Catalyst

pHa

Temp. K

Time of addition of M g c~trate.c

Vigorous stirring

1-7 1-11

4.4 4.4

278 278 d

lh lh

yes yes

1-5 1-13

4.4 4.4

278 278

nonee none

yes yes

7.0b R.T. 0.5 h B-10B no 8.0b R.T. 0.5 h B-15 no pH 0.5 h after mixing Au solution and AI203. After mixing Au solution and A1203, the pH was adjusted to this value by addition of N~CO~. Time after Au solition was mixed with AI203. Temperature was slowly raised to room temperature after addition of Mg citrate. No Mg citrate was added.

710 REFERENCES .

2. 3. .

5. .

.

10.

11. 12. 13.

14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25.

26.

M. Iwamto and H. Hamada, Catal. Today 10(1991) 57. K.C.C. Kharas, H.J. Robota, and D. Liu, Appl. Catal. B, 2 (1994) 225. S. Matsumoto, K. Yokata, H. Doi, M. Kimura, K. Sekizawa, and S. Kasahara, Catal. Today, 22 (1994) 127. J.Y. Yan, G.-D. Lei, W.M.H. Saehtler and H.H. Kung, submitted for publication. K. Maeda, F. Mizukami, S.-i. Newa, M. Toba, M. Watanabe, and K. Masuda, J. Chem. Soc. Farad. Trans., 88 (1) (1992) 97. K. Masuda, T. Sano, F. Mizukami, T. Takezaki and K. Kun, Appl. Catal. Catal. B: Envir., 4 (1994) 187. H. Hamada, Y. Kintaiehi, T. Yoshinari, M. Tabata, M. Sasaki, and T. Ito, Catal. Today, 7 (1993) 111. T. Miyadara and K. Yoshia, Chem Lett. (1993) 1483. A. Obuehi, A. Ohi, M. Nakamure, A. Ogata, K. Mizuno, and H Obuchi, Appl. Catal. B:Envir., 2 (1993) 71. S. Tsubota, A. Ueda, H. Sakurai, T. Kobayashi, and M. Hanata, in Environmental Catalysis, ACS Symposium Series No. 552, J. Armor Ed., American Chemical Society, Washington, D.C., 1994, p. 420. K. A. Bethke, MC. Kung, B. Yang~ M. Shah, D. Alt, C. Li, and H.H. Kung, Catal. Today, 26 (1995) 169. M. Haruta, N. Yamada, T. Kobayashi and S. Iijima, J Catal., 115 (1989) 301. S. Tsubota, M. Haruta, T. Kayashi, A. Ueda and Y. Nakahara in "Preparation of Catalysts V," G. Poneelet, P.A. Jacobs, P. C~ange and B. Dalmon, Eds., Elsevier Sei. Publ., 1991, p. 675. M.C. Kung, K.A. Bethke, and H.H. Kung, Preprint ACS Div. Petrol Chem., 39 (1994) 154. K.A. Bethke, unpublished results. D.G. Duff and A. Baiker, in Preparation of Catalysts VI, Studies in Surface Science and Catalysis, Vol. 9 I, G. Poncelet et al. ed., Elsevier Science Publ., Amsterdam, 1995, p. 505 U. Kreibig, and L. Genzel, Surf. Sci., 156 (1985) 678. M. Haruta, S. Subota, T. Kobayashi, H. Kageyama, M.J. Genet, and B. Delmon, J. Catal. 144, 175 (1993). M. Haruta, S. Tsubota, A. Ueda, and H. Sakurai, in New Aspects of Spillover Effect in Catalysis, T. Inui et al. ed., Elsevier Sci. Publ., Amsterdam, 1993, p. 45. M. Shelef, C.N. Montreuil, and H.W. Jen, Catal. Lett. 26 (1994) 277. J.O. Petunchi and W.K. Hall, Appl.Catal. B, 2 (1993) L 17. K.A. Bethke, C. Li, M.C. Kung, B. Yang and H.H. Kung, Catal. Lett. 31 (1995) 287 Z. Chajar, M. Primet, H. Praliaud, M. Chewier, C. Gauthier and F. Mathis, Catal. Lett. 28, (1994) 33. C. Yokoyama and M. Misono, J. Catal., 150 (1994) 9. M.C. Kung, K.A. Bethke, D. Alt, B. Yang and H. H. Kung, in NOx Reduction, ACS Symposium Series NO. 587, U. Ozkan et al. ed., American Chemical Society, Washingon, D.C., 1995, p. 96. H. Hamada, Y. Kintaichi, M. Sasaki, T. Ito and M. Tabata, Appl.Catal., 70 (1991) LI5.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

711

T h e role of surface-generated gas-phase methyl r a d i c a l s in the reduction of N O by C H 4 over a Sr/La203 catalyst Shuibo Xie, Todd H. Ballinger, Michael P. Rosynek and Jack H. Lunsford Department of Chemistry, Texas A&M University College Station, Texas 77843 USA Methyl radicals have been detected in the gas phase over a Sr/La203 catalyst during the reaction of CH4 with NO, provided 02 is present in the system. In the absence of 02 the concentration of CH3" radicals decreases almost to the background level. The results indicate that the enhanced effect of 02 on the reduction of NO by CH4 may be due to surface-generated gas-phase CH3" radicals, but in the absence of 02 another reaction pathway may be dominant. Evidence has been found for the presence of CH3NO, a likely intermediate in the radical reaction, at temperatures up to 800 ~C.

1. INTRODUCTION The removal of NO from exhausts of lean-bum engines and stationary power generators continues to be a challenging problem in heterogeneous catalysis. Vannice and co-workers [1-4] have recently demonstrated that many of the strongly basic oxide catalysts which are effective for the oxidative coupling of methane also are active for the reduction of nitric oxide by methane. Methyl radicals are known to be intermediates in the oxidative coupling reaction [5-7], and, by analogy, it has been suggested that these mdic~s may be involved in the reduction of nitric oxide [1-4]. Indirect evidence for the reaction of CH3" radicals with NO comes from the observation that the formation of C2I-I~, a coupling product, is strongly inhibited when NO is introduced during the reaction of CH4 and 02 over a catalyst. Presumably, the reaction intermediate CH3NO is favored over the coupling product. In homogeneous gas phase studies, carried out at room temperature, it has been shown that nitrosomethane reacts further with nitric oxide to form nitrous oxide and methoxy radicals [8]; whereas, at much higher temperatures (> 1,000 ~ hydrogen cyanide is the principal product [9]. The direct measurement of CH3" radicals, using a matrix isolation electron spin resonance system (MIF.SR), has previously demonstrated that surface-generated CH3" radicals are produced during the oxidative coupling of CH4 and emanate into the gas phase

712 where most of the coupling occurs to yield ~ [5]. The oxidative coupling reactions are typically carded out at temperatures between 650 *C and 800 ~ Vannice and coworkers have observed, however, that the reaction of CH4 with NO occurs at temperatures as low as 500 ~ and they suggest that the mechanism may involve a reaction between surface methyl ~ e s and adsort~ NO. The location of the reaction, i.e., gas phase vs. surface, is of considerable interest in establishing the details of the mechanism. The present study was undertaken to determine the role of gas phase CH3" radicals, formed on the surface of a Sr/La203 catalyst, during the reduction of NO by CH~. This material was selected because it is one of the most active catalysts in the reduction reaction, and, under suitable conditions, it is very effective in producing CH3" radicals [10].

2. EXPERIMENTAL A variable ionization energy mass spectrometer (VIEMS), similar to that described by Stair and co-workers [11], was used to detect radical intermediates and stable products. At a nominal electron-impact energy of 16-19 eV, it is possible to selectively ionize CH3" radicals in the presence of a very large excess of CH4. In addition to CH3" radicals (9.8), it also is possible to ionize C2H4 (10.5), C~I-I~(11.5) and NO (9.2) at 16-19 eV, but not CI-L (12.8) CO (14.1), CO2 (13.9), N2 (17) or N20 (12.9). Molecular oxygen (12.1) can be detected at 19 eV, but not at 16 eV. The numbers given in parentheses are the appearance potentials in eV for the respective molecules. Electron-impact energies considerably greater than the appearance potentials were used in these experiments to improve the sensitivity. As will be shown in the subsequent section, it was possible to discriminate between molecules having different appearance potentials, even when these larger electron impact energies (16-19 eV) were employed. For this study a UTI Model 10(K~ quadrupole mass spectrometer was modified to allow the electron-impact energy to be varied. The methodology used to detect CH3" radicals formed during the partial oxidation of CI-L is similar to that employed by Gutman and co-workers [7]; however, in their system, selective ionization was carried out by photons instead of electrons.

Figure 1. Reactor used in VIEMS system

The catalytic reactions to produce CH3" radicals and other products were carried out in the fusedquartz reactor depicted in Figure 1. Approximately 200 mg of catalyst was placed near the exit of the reactor. The reactor was connected via a skimmer to

713 the mass ~ t r o m e t e r . The pressure in the reactor was mass spectrometer was c a . 3 X 10"6 Torr.

ca.

500 mTorr, while that in the

In a separate set of experiments designed to follow the gas phase reactions of CH 3" radicals with NO, CH3" radicals were generated by the thermal decomposition of a z o m ~ e , CH~I=NCH3, at 980 *C. The CH3" radicals were subsequently allowed to react with themselves and with NO in a Knudsen cell that has been described previously [12]. Analysis of intermediates and products was again done by mass spectrometry, using the VIEMS. Calibration of the mass spectrometer with r e ~ t to CH3" radicals was carded out by introducing the products of azomethane decomposition directly into the high vacuum region of the instrument. Some reaction results also were obtained in a conventional single-pass flow reactor ~ t i n g at atmospheric pressure. The reactor consisted of a 4 mm i.d. fused-quartz tube. The reagent gases were mixtures of 9.9% 02 in He, 4.1% NO in He and 1.1% CH4 in He, all of which were obtained from Matheson. The products were analyzed by gas chromatography. The 1 wt % Sr/La203, provided by Amoco, has been used in previous studies [10]. In both the low- and high-pressure experiments the catalyst was pretreated in flowing 02 for 1 h at 800 *C. The surface area of the Sr/La~O3 after pretreatment was 5.4 m2/g.

3. R E S U L ~ AND DISCUSSION 3.1. Detection of CH3" radicals during the partial oxidation of a large excess of CH4 The effect of changing the electron-impact energy on the observed mass spectrum, following the reaction of CH4 and 02 over the Sr/La~O3 catalyst at 800 *C is shown in Figure 2. The reagents consisted of a CHdO2 ratio of I0; therefore, 02 was the limiting reagent, and after reaction the effluent consisted mainly of CH4. At 70 eV the typical fragmentation pattern of CI-L is evident, with w.aks at 12-16 amu. Most significant is the relatively large peak due to CH4 at 15 amu, which would mask the contribution from a small concentration of CH3" radicals. By contra~_.st, at an ionization voltage of 16 eV the dominant peak, due to CH3" radicals, is at 15 amu, and there is very little contribution due to CI-L, either at 16 amu or at 15 amu, even though the CH4/CH3" concentration ratio was approximately 100. Because of the high level of amplification, there is always a small background at 15 amu, and in reporting subsequent results this background has been subtracted out. In addition to the peak at 15 ainu in spectrum c of Figure 2A, there is also a major peak at 28 ainu, with minor peaks at 29 and 30 amu. These peaks result from the ionization of C2H4 (28 amu) and the ionization and fragmentation of C2H6 (28-30 ainu). Based on calibration results using pure components, the C2H4/C2H6 ratio was c a . 0.17.

714 Under the conditions of this experiment, the small amount of C2H 4 probably results from the catalytic dehydrogenation of CaI-I~ at 750 ~ rather than from the reaction 2CH3" -" Calls" + H-, which competes with the coupling of CH3" radicals in the gas phase

A

ai

100

.~

so

.

.

.

.

.

.

I 'v

.

.

.

.

.

.

.

.

I

9

.

o

o.

E

_

il

b ~x10 a

10

20

Temperature,

• 10"6 __.~, x lO -'s. _..

, . . . . . , , . I , . . ,

Number,

20

, . , , I

,,

..

30

ainu

apparent activation energies for CH3" radical and C2 formation, obtained in the temperature range 520 ~ to 670 ~ were 13 kcal/mol and 28 kcal/mol, respectively. The effect of temperature on the formation of CH3" radicals agrees nicely with the results previously obtained over this catalyst using the MIESR system [ 10].

.

!

.

__

i

600

a

Figure 2. Mass spectra obtained at different electron impact energies: (a) background ; (b) CH, and 02 at 25 ~ (e) after passing CH4 and O2 over a Sr/LaaO3 catalyst at 750 ~ The CH4/O2ratio was I0:1.

o

500

~

30 10

Mass

1

400

t

-~ . . . .

A__.,~xlO "1I

9. , , , , , , . I , , , , , , , , , I , , ,

1

-

70 eV

"o

Methyl radicals and Ca products were detected over the Sr/La203 catalyst at temperatures as low as 520 ~ as shown in Figure 3. The concentrations of these two species increased up to 750 ~ and then decreased at higher temperatures (i) because the 02 was essentially depleted, and (ii) secondary reactions became more significant. The .

B

J

[7].

.

16eV

700

....

800

~

Figure 3. Mass spectra obtained at (A) 16 eV and (B) 70 eV a t ~ passing CH 4 and 02 in a I 0:1 ratio over Sr/La203 at different temperatures: O, mass 15; A, mass 28; II, mass 16; V, mass 27; e , mass 28; t , mass 30; 6, mass 32; D, mass 44.

3.2. Effects of adding NO The effect of adding increasing amounts of NO to the CH4 and 02 reagents is depicted in Figure 4. Here, the actual pressures of CH3" and C2 products are given; however, it should be pointed out that because of uncertainties in calibration, the

715 results are semiquantitative. With no NO present, the CH3" radical concentration was ca. 3 X 1013 molecules/cm 3. Clearly, the presence of NO caused a decrease in the concentration of CH3" radicals and, consequently, in the formation of C2 products. The nonlinear decrease in CH3" radicals suggests that at low concentrations the NO may interfere with the production of the radicals at the surface, and at all concentrations, reactions with NO in the gas phase may destroy CH3" radices. The decrease in C2 products is consistent with the results reported at higher pressures [4]. 3.0

.

.

.

.

14

.

12

2.5 =3

2.0

910

CH4+O2+Ar (10:1:10) - r n CH4+O2+NO (10:1:10) E l CH4+Ar+NO (10:1:10) -

8

1.$ C

1.0

c

6

c --

4

x 5000

m

L 0.0

.....

0.0

0~

0.4

0.6

0.8

1.0

NOICH 4 Ratio

Figure 4. Effect of adding NO to a CH4 + 02 reaction mixture (CH4/O2 = 10:1) over a Sr/I.,a203 catalyst at 735 ~ and a total pressure of 500 mTorr, O, CH3", A, C2I-I6. In the absence of O2: @, CH3"; ~k, C2H6. Mass spectra were obtained at 16 eV electron impact energy.

15

26

27

28

Mass Number, amu

Figme 5. Variation in concentration of products over Sr/La203 as a function of the reagent gas. comtx~itiort Results were obtained at 16 eV for CH3" radicals and at 70 eV for stable products: P(total) = 490 reTort, T = 735 ~

It is significant that in the absence of 02 (solid points in Figure 4) almost no radicals were formed; the amount reported is close to the detection limit of the instrument. In one sense, this observation provides an explanation for the positive effect that 02 has on the rate of reaction between NO and CH4 [3,4]; i.e., 02 enhances CH3" radical formation. However, the results also indicate that NO itself is not very effective in generating active sites which are responsible for CH3" radical production. This means that the reaction of NO with CH4, in the absence of added 02, may occur via a nonradical pathway. In a similar set of experiments, the gas phase reactant concentrations were varied, and the relative concentrations of the important species were determineA. The stable products were detected at 70 eV, which allowed one to determine the contribution of N2, if present,

716 to the peak at 28 ainu. As shown in Figure 5, with only CH4 and 02 as reagents, significant amounts of CH3" radicals and C2 products were observed. But when the Ar diluent was replaced by NO, there again was a decrease in CH3" radical and C2 products. (The latter is based on the peaks at 26 and 27 ainu). The increase in the mass 28 peaks in this experiment confirms that a relatively large amount of N2 was produced. Finally, the NO concentration was held constant and the 02 was replaced by Ar. Again N2 was formed, although in smaller amounts, but the concentrations of CH3" radicals and C2 products decreased further. It seems unlikely that the NO simply reacted with most of the CH3" radicals in this case, because the same m o u n t of NO was present in both of the latter two experiments. Rather, NO a ~ s to be much less effective in generating the active sites for the formation of CH3" radicals on the surface. Previous studies on the generation of CH3" radicals over a Li/MgO catalyst demonstrated that N20 was much less effective than 02 at comparable concentrations. At much greater concentrations, however, N20 could be used as an oxidant, and, in fact, at comparable levels of CH4 conversion, it resulted in greater C2 selectivity [13]. 3.3. Gas phase reactions of CH3" radicals with NO Although the reaction of C H 3" radicals with NO in the gas phase has been extensively studied by indirect methods, we are aware of only one report in which the intermediate CH3NO was directly determined, and in this case the reaction was carried out at room temperature [14]. In order to gain insight into the thermal stability of CH3NO, a cursory study was carried out in a Knudsen cell over the temperature range from 25 ~ to 800 ~ Unfortunately, in the Knudsen cell we were restricted to pressures of a few mTorr. Nevertheless, as shown in Figure 6, there 7 1 1 9 is evidence for the presence of CH3NO (45 6 ainu) at temperatures up to 800 ~C. :i

s

E "r-I

0

2oo

4o0

6oo

800

T e m p e r a t u r e , *C

Figure 6. Formation of CH3NO during the gas phase reaction of CH3" radicals with NO in a Knudsen cell: O, mass 15; A, mass 28; El, mass 45. Results were obtained at P(CH3") = 0.7 mTorr, P(NO) = 1.4 mTorr and an electron impact energy of 19 eV.

As the temperature increased, the concentration of CH 3" radicals increased up to about 700 ~ and the concentration of C2 products correspondingly decreased. This phenomenon was observed even in the absence of NO, and it is partially explained by the fact that the residence time decreased from 0.055 s at 90 ~C to 0.036 s at 600 oC. Of greater significance is the fact that the activation energy for CH3" radical coupling determined from this experiment was - 90 cal/mol. Both positive and negative activation energies have been reported for this reaction in the literature [ 15].

717 The concentration of CH3NO decreased monotonically with temperature, but even at 800 *C a small concentration could be detected. The decrease could be due to unimolecular reactions, to biomolecular reactions with either CH3" or NO, or to a reverse reaction, CH3NO --" CH3" + NO. It is reasonable to assume that similar gas phase reactions would occur at much higher pressures; however, we have not detected CH3NO during our catalytic experiments over the Sr/La203 catalyst. Perhaps at much larger concentrations of NO the reaction CH3NO + NO -. N20 + CH30" rapidly removes the nitrosomethane. 3.4 Catalytic reactions at atmospheric pressure In order to obtain reasonable concentrations of CH 3" radicals in the low pressure studies, reactions were carried out at temperatures > 700 *C; whereas, the conventional catalytic studies of Vannice and co-workers have been carried out at temperatures < 700 ~ Therefore, for comparison with our low pressure results, the effect of temperature on the rate of N 2 formation at a total pressure of 760 Torr has been investigated up to 875 ~ for three gas compositions: (i) 1% NO, (ii) 1% NO + 0.25 % CH4 and (iii) 1% NO + 0.25 % CI-I~ + 0.5 % 02. The results are summarized in Figure 7. At 700 ~ the trends with respect to the effect of gas composition are qualitatively the same as those reported by Vannice; i.e., the addition of CH4 enhanced the rate of N2 formation, and the addition of 02 enhanced the rate even more. At 700 ~ the effect of adding CH4 was to increase the rate two-fold, while the effect of adding CH4 + 02 was to increase the rate five-fold. Unexpectedly, the rate of N 2 formation in the presence of CI-I4 and 02 went through a maximum at 700 *C. Furthermore, in the temperature range 700 *C - 800 *C, the conversion of CH4 approached nearly 100%. Although it is not possible to measure directly the amount of O2 under reaction conditions, one may presume that some remains even at temperatures of ca. 700 ~C. Nevertheless, at higher temperatures as 02 is consumed its positive effect diminishes. Moreover, as shown in Figure 3, the production of CH3" radicals reached a maximum at 750~ The decrease in oxygen concentration and the maximum in CH3" radicals formation may cause the maximum in the rate of N2 formation that is described in figure 7. In addition to the maximum at 700 ~ the rate of N2 production also went through

~,

0.08

,

,

,

,

500

600

700

800

E eo

"~ 0.06 E

o.o4 c o ~ 0.02 o u_ 0.1~ 400

90

Teml~mture,~ Figure 7. Rate of N 2 formation over Sr/La203 in a conventional catalytic reactor with the following reagent gases:, O 1% NO;A, 1% NO + 0.25% CH4; ll, 1% NO + 0.25% CH 4 + 0.5% 02. The total pressure was 760 Torr.

718 a minimum at 800 ~ Such a minimum is expected since, at sufficiently high temperatures, the direct decomposition of NO would become significant. In fact, by 850 ~C, the direct decomposition of NO was about twice as fast as when CH4 and 02 were added as reagents. This apparent negative effect of CH4 and 02 probably is a result of CO2, which is known to be a strong poison for these basic oxide catalysts during oxidative coupling reactions [16,17]. It is relevant to the low pressure studies, particularly those of Figure 5, that at 735~ CI-L + 02 had a positive effect on the rate of N2 formation and on the rate of CH3" radical production, relative to the case when only CH4 + NO were the reagents. These results are consistent with the view that CH3" radicals are intermediates in the reaction of CH 4 with NO when 02 is present, but in the absence of 02, the reduction of NO by CH4 over Sr/La203 may occur via another pathway that does not involve CH3" radicals.

ACKNOWLEDGMENT This research was supported by a grant from the U. S. Department of Energy, Office of Basic Energy Sciences. The authors wish to thank Professor M. A. Vannice for providing us with preprints of his unpublished papers and for helpful discussions.

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