Reaction Kinetics and the Development of Catalytic Processes (Studies in Surface Science and Catalysis)

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Studies in Surface Science and Catalysis 122 REACTION KINETICS AND THE DEVELOPMENT OF CATALYTIC PROCESSES

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Studies in Surface Science and Catalysis Advisory Editors: B. Delmon and J.T. Yates

Vol. 122

REACTION KINETICS AND THE DEVELOPMENT OF CATALYTIC PROCESSES Proceedings ofthe International Symposium, Brugge, Belgium, April 19-21, 1999 Editors

G.E Froment TexasA & M University, Department of Chemical Engineering, CollegeStation, TX 77843, USA K.C. Waugh UMIST, Department of Chemistry, Manchester, UK

0

1999

ELSEVIER A m s t e r d a m - - Lausanne-- New Y o r k - - O x f o r d - - S h a n n o n - - Singapore m Tokyo

ELSEVIER SCIENCE B.V. Sara Burgerhartstraat 25 P.O. Box 211, 1000 AE Amsterdam, The Netherlands 91999 Elsevier Science B.V. All rights reserved. This work is protected under copyright by Elsevier Science B.V., and the following terms and conditions apply to its use: Photocopying Single photocopies of single chapters may be made for personal use as allowed by national copyright laws. Permission of the publisher and payment of a fee is required for all other photocopying, including multiple or systematic copying, copying for advertising or promotional purposes, resale, and all forms of document delivery. Special rates are available for educational institutions that wish to make photocopies for non-profit educational classroom use. Permissions may be sought directly from Elsevier Science Rights & Permissions Department, PO Box 800, Oxford OX5 1DX, UK; phone: (+44) 1865 843830, fax: (+44) 1865 853333, e-mail: [email protected]. You may also contact Rights & Permissions directly through Elsevier's home page (http://www.elsevier.nl), selecting first 'Customer Support', then 'General Information', then Permissions Query Form'. In the USA, users may clear permissions and make payments through the Copyright Clearance Center, Inc., 222 Rosewood Drive, Danvers, MA 01923, USA; phone: (978) 7508400, fax (978) 7504744, and in the UK through the Copyright Licensing Agency Rapid Clearance Service (CLARCS), 90 Tottenham Court Road, London W l P 0LP, UK; phone (+44) 171 436 5931; fax: (+44) 171 436 3986. Other countries may have a local reprographic rights agency for payments. Derivative works Tables of contents maybe reproduced for internal circulation within their institutions but permission of Elsevier Science is required for external resale or distribution of such material. Permission of the publisher is required for all other derivative works, including compilations and translations. Electronic Storage or Usage Permission of the publisher is required to store or use electronically any material contained in this work, including any chapter or part of a chapter. Contact the publisher at the address indicated. Except as outlined above, no part of this work may be reproduced, stored in a retrieval system or transmitted in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, without prior written permission of the publisher. Address permissions requests to: Elsevier Science Rights & Permissions Department, at the mail, fax and e-mail addresses noted above. Notice No responsebility 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. Because of rapid advances in the medical sciences, in particular, independent verification of diagnoses and drug dosages should be made.

First edition 1999

Library of Congress Cataloging in Publication Data A catalog record from the Library of Congress has been applied for.

ISBN: 0-444-50073-1

~The paper used in this publication meets the requirements of ANSI/NISO Z39.48-1992 (Permanence of Paper). Printed in The Netherlands.

CONTENTS Preface KEYNOTES

Catalysis from first principles J.K. Norskov Complexities and dynamics of the enantioselective active site in heterogeneous catalysis R. Raval, C.J. Baddeley, S. Haq, S. Louafi, P. Murray, C. Muryn, M. Ortega Lorenzo and J. Williams

11

Intrinsic activities and pore diffusion effect in hydrocarbon cracking in steamed Y zeolite H.H. Kung

23

Internal reforming in solid oxide fuel cells R.M. Ormerod

35

Microchemical reactors for heterogeneously catalyzed reactions D. H6nicke

47

TRANSIENT STUDIES

Novel frequency response techniques for the study of kinetics in heterogeneous catalysis M. Cavers, J.M. Davidson, I.R. Harkness, G.S. McDougall and L. V. C. Rees

65

Reactivity of novel metal substituted heteropolyacid catalysts using Steady-state and transient response kinetics H.T. Randall, P.L. Mills and K. Kourtakis

73

Use of isotopic transient methods for mechanistic analysis of ethylene hydroformylation over 4 wt% Rh/SiO2catalyst S.S.C. Chuang, S.A. Hedrick and M.A. Brundage

83

CO oxidation over a supported Pt catalyst:transient kinetics using temporal analysis of products (TAP) A.H.J. Colaris, J.H.B. Hoebink and J.C. Schouten

93

Transient kinetics of ethylene and carbon monoxide oxidation for automotive exhaust gas catalysis : experiments and modeling J.M.A. Harmsen, J.H.B. Hoebink and J.C. Schouten

I01

vi

Dynamics of NO adsorption and transformation over supported Pt catalysts for the treatment of lean burn engine emissions G.E. Arena, A. Bianchini, G. Centi, F. Vazzana a n d P. Vitali

109

Effect of catalyst deactivation on the dynamics of cyclic reactive processes D.O. Bofio, N.S. Schbib and J.E. Gafica

117

Propyne hydrogenation over a silica-supported platinum catalyst studied transient conditions D.R. Kennedy, B. Cullen, D. Lennon, G. Webb, P.R. Dennison and S.D. Jackson

125

A TAP reactor investigation of the oxidative dehydrogenation of propane over a V/MgO catalyst: experiment and modeling Y. Schuurman, T. D~camp, J.C. Jalibert and C. Mirodatos

133

Transient behavior of an industrial acetylene converter N.S. Schbib, A.F. Errazu and J.A. Porras

141

Modeling of alkane dehydrogenation under transient and steady state 149 conditions over a chromia catalyst using isotopic labelling S.D. Jackson, J. Grenfell, I.M. Matheson and G. Webb The development of a model capable of predicting diesel lean NOx catalyst performance under transient conditions G.P. Ansell, P.S. Bennett, J.P. Cox, J.M. Evans, J.C. Frost, P.G. Gray, A.-M. Jones, M. Litorell, R.R. Rajaram, G. Smedler and A.P. Walker

157

DYNAMICS OF SURFACES

Monte Carlo and lattice-gas studies of the kinetics of hydrocarbon hydrogenation reactions A.S. M c L e o d a n d L.F. Gladden

167

Time distribution of adsorption energies, local monolayer capacities, local isotherms and energy distribution functions on catalytic surfaces Ch. Abal-zoglou and N.A. Katsanos, A. Kalanlzopoulos and F. Roubani-Kalantzopoulou

175

Molecular dynamical studies of the mobility of benzene and water on silica surfaces : correlation with the influence of surface chemistry and morphology S.P. Rigby a n d L.F. Gladden

183

Experimental study of reaction instability and oxcillatory behavior during CO oxidation over Pd supported on glass fiber catalysts I. Yuranov, L. Kiwi-Minsker, V. Barelko and A. Renken

191

Enhancement of selective conversion in spatially patterned reactors A.S. C6t~, W.N. Delgass and D. Ramkrishna

199

vii NOVEL REACTORS

A novel laboratory reactor for gas-phase transient kinetics based on time-of-flight and quadrupole mass spectroscopy H.T. Randall, P.L. Mills a n d J.S. McCracken

209

Development of new photo-catalytic methods and reactors for waste water treatment A. Starosud, A. Bhargava, C.H. Langford and A. Kantzas

219

An integrated dehydrogenation-hydrogenation membrane reactor for 229 the simultaneous production of styrene and cyclohexane T.M. Moustafa, I. Ashour a n d S.S. Elnashaie A simple and flexible micro reactor for investigations of heterogeneous 237 catalytic gas phase reactions G. Veser, G. Friedrich, M. Freygang and R. Zengerle The use of a catalytic wall reactor for studying highly exothermic reactions B. Amon, E. Klemm a n d G. Emig

247

A novel reactor to activate chromium-catalysts H. Sch6nfelder, M. K~mmerer a n d W. de Lange

255

KINETIC MODELING

Kinetics and reactor simulation for polyethoxylation and polypropoxylation reactions E. Santacesafia, M. Di Sefio a n d P. lengo

267

On the use of response reactions in the kinetic modeling of complex heterogeneous catalytic reactions I. Fishtik a n d R. Datta

275

Kinetic effects of chemical modifications of PMo12catalysts for the selective oxidation of isobutane M. Sultan, S. Paul a n d D. Vanhove

283

Kinetic based deactivation modeling of an isothermal propane dehydrogenation reactor E.H. Sfitt, S.D. Jackson a n d F. King

291

Development of kinetic models for reactions of light hydrocarbons over ZSM-5 catalysts. Experimental studies and kinetic modeling of ethene transformation and deactivation of HZSM-5 catalyst D.B. Lukyanov

299

Two-dimensional reactor modeling of the pure dehydrogenation of methanol to formaldehyde S. Ruf, S. Schunk, G. Emig, Th. Weber, S. Braun, G. Brenner a n d F. Durst

307

viii

Kinetic modeling of enzymatic chiral resolution of ~)2-methyl butyric acid R. Garcfa, M. MarEnez, T. Garcfa and J. Aracil

317

Regioselective synthesis of monoglycerides. Kinetic modeling T. Garcfa, M. Marffnez, D. Garcfa and J. Aracil

325

Kinetic modeling of paraffins hydrocracking based upon elementary steps and the single event concept G. Martens a n d G.F. Froment

333

A comparative kinetic study of CH4oxidation by NiO/AI2Q, Pt/AI203 and NiO-Pt/AI2Q catalysts T.-N. Angelidis a n d V. Tzitzios

341

CATALYSIS IN PROCESSES

Kinetic and mass-transfer effects in the hydrogenation of xylose to xylitol 351 J.-P. Mikkola, T. Salmi a n d R. Sj6holm Methylation of biphenyl over zeolite H-ZSM-5 in gas phase with 359 methanol in presence of water : effect of the catalyst impregnation by tetraethyl orthosilicate S. Dubuis, R. Doepper and A. Renken Kinetics and mechanism studies of the catalytic dehydrogenation of isobutane on Platinum-Indium catalyst D. Casanave, K. Fiaty, J.A. Dalmon and M. Forissier

367

Synthesis of ethylbenzene from diethylbenzenes in the presence of benzene using triflic acid as catalyst M.C. AI-Kinany a n d S.H. AI-Khowaiter

375

Elementary steps of reaction pathway in the pilot plant photomineralisation of s-triazines on to photocatalytic membranes immobilising titanium dioxide and promoting photocatalysts A. Moroni, I.R. Bellobono and B.M. Gawlik

385

Methane oxidation over supported nickel catalysts A.M. Diskin, R.H. Cunningham and R.M. Ormerod

393

POSTERS

Novel sensor for studying the transient behaviour of an iron antimonate 403 partial oxidation catalyst D. Barth, M. Sahibzada and i.S. Metcalfe Chemical equilibria in direct synthesis of dimethyl ether M. Grzesik a n d J. Skrzypek

407

ix

Thermodynamics and kinetics of the synthesis of higher aliphatic alcohols M. Grzesik, M. Kulawska, J. Skrzypek and M. Witczak

411

Kinetics of esterification of acrylic acid with C 3and C 4 aliphatic alcohols 415 in the presence of sulfuric acid as a catalyst M. Grzesik, J. Skrzypek and M. Witczak I

Hydrogenation of carbonaceous adsorbed species formed during the 419 CO/H 2 reaction on a Ru/AI2Q catalyst: experiment and kinetic modeling H. Ahlafi, M. Nawdali and D. Bianchi Application of the continuous two impinging streams reactors in chemical absorption M. Sohrabi and A.M. Jamshidi

423

A kinetic study of Heck reaction of iodobenzene and methyl acrylate using homogeneous Pd/TPP catalyst F.-G. Zhao, B.M. Bhanage, M. Shirai and M. Arai

427

Kinetic and catalytic aspects in the synthesis of polyethylene terephtalate (PET), also through the use of model molecules B. Apicella, E. Santacesaria and M. Di Serio

431

H/D isotopic exchange between oxide surface and spiltover hydrogen 435 on nickel supported catalysts V. Almasan, M. Lazar and P. Marginean Transient studies of adsorption kinetics J. Kanervo, L.B. Backman, A. 0.1. Krause and S.-L. J6msd~-Jounela

439

Modeling the Voltammetric behaviour of cobalt cations inside zeolites 443 M.A.N. Lemos, P. Sousa, F. Lemos, A.J.L. Pombeiro and F. Ram6a Ribeiro Modeling transient tracer studies of complex activation mechanisms of two-atom labelled molecules A.A. Shestov, R. Burch, J.A. Sullivan and V.S. Muzykantov

447

Preparation of ZSM-5 zeolite film on metal support A. Brehm, U. Antons and A. Bekurdts

451

A steady state isotopic transient kinetic analysis of NO reduction over Pt/SiO2 under lean burn conditions A.A. Shestov, R. Burch and J.A. Sullivan

455

Modeling the dynamics of the surface of a carbon C. Palma, I.' Santos Silva, F. Lemos and L. Sousa Lobo

459

Investigation of isomerization kinetics of m-xylene over zeolite based catalysts O. Akpolat a n d G. G[Jnd[Jz Authors' index

463 467

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•

The Symposium "Reaction Kinetics an~d the Development of Catalytic Processes" is the continuation of the very succesful International Symposium "Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis", held in September 1997 in Antwerpen, Belgium. A unique series of top level plenary lectures, but also oral communications and posters mainly focused on the dynamics of catalytic surfaces, the interaction of the reacting molecules with the solid catalyst, the elementary steps of reaction pathways and molecular kinetics. Surface science techniques, molecular modeling, transient kinetic studies, sophisticated and specific reactors are included to a growing extent in the kinetic modeling and the development of catalytic processes. How this is practiced today and how it will evolve in the coming years, what benefit can be expected from a more fundamentally based approach is the aim and scope of the Symposium. G.F. Froment K.C. Waugh

The International Symposium "Reaction Kinetics and the Development of Catalytic Processes" was organized by : The Technological Institute associated with the Royal Flemish Society of Engineers (TI - K VIV). The Technological Institute was founded in 1940 with the aim of disseminating information on scientific and technological development by means of seminars, lectures, courses and conferences. Address:Technological Institute vzw Desguinlei 214, B - 2018 Antwerpen tel: +32 3 216 09 96 fax: +32 3 216 06 89 e-mail : [email protected]

Technologisch Instituut

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Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

Catalysis from first principles J. K. N C r s k o v

Center for Atomic-scale Materials Physics, Department of Physics, Technical University of Denmark, DK-2800 Lyngby, Denmark

Abstract Recent progress in the theoretical description of elementary reactions on transition metal surfaces is discussed. Calculations based on density functional theory and a non-local description of exchange and correlation effects can now be used to predict changes in reactivity from one system to the next. On the basis of the calculations, models can be developed elucidating the "electronic factor" in catalysis.

In the present lecture I will discuss some of the recent progress in a "first principles" understanding of elementary steps in simple catalytic processes. By a "first principles" understanding I mean an understanding based on a detailed quantum mechanical description of the interactions in the reacting system. The theoretical description of interatomic interactions at surfaces has developed immensely over the last few years. The main reason is that density functional theory (DFT) calculational methods have reached a point where the complex systems of interest to catalysis can be treated with a reasonable accuracy. These calculations are parameter free, and have given detailed reaction mechanisms and reaction energetics for a number of elementary reaction on metal surfaces. This in itself gives new possibilities of understanding in detail processes at the molecular level. An equally important development is that along with the new computational possibilities, the models describing variations in adsorption energies of activation barriers for reaction has been developed further. This means that we are now approaching the point where we can make predictions about the direction in which the rate of a reaction should vary form one surface

to the next. Or to put it differently, we are beginning to understand which properties of the clean metal surface governs its reactivity. In the following I will summarize a large number of calculations of adsorption energies and barriers for reactions on transition metal and noble metal surfaces. I will show that the variations can be understood largely as governed by a few metal surface parameters relating to the energy of the local d states and the size of the coupling matrix elements with the relevant adsorbate states.

Calculational details All the calculations presented here are based on slab models of the surface. Ionic cores are described by soft (or ultra-soft) pseudopotentials [1] and the Kohn-Sham one-electron valence states are expanded in a basis of plane waves with kinetic energies below 50 (or 25) Ry. The exchange-correlation energy and potential are described by the generalized gradient approximation (PW91) [2, 3]. The self-consistent PW91 density is determined by iterative diagonalization of the Kohn-Sham Hamiltonian, Fermi-population of the Kohn-Sham states (kBT=0.1 eV), and Pulay mixing of the resulting electronic density [4, 5]. All total energies have been extrapolated to kBT=O eV. In the calculations the adsorbate(s) and the uppermost atomic layers of the metal slab have been allowed to relax to the energetically most favorable position. Transition states for reactions are found by varying a single coordinate (the reaction coordinate) and minimizing all other degrees of freedom of the reactants. In cases where it was not evident which coordinate would be the reaction coordinate, the latter was found by the iterative method of Ulitsky and Elber[6].

2

Results

In Fig. 1 I have collected a large number of data from the literature[7, 8, 9, 10, 11, 12, 13], all extracted from DFT calculations. These data represent calculated adsorption energies of atomic and molecular adsorbates as well as activation energies for surface reactions. They all describe a situation where

0.0

~, ~,.

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"~ ~ - 1 . 0

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t

Cu(11e~,,~Cu:Cu,Pt(111 ) Cu:Ni@Cd~~ Ft_~x~d/Ru(100) 9Cu/Pt(111) ~

Pd(111) Ni(1111

]

Pt-he' ?t(10~ t:Cu3Pt(111) Ni:Ni@~Cn11 ~ Ru(0001)-2.70A

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,,

I

"" I

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/Pd(21l)--terr~ce ~O/Pd(lll) N//Pd(211)-terrace" e O/Pd(211)-st~

-4.5

N/Pd(lll)~ N/Pd(211)-step

2 -5.0

O/Ru(0001)-2.70A

--....

-5.5

O/Ru(0001)-2..80A I

I

I

I

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CO/Ru(0001)-2.70A ~'"t ~m CH~[I@Cu(lll

1.0

HJCu~ o-.

~CO/Ru(000,)--~.,4A

H~/Cu(100)CO/Ru(0001)-2.78A

0.5

0.0

--0.5

~d(110/P~i11) I

-3.0

!

I

!

-2.5

-2.0

-1.5

ed

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-1.0

Figure 1' Molecular ( ~ l m ) and atomic ~,a-Jchern {~-atornic~ } binding energy as a function of the d-band center (ed)of the metal surface (top and middle panel, respectively). The barrier for dissociation of small molecules, referenced to gas phase zero, as a function of ed is shown in the bottom panel. C o m m o n colors are used for data corresponding to the same metal throughout the three panels. Lines drawn represent best linear fits. Specific data points are taken from: [7] for N, O, and N O on Pd, [10, 11] for C O on Pt, Ni, Cu, and Pd, [9, 12] for CH4 on Ni, [8] for/'/2 on Cu.

From [13].

the adsorbate interacts with the same kind of metal atom(s) in the same local geometry, but the environment varies. The environment has been changed four different ways: i. By straining the surface. 2. By considering different facets and stepped surfaces. 3. By considering one metal as a pseudomorphic

overlayer on another.

4. By considering surface alloys. The calculated values of adsorption energies and barriers for dissociation show large variations. It can for instance be seen that CO bound on top of a Pt atom can have adsorption energies varying by about 1 eV (I00 kJ/mol) depending on the surroundings of the Pt atom in question. The largest adsorption energies are found for adsorption on steps and kinks and the lowest on the compressed hexagonal overlayer structures on Pt(100). Such effects have been observed experimentally. An increased adsorption energy for CO at steps has been observed using thermal desorption spectroscopy[14], and a similar effect is observed in small metal clusters on supports[16]. The relative inactivity of the compressed hexagonal overlayer on Pt(100) has also be observed experimentally[17]. Steps on Pd surfaces (Pd(211) consists of (III) facets and steps) are also seen to have a lower barrier for dissociation of NO than the fiat (III) surface, again in agreement with experiment[15]. Changes in adsorption energy and dissociation barriers for overlayer structures and surface alloys have also be observed in many cases[18, 19, 20]. Most recently the effect of strain on adsorption properties has been observed directly[21]. In Fig. 1 we show the data as a function of the center of mass of the density of states projected onto the atomic d states of the clean surface. For convenience, we use all the d states here, instead of the ones with the correct symmetry for bonding with the various adsorbates. This makes no major difference, when the adsorption geometry remains similar. Fig. 1 shows clearly that for a given adsorbate or reaction and for a given metal, the variations in adsorption energies or activation energies are governed largely by the variations in the energy of the surface d-states. There are good reasons for this[22]. The interaction between the adsorbate states and the metal d states is an important part of the interaction

energy, and while the sp bands of the metal are broad and structure-less, the d bands are narrow, and small changes in the environment can change the d states and their interaction with adsorbate states significantly. The d-band center (Cd) is the simplest possible measure for the position of the d states. The correlation between interaction strength (adsorption energy or activation energy barrier) and Cd holds for many different adsorbates and metals. Calculations for/-/2 dissociation on transition and noble metals have shown a similar relationship to hold also when different metals are compared [22, 23]. The generality of this correlation is a simple manifestation of the fact that the coupling to the d states depends on the position of the d states relative to the Fermi level. This tendency is also elucidated by simple models describing interactions between atomic or molecular adsorbates and transition states with metal surfaces[22]. In addition, the correlation between the interaction strength and the dband center found in the framework of these simple models appears to be independent of the adsorbate and the metal, in agreement with the trends revealed in the large-scale total energy calculations, as illustrated with the data in Fig. 1. The identity of the metal involved shows up in the strength of the effect, that is, the slope of E(cd) through the size of the coupling matrix element. The relative ordering in the coupling strength is 5d > 4d > 3d following the relative sizes of the d-wavefunctions[22]. Having established that the (local) transition metal d band center, Cd, is an important parameter determining the ability of a metal atom to interact with a reacting atom or molecule, the question arises what determines the variations in Cd from one system to the next. First, consider the case where the metal is simply strained. When the lattice is expanded parallel to the surface, the overlap between the d-electrons on neighboring metal atoms becomes smaller, the band width decreases and to keep the d-occupancy fixed, the d states have to move up in energy (for a more than half-filled band). When the structure of the surface changes for instance by introducing a step, the local d-projected density of states is not changed due to strain, but due to a change in the number of metal neighbors, the metal coordination number. The general rule is that the lower the coordination number, the smaller the local band width and the higher the Cd (for metals with more than half-filled d-bands). For alloys and overlayers, a large portion of the change in Cd can be attributed to changes in the metalmetal distances in the surface[24], as it is the case for the strained slab.

3

Conclusion

In conclusion, I have illustrated that large scale density functional calculations can now be used to predict adsorption energies and barriers for reactions at metals surfaces. In addition we can begin to understand the physical origin of variations in adsorption energies and barriers from one system to the next. There is a general correlation between the reactivity of a surface shifts in the center of the metal d-bands (Cd). The coupling strength between the adsorbate states and the d states is another important parameter. Together with the number of d-eletrons it determines the variations on interaction strengths from one metal to the next[22]. In the discussion here I have deliberately always considered classes of systems where the local bonding geometry is the same. This has allowed me to single out the electronic effects. An additional factor is the geometry both the arrangement and identity of the surface atoms. The geometrical effects can be as important as the purely electronic effect[8]. When combined with experimental investigations of surface structure and reactivity, the new developments in the theory of surface reactivity makes it possible to suggest new catalytically interesting systems. An example of such a catalyst design from first principles have been reported recently[25]. The present work was in part financed by The Danish Research Councils through The Center for Surface Reactivity and grant #9501775. The Center for Atomic-scale Materials Physics is sponsored by the Danish National Research Foundation.

References [1] D. H. Vanderbilt, Phys. Rev. B, 41, 7892, (1990). [2] J. P. Perdew, et. el., Phys. Rev. B, 46, 6671, (1992). [3] J. A. White and D. M. Bird, Phys. Rev. B, 50, 4954, (1994). [4] G. Kresse and J. Forthmiiller, Comput. Mat. Sci., 6, 15, (1996). [5] B. Hammer, L. B. Hansen, and J. K. N~rskov, Phys. Rev. B, to be published.

[6] A. Ulitsky and R. Elber, J. Chem. Phys. 92(2), 1510 (1990). [7] B. Hammer, Faraday Discuss. Chem. Soc., (submitted). [8] P. Kratzer, B. Hammer, and J. K. Ncrskov, Surf. Sci. 359, 45 (1996). [9] P. Kratzer, B. Hammer, and J. K. NCrskov, J. Chem. Phys. 105, 5595 (1996). [10] B. Hammer, O. H. Nielsen, and J. K. NCrskov, Catal. Lett. 46, 31 (1997). [11] B. Hammer, Y. Morikawa, and J. K. Ncrskov, Phys. Rev. Lett. 76, 2141 (1996). [12] P. M. Holmblad, et. al., Catal. Lett. 40, 131 (1996). [13] M. Mavrikakis, B. Hammer, and J. K. Ncrskov, Phys. Rev. Lett. 81, 2819 (1998). [14] M. R. McClellan, J. L. Gland, F. R. McFreeley, Surf. Sci. 112, 63 (1981); B. E. Hayden, K. Kretzschmar, A. M. Bradshaw, and R. G. Greenler, Surf. Sci. 149, 394 (1985); J. T. Yates, J. Vac. Sci. Technol. A 13, 1350 (1995). [15] Q. Gao, R. D. Ramsier, H. Neergaard Waltenburg, and J. Yates, Jr., J. Am. Chem. Soc., 116, 3901 (1994). [16] C. R. Henry, Surf. Sci. Reports 31,231 (1998). [17] Y. Y. Yeo, C. E. Wartnaby and D. A. King, Science 268, 1731 (1995). [18] J. A. Rodriguez, and D. W. Goodman,

Science 257, 897 (1992).

[19] E. Kampshoff, E. Hahn, and K. Kern, Phys. Rev. Lett. 73, 704 (1994). [20] J. H. Larsen and I. Chorkendorff, Surf. Sci. 405, 62 (1998). [21] M. Gsell, P. Jakob, and D. Menzel, Science 280, 717, (1998). [22] B. Hammer and J. K. Ncrskov in: "Chemisorption and Reactivity on Supported Clusters and Thin Films", 285-351, R. M. Lambert and G. Pacchioni (Eds.), (Kluwer Academic Publishers, The Netherlands, (1997).

10 [23] A. Eichler, G. Kresse, J. Hafner, Surf. Sci. 397, 116 (1998). [24] A. Ruban, et. el, J. Mol. Catal. A, 115, 421, (1997). [25] F. Besenbacher, I. Chorkendorff, B. S. Clausen, B. Hammer, A. M. Molenbroek, J. K. Ncrskov and I. Stensgaard, Science 279, 1913 (1998).

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

11

Complexities and Dynamics of the Enantioselective Active Site in Heterogeneous Catalysis R.Raval 1'2., C.J.Baddeley ~, S. Haq 2, S.Louafi ~, P. M u r r a y 2, C.Muryn 2, M.Ortega Lorenzo ~ and J.Williams x'2 1Leverhulme Centre and 2Surface Science Centre, D e p a r t m e n t of Chemistry, University of Liverpool, Liverpool L69 3BX. Abstract There is now renewed interest in utilising chirally modified metals as heterogeneous catalysts for enantioselective reactions. Effective systems have been produced by first modifying a metal surface with chiral molecules and, subsequently, conducting the enantioselective reaction on the modified surface. How stereocontrol is achieved by the chirally modified metal is still a m a t t e r of speculation and a lack of sensitive spectroscopic data of the modified surface has precluded the development of detailed models. In this work we report results relevant to the asymmetric hydrogenation of ~-ketoesters over R,Rtartaric acid and S-alanine modified metal surfaces. Using a powerful range of surface spectroscopies on modified surfaces created under ultra-clean a n d controlled environments, we show that the adsorbed modifiers, R,R-tartaric acid and S-alanine, display rich and complex phase d i a g r a m s in which the chemical n a t u r e and 2-dimensional organisation of the chiral units varies considerably as a function of surface coverage and t e m p e r a t u r e . From this work it would seem that the models suggested for stereocontrol involving oneto-one interaction between the adsorbed chiral modifier and the r e a c t a n t (substrate) are, perhaps, too simplistic and careful account needs to be taken of the complexities and dynamic interplay between different modifier phases at a surface in order to fully understand the nature and scope of stereocontrol. 1. I N T R O D U C T I O N The enantioselective reaction, undoubtedly, represents the most s t r i n g e n t control of a reaction at the molecular level, stereodirecting a chemical process so that only one optical component is produced. As the industrial need for optically pure compounds increases at a fast rate, there is a strong impetus to design new enantioselective catalysts, of which the heterogeneous catalyst promises a n u m b e r of practical advantages over the more widely-used homogeneous systems. One of the most successful method of bestowing stereodifferentiation to a heterogeneous catalytic system is to attach chiral modifier molecules to the reactive metal surface. The enantioselective hydrogenation of ~-ketoesters over modified nickel surfaces is a case in point and s p e c t a c u l a r *corresponding author, e-mail: [email protected]

12 success has been reported for the asymmetric hydrogenation of acetylacetone and methyl acetoacetate (MAA) over R,R-tartaric acid and S-alanine modified Ni surfaces, with enantioselective excesses of >90% reported for the R-product [1-3]. Despite this success, and despite the fact that this is one of the most intensively studied enantioselective systems, the central question in this ~field still r e m a i n unanswered: what is the nature of the active site present on the modified surfaces? The formulation of a clear answer has been h a m p e r e d by the sheer complexity of the working catalyst system, the large n u m b e r of interrelated variables which affect optical yields, ranging from modification time, modification temperature, modifier concentration and pH, and by the difficulty of utilising suitable surface-sensitive probes to monitor the nature of the modified adlayer. In a bid to answer this question, we have launched a study of model chiral surfaces created by the adsorption of optically active modifiers on defined metal single crystal surfaces and scrutinised using the formidable array of surface analytical tools that are now available. In this paper we report and review our findings with regard to the 2-dimensionally anisotropic Cu(ll0) surface. Three main in situ techniques were employed to probe the modifiermetal interface: Reflection Absorption Infrared Spectroscopy (RAIRS), Low Energy Electron Diffraction (LEED) and Scanning Tunnelling Microscopy (STM). Each technique provides complementary information on the modified interface, enabling a detailed picture of the interface to be constructed. The high sensitivity, high resolution vibrational data obtained by F o u r i e r - t r a n s f o r m RAIRS provides direct information on the chemical nature of the adsorbed modifier molecules and their perturbation by the surface, while application of the strict dipole selection rule that operates for the technique (i.e only vibrations that produce a dynamic dipole normal to the metal surface are observed) allows the orientation of the species to be determined. LEED monitors the long-range 2dimensional periodicity of the adlayer while STM provides information on the local a r r a n g e m e n t s of the modifier molecules within the domains formed at the surface. It is emphasised that the sensitivity of all these techniques allows the detection of submonolayers of modifier at the metal surface which, on a typical single crystal sample with a surface area of 1 cm 2, represents the ability to sense a nanomole or less of the modifier. 2. EXPERIMENTAL Experiments were carried out in two separate surface analysis i n s t r u m e n t s , each with a base pressure of better than 2 x 101~ mbar. The first chamber was interfaced, via auxiliary optics, with a Mattson 6020 FTIR spectrometer to allow RAIRS experiments to be conducted by the single reflection of IR light, at near-grazing incidence, from the Cu(ll0) surface. RAIR spectra were recorded throughout a continuous dosing regime as s a m p l e single beam infrared spectra and subsequently ratioed against a reference background single beam spectrum from the clean Cu(ll0) surface in order to obtain AR/R ~ spectra. A liq. N 2 cooled HgCdTe detector allowed IR data to be collected over the 4000-800 cm 1 region and a polariser placed in front of the detector ensured that only p-polarised light was detected. All spectra were

13 recorded at 4 cm 1 resolution with the coaddition of 256 scans. The chamber w a s also equipped with LEED, quadrupole mass spectrometry, Auger Electron Spectroscopy (AES) and sample cleaning facilities. LEED patterns displayed on the phosphor screen in the chamber were captured and digitised by a CCD video camera interfaced to a computer. STM experiments were carried out in a separate Omicron V a k u u m p h y s i k chamber which was also equipped with STM, LEED, AES and sample c l e a n i n g faclities. The STM experiments were carried out by creating the required adlayer by specific exposures of the modifier molecules at the required t e m p e r a t u r e and then cooling to room t e m p e r a t u r e to record the data. All STM images were acquired in constant current mode. Depending on the tolerance of the adlayer to electron beam damage, LEED experiments were conducted before or after the STM experiments, in order to provide a direct correlation between the two sets of data. Prior to all experiments, the Cu(ll0) crystal was cleaned by cycles of Ar § sputtering and annealling at 600K. The sample cleanliness and surface order were monitored by AES and LEED, respectively. S-Alanine (99%) was obtained from Aldrich Chemical company and R,RTartaric acid (99%) from Fluka Chemical Company and were used without further purification. The required modifier was contained in a s m a l l electrically heated glass tube separated from the main v a c u u m chamber by a gate valve and differentially pumped by a turbo molecular pump. Before sublimation the modifier sample was pumped for a few hours and outgassed at a t e m p e r a t u r e o f - 3 5 0 K . The modifier was then heated to a t e m p e r a t u r e of - 3 7 0 K and exposed to the copper crystal. During sublimation the main c h a m b e r pressure w a s - - 2 x 10 .9 mbar, ensuring ultra-clean deposition conditions. 3. R E S U L T S AND DISCUSSION The enantioselective hydrogenation of MAA to give R-methyl-3hydroxybutrate or S-methyl-3-hydroxybutrate over modified Ni is shown in Figure 1.

H

OH

0

N icatalyst

I IoI~ I ~

S-Meth yl-3-hyd roxyb uty rate

Nicatalyst

prfferreaface for attack

on Ni premodified by R,R-Tartar ic acid or S-Alanine

Methylacetoacetate

R-Me thyl-3-hyd roxyb u tyrate

Figure 1. Hydrogenation reaction of MAA on modified Ni catalysts.

14 Since MAA is a planar molecule, the optical activity of the product is simply determined by which molecular face the hydrogen attack occurs at. For the unmodified Ni, attack on each face occurs with equal probability, resulting in a racemic product mixture. However, for Ni premodified by R,R-tartaric acid or S-alanine, the reaction is stereodirected to give almost exclusively the Rproduct. It has been established by many workers [1-3] that enantioselectivity is conferred prior to the hydrogenation step. In other words, the modified surface m u s t align the MAA molecule so that hydrogenation is effected along one selected face only. Therefore, the architecture of the active site, composed of the modifier plus metal, must be intricately implicated in the enantioselectivity. In this paper, we show that chiral modifiers, such as R,R-tartaric acid and S-alanine, adsorbed at metal surfaces possess a rich and complex p h a s e diagram, where the nature of the modified surface and, hence, the n a t u r e of the available active site varies critically as a function of coverage a n d temperature. Although the work presented here is confined to the defined Cu(ll0) surface, the general conclusions have direct implications for the catalytically active Ni surface. 3.1 C h e m i c a l n a t u r e and self-organisation of S-Alanine on C u ( l l 0 )

The S-alanine molecule can exist in four different forms: the neutral, the cationic, the anionic and the zwitterionic. Therefore, the first question that arises is which form is stabilised at a metal surface. Figure 2 displays the RAIR spectra obtained for adsorption carried out at 300K up to a m a x i m u m first layer saturation coverage of just over 0.33 monolayer (ML), representing 1 chiral molecule per three surface metal atoms. Although complex vibrational spectra are observed, identification of the adsorbed species is readily made since the frequency of the observed bands are almost identical to those reported for the metal-alanine complexes, Cu-(ala) 2 and Ni-(ala) 2 [4] in which alanine exists in its anionic form. Direct evidence for the anionic species is provided by the symmetric vs(COO-) vibration at 1411 cm 1 and the asymmetric Vas(COO') vibration at 1626 cm 1, attributed to the deprotonated carboxylate functionality. In addition XPS data for the adsorbed phase show a N l s binding energy at between 399.5 and 399.9 eV, consistent with a NH 2 unit rather t h a n the protonated NH3 § group which displays a N ls binding energy that is almost 2 e V higher [5]. By application of the strict RAIRS dipole selection rule, the orientation of the alanine molecule can be pieced together by considering vibrations of the distinct functional groups on the molecule. A careful analysis, reported in detail elsewhere [6], shows that at low coverages, the molecule essentially straddles across the close packed metal rows of the Cu(ll0) surface, with both the carboxylate and the NH 2 functionalities in close proximity to the surface and contributing to the bonding with the metal. The two carboxylate oxygens are placed equidistant from the surface, rendering the v~(COO) vibration active and the Va~(COO') vibration inactive, Figure 2 a-c. Similarly, it can be deduced that the NH 2 plane is almost parallel to the surface and the CH and CH 3 units are held at an angle away from the surface normal. A detailed molecular

15

Figure 2. RAIR spectra recorded with increasing coverage in the first layer of S-alanine on C u ( l l 0 ) a t 300K. The molecular orientations adopted at low coverage (spectra a-c) and high coverage (spectra d,e) are depicted on the right. orientation is depicted in Figure 2, top right. Recent scanned energy mode N l s and Ols Photoelectron Diffraction studies [7] of glycine, the simplest aminoacid, adsorbed on Cu(ll0), have also come up with a strikingly s i m i l a r structural model. The detailed adsorption site information available from this technique shows that the carboxylate group is aligned along one row of Cu(ll0) atoms with the oxygens atoms located in on-top metal sites. The m o l e c u l a r skeleton then essentially 'bends' over to bring the NH 2 unit close to the adjacent Cu row. Theoretical calculations [8] on the glycine/Cu(ll0) system also favour such an adsorption geometry. As coverage is increased in the first layer, a second type of anionic species is stabilised alongside the already adsorbed alanine, leading to the emergence of a strong Vas(COO') vibration at 1626 cm 1 in Figures 2 d,e. A detailed analysis [6] of the changes observed reveals that a new, differently oriented alanine species is now coadsorbed on the surface, with the carboxylate unit tilted with respect to the surface and the methyl group held almost vertical along the surface normal. A schematic orientation is shown in Figure 2, bottom right. It is interesting to note that this tilted adsorption configuration is almost identical to that proposed on modified supported Ni catalysts [1,2]. The obvious question then to ask is which configuration is the more stable and under what conditions is it stabilised. In a series of experiments to follow the effect of t e m p e r a t u r e and coverage on adsorption behaviour, we find that the geometry initially adopted

15 in the low coverage phase is the thermodynamically favoured one throughout the temperature and coverage range studied. For example, heating the h i g h coverage monolayer created at 300K to 470K leads to the conversion of all the tilted species to the adsorption geometry stabilised at low coverages. This configuration is then maintained upon recooling to 300K. We now turn to a second aspect of chirally modified metal surfaces w h i c h , hitherto, has not been directly addressed in the literature, namely the selforganisation of the modifier species at the surface. In particular, we wished to investigate w h e t h e r strong lateral interactions in the adlayer lead to the creation of ordered, self-assembled arrays. LEED and STM experiments c a r r i e d out as a function of coverage and t e m p e r a t u r e show that, in fact, a r e m a r k a b l e degree of self-organisation is demonstrated by the S-alanine/Cu(ll0) system. I n particular, two main ordered structures are formed: a plgl(3x2) s t r u c t u r e which forms at high t e m p e r a t u r e and at saturation monolayer coverage, and another phase which can be described in matrix notation as the (5 3, -2 2) s t r u c t u r e where the relationship between the metal surface net, given by the vectors a, b and the overlayer net vectors a,' b' is given by:

(;:)

(o1

For convenience, the m a t r i x formation shown above will be r e p r e s e n t e d throughout in the text as (G~ G~2 , G2~ G22). The (5 3, -2 2) structure is formed when adsorption is carried out between 300-420K. This structure is m a i n t a i n e d over a wide coverage range and, therefore, must describe a general motif at the surface capable of supporting a varying local domain coverage. Figure 3 shows the STM image and the schematic adlayer of the (5 3, -2 2) structure formed at 400 K. STM data shows that the two-dimensional order of the S-alanine extends over long distances (> 400.,~) across the surface. Another interesting fact that emerges is t h a t the S alanine molecules seem to be arranged in regular g r o u p s of eight, which are aligned at a definite angle along the surface. This growth

Figure 3. a) STM image (100/~ x 100/k) and b) schematic of the (5 3, -2 2) S-alanine structure on Cu(110).

17 axis is not coincident with either of the major directions of the metal surface, thus destroying the two mirror planes that exist at the fcc(ll0) surface. In other words, the growth direction of the S-alanine directly imposes chirality on the metal surface! A final aspect to note is that this structure now n a t u r a l l y possesses empty, chiral channels which may be important in stereocontrol [9]. The plgl(3x2) structure which can only be created at saturation first layer coverage and high temperatures of >460K displays a very different twodimensional structure, shown in Figure 4. Again, there is a preferred growth direction, but now the molecules at the corners of the primitive (3x2) mesh are aligned so that the mirror plane along the [001] crystallographic direction is retained. New STM data of this structure, however, show that there is another molecule placed within the (3x2)mesh, causing the mirror plane to be transformed into a glide plane [9].

Figure 4. Schematic diagram with Figure 5. Schematic diagram of the various arrangement of molecules in the phases that can be created by S-alanine plgl(3x2) structure. The glide on Cu(ll0). plane is marked. The stylised phase diagram shown in Figure 5 summarises the different types of modified surfaces that can be created by S-alanine as a function of coverage and temperature. It is important to emphasise that there is considerable interplay between the different structures. For example, m a n y structures are only stabilised under kinetically hindered conditions. The most favoured molecular orientation is the one adopted at low coverages, Figure 2, top right. Our work suggests that growth of the (5 3, -2, 2) structure allows this optimum geometry to be realised. However, once this structure is completed new molecules can only be accommodated in the empty channels if they possess a smaller 'footprint'. As a result, the tilted species is formed. To allow all molecules to adopt the preferred orientations would require dismantling of the (5 3,-2 2) structure to form the new plgl(3x2) structure. This process has a high activation barrier associated with it and, therefore, only becomes possible at high temperatures (T > 460K).

18 3.2 C h e m i c a l n a t u r e a n d self-organisation of R,l~Tartaric Acid on C u ( l l 0 )

Like S-alanine, R,R-Tartaric acid is also versatile in the chemical structures it possesses, ranging from the neutral form, to the monotartrate where one of the carboxylic acid group deprotonates to give a carboxylate, to the bitartrate where two carboxylate functionalities exist on the molecule. I n addition, carboxylic acids are also well known for their ability to dimerise. RAIRS data obtained after adsorption at 300K, Figure 6, show bands due to the v(C=O) and the vs(COO) vibrations at 1703 and 1436 cm 1, respectively, consistent with monotartrate formation, involving the deprotonation of one acid group only. At 300K, the monotartrate form persists throughout the monolayer regime while multilayer formation is accompanied by condensation of neutral acid molecules [10], producing a v(C=O) fingerprint vibration at 1757 cm 1.

M Low o nCoveraoe ol~av---er'-':~

(D

0

C).. x W

Low coverage monotartrate

(~0)

High coverage monotartrate

(~)

Saturation monolayer

.c_ 0o C) },,., 0

_c

Multilayer; neutral acid molecules

1 .

3850

.

.

.

3300

.

.

.

2750

.

.

.

2200

.

.

1650

.

1100

Wavenumber, cm -1

Figure 6. RAIRS data of R,R-tartaric acid adsorbed on Cu(ll0) at 300K to form a monolayer and then multilayers.The main phases observed are outlined. R,R-tartaric acid also displays strong self-organising behaviour and LEED and STM experiments show that an extremely rich phase diagram exists for this modifier at the metal surface, resulting in the creation of four different types of 2-dimensionally ordered adlayers between 300 and 480K. Two of these phases, with the matrix structures (4 0, 2 3) and (4 1, 2 3), are formed upon adsorption at 300K, and their existence range is shown in Figure 6. The (4 0, 2 3) structure exists at low coverages while the (4 1, 2 3) structure is formed at n e a r saturation monolayer coverages. It can be seen from the matrix notation that the two structures are closely related and transition from low to high coverage essentially involves a contraction and rotation of the surface unit mesh.

19

Figure 7 a) STM image of a tartrate island with (4 0, 2 3) structure; b) s c h e m a t i c of the t a r t r a t e overlayer and; c) proposed monotartrate orientation at surface. Details of the (4 0, 2 3) structure displayed in Figure 7 show that it involves very closely packed monotartrate molecules. The monotartrate entities are thought to be adsorbed to the metal via the carboxylate group, with the acid group held away from the surface, Figure 7c. This allows the modifier to exert a smaller footprint and, thus, allows a high packing density to be m a i n t a i n e d at the metal surface, calculated to be 0.25 ML (i.e 1 molecule per 4 surface atoms). The density of this structure precludes the accommodation of a bigger molecular footprint which would allow the second acid group to lie in close proximity to the surface.

Figure 8 a) RAIRS and b) STM data showing transformation of the (4 0, 2 3) monotartrate structure into a bitartrate (9 0, 1 2) phase.

20 Experiments conducted over a wide coverage and t e m p e r a t u r e phase space [10] reveal that, again, a considerable amount of interplay exists between the modifier structures. The most remarkable of these is the transformation of a n adlayer consisting of islands of the (4 0, 2 3) structure. At 300K, this t r a n s f o r m a t i o n is very slow and undetectable during the normal course of a n experiment. However, at 350 K the relaxation can be easily followed by I R spectroscopy, Figure 8a, which shows the gradual attenuation of the v(C=O) vibration at 1700 cm 1, accompanied by the appearance of new bands in the vs(COO) region at 1412 cm 1. We interpret this as a transformation of the m o n o t a r t r a t e form into the bitartrate form. Clearly, this process involves a n activation barrier. STM data, Figure 8b, provide a beautiful record of this p h a s e evolution. It can be seen that, over time, molecules in the high local density (4 0, 2 3) structure migrate out onto clean metal areas and form a new, lower density phase with a (9 0, 1 2) structure. At 300K, the mobility of the modifier molecules is so low t h a t the (4 0, 2 3) islands remain kinetically frozen in. However, as the t e m p e r a t u r e is increased to 350K, significant mobility is initiated and over 15 to 20 minutes, the transformation is seen to occur in the RAIRS and the STM experiments. When the adsorption is carried out at 430 K, the mobility of the impinging molecules allows the new (9 0, 1 2) structure to form immediately.

Figure 9. Schematic diagram of the various phases that can be created by R,Rtartaric acid on C u ( l l 0 ) and the interconversions between them.

21 Interestingly, an equilibrium exists between the monotartrate and bitartrate species. The bitartrate form is favoured at low local coverages only and adsorption of further molecules onto the fully converted (9 0, 1 2) bitartarate structure, causes a reconversion of the modifier to the monotartrate form. A balance, therefore, exists between optimising modifier-metal interactions and optimising modifier density at the surface. In this sense the tartaric acid modifier system is very different from the S-alanine system, where the anionic form in a particular adsorption geometry is always thermodynamically favoured. The complex and dynamic transformations of the tartaric acid/Cu phases are best represented by the schematic phase diagram in Figure 9. 4. CONCLUSIONS AND IMPLICATIONS FOR ENANTIOSELECTIVE CATALYSIS Using a range of ultra-sensitive surface spectroscopies, we have shown that a metal modified by the chiral molecules S-alanine and R,R-tartaric acid displays a rich and complex phase diagram where the description of the available active site varies significantly with coverage and temperature. The variation is manifested in two major aspects: (i) the chemical form and orientation of the adsorbed modifier and (ii) the 2-dimensional array created by the self-organisation of the modifiers at the surface. In addition, a delicate balance often exists between the different modifier phases and the conditions required to stabilise each one are different. This complex surface chemistry may be one of the major underlying reasons for the strict modification conditions required to achieve high enantioselectivities. Another important conclusion that arises from this work, is that the single molecule description of the active site put forward in the literature for these modifiers may not be sufficient to describe the enantioselective site and that account needs to be taken of the larger molecular templates and channels that are generated at the surface which may be crucial in stereocontrol of the reaction.

Acknowledgements

We are grateful to the EPSRC and the University of Liverpool for equipment, postdoctoral researcher and PhD studentship grants which have allowed this research to be funded. 5. REFERENCES Y. Izumi, Adv. Catalysis, 32 (1983) 215. A. Tai and T. Harada in 'Tailored Metal Catalysts' Y. Iwasa (ed), D. Reidel Publishing Co., 1986. G. Webb and P.B. Wells, Catalysis Today, 12 (1992) 319. G.C. Percy and H. Stretton, J. Chem. Soc., Dalton Trans. (1976) 2429. D.T. Clark, J. Peeling and L. Colling, Biochimica et Biophysica Acta, 453 (1976) 533.

22 6 7 8 9 10

J. Williams, S. Haq and R. Raval, Surface Science, 368 (1996) 303. N.A. Booth, D.P.Woodruff, O. Schaff, T. Gei~el, R. Lindsay, P. Baumgartel and A.M. Bradshaw, Surface Science, 397 (1998) 258. J. Hasselstrom, O. Karis, M. Weinelt, N. Wassdahl, A.Nilsson, M.Nyberg, L.G.M. Pettersson, M.G. Samant and J. Stohr, Surface Science, 407 (1998) 221. J. Williams, S. Haq, S.Louafi, P. Murray and R. Raval, J. Phys. Chem., to be submitted. M. Ortega Lorenzo, S. Haq, T. Bertrams, C. Muryn, R. Raval and C.J. Baddeley, J. Phys. Chem., to be submitted.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

23

Intrinsic activities and pore diffusion effect in hydrocarbon cracking in steamed Y zeolite H.H. Kung Department of Chemical Engineering, Northwestern University, Evanston, Illinois 60208-3120, U.S.A.

Abstract In the literature, it has been shown that the hydrocarbon cracking activity of a zeolite particle can be greatly enhanced by steam treatment. The extent of enhancement depends on the process condition. Recently, it has been shown that under conditions where the reaction is dominated by the monomolecular mechanism, the turnover frequency and the intrinsic activation energy for cracking are practically unchanged by steaming. However, a much larger difference in activity can be observed under conditions where bimolecular and oligomeric cracking dominate. Thus, the phenomenon of enhanced activity by steaming can be explained by the possibility that the bimolecular and oligomeric cracking reactions are pore diffusionlimited. Steaming generates defects in the zeolite particle, increases the external surface area for diffusion, which then leads to the large enhancement in the observed activity.

1. I N T R O D U C T I O N Catalytic hydrocarbon cracking is an important step in the conversion of crude oil to gasoline. Industrially, this reaction is carried out in a fluidized bed of catalyst that is composed of USY (ultrastable Y) zeolite in a matrix. It has been well established that the reaction requires strong Bronsted acid sites [1]. It has also been known that steaming, that is, treatment with high temperature water vapor, has a strong effect on the activity of the Y zeolite [2,3]. An appropriate steam-treatment of a HY zeolite could enhance the cracking activity by over two orders of magnitude [4]. However, excessive steaming leads to decline in activity from the optimal value. It is of great interest and practical importance to understand the origin of the enhanced activity due to steaming. A prevalent explanation is that steaming generates a small number of highly active sites [5,6,7,8,9]. These sites can be selectively poisoned by alkali metal ions or ammonia. The nature of these sites are not known. One proposal is that they are framework Bronsted acid sites which are made stronger by extraframework A1 species which delocalize the framework negative charge at the site. The extraframework A1 species are generated by steaming. Another explanation is that steaming generates new sites which, through synergistic effects with the framework Bronsted acid sites, effectively enhance the cracking activity [10]. The nature of this synergistic effect is not known. One possibility is that the new sites enhance dehydrogenation of the alkane reactant, and, consequently increase the surface carbenium ion concentration. The result is faster hydride transfer between the alkane reactant and surface

24 carbenium ions. However, there is evidence that the dehydrogenation activity associated with extraframework A1 species does not contribute significantly to the steady state activity [ 11 ]. There are numerous studies reported in the literature in attempts to understand this phenomenon of steaming. However, when examining the results from different studies, it is often found that there are important discrepancies that cannot be easily explained using these models. For example, in the selective poisoning experiments using alkali ions or ammonia to determine the density of unusually active sites [7,8,9], the density obtained differs somewhat depending on the poison used. Another example is that it has been suggested that the 3600 cm 1 hydroxyl band in the infrared is due to the unusually active Brtmsted acid sites in a steamed Y zeolite. Its intensity correlates with the catalytic activity [7]. However, on a La-exchanged Y zeolite, such a correlation cannot be made [12]. The role of extraframework A1 species is not clear either. Some studies reported a linear relationship between the density of extraframework A1 species and the cracking activity [ 13,14]. However, attempts to quench the enhanced activity by extracting the extraframework A1 species did not provide conclusive results. In some studies, a small decrease in activity was observed [15,16], whereas in some other studies, either no change in activity [17], or a slight increase was observed [ 18,19]. The examples cited above illustrate the apparent variations in the literature results and the confusion they might generate in attempts to understand the phenomenon. One complication in comparing results from different studies or results for different catalysts in a study is the different conditions under which the measurements were made. In fact, the extent of enhancement of activity by steaming appears to depend on the process variables, including the hydrocarbon used and the time.on-stream [20], the latter is because the zeolite catalyst often deactivates in the cracking reaction. Further complicating the situation is the fact that different mechanisms participate in the cracking reaction [21,22]. The relative contributions of the different mechanisms to the observed rate could very well depend on the process conditions and the state of the catalyst. Thus, it appears that in order to quantitatively compare the true activities of different catalysts, it would be much more meaningful if the activity data were obtained when the catalysts are at the same state and for the same reaction mechanism (as explained below). Here, the results of our attempt to make such a comparison are summarized. 2. REACTION MECHANISM It has been suggested that three mechanisms participate in hydrocarbon cracking over zeolite catalysts: monomolecular, bimolecular, and oligomeric [21,22,23]. In monomolecular cracking, the reactant alkane molecule interacts with a Bronsted proton of the zeolite to form a high energy transition state which consists of a hydrocarbon containing a five-coordinated carbon atom and a substantially lengthened Bronsted O-H bond (I) [24]. Cracking of this transition state leads to the formation of dihydrogen, smaller alkanes and alkenes. The formation of the high energy transition state is believed to be the rate limiting step. In the bimolecular mechanism, conversion of the reactant alkane proceeds via hydride transfer between an alkane molecule and a surface carbenium ion, which may be more appropriately described as an alkyl group o-bonded to a surface oxygen ion [25]:

25

Rk,.i2 ?' I I i

0

I~.H + 1~.+(ad)-- ~+ (ad) + ~ H

(1)

Oligomeric cracking is similar to bimolecular cracking, except that the hydride transfer involves a surface carbenium ion that is larger than the reactant alkane, and is probably formed by alkylation of a surface carbenium ion by an alkene molecule. Because very different molecular transformations are involved in the monomolecular mechanism than the other two mechanisms, it is reasonable to expect that the products and the kinetics of these reactions would be different also [21,23]. Therefore, it should be possible to use product distributions in different experiments to indicate whether the data in the different experiments were obtained for the same mechanism, and if active sites of different chemical properties are present on different catalysts. Detailed product distribution information is available in a number of literature reports. Interestingly, these results show that the product distributions is a strong function of conversion of alkane, but only a weak function of the activity of the zeolite, i.e., whether it has been steamed or the extent of deactivation. Table 1 shows some illustrative data for the cracking of 2methylpropane (isobutane) [26]. In this experiment, a steamed USY zeolite was poisoned to different extents by Na ions. The cracking activities of the resulting samples were very different, as indicated by the different flow rates (F) and weights of catalyst (W) needed to obtain the same 2-methyolpropane conversion. In spite of their very different activities, the product distributions were practically identical when compared at the same 2-methylpropane conversion. This is unexpected, because, presumably, Na ions selectively poison the unusually active sites, as was indeed shown by the large decrease in activity when a proportionally small amount of Na was introduced. Therefore, the identical product distributions imply that if there were unusually active sites present, the chemical properties of these sites are identical to the sites with normal activities. A similar conclusion was obtained in the cracking of hexane. It was observed that over three different samples (a steam-treated Y, a steam-treated Y followed by extraction of A1 ions, and a chemically dealuminated Y), the rates of formation of C~ and C2 products relative to the rate of formation of 2.methylpropane product followed the same dependence on the rate of hexane cracking [17], independent of whether the data were collected after 5 minutes time-onstream or twenty four hours. Another very similar type of observation was made in the cracking of 2.methylpentane [22]. The molar ratio of C3 hydrocarbons to the sum of 2-methylpropane and C5 hydrocarbons in the products followed a common trend as a function of conversion for three different zeolites (a H-USY sample, a chemically dealuminated Y, and an HY) and for data

26 Table 1 Product distribution (molar concentrations relative to n-C4H10) in isobutane cracking at 430~ and 2.33% conversion on a steamed USY zeolite containing different Na contents. From ref. 26. Na20 (wt. %) Na/A1 F/W (10-6 mol/g-s)

2.03 0.17 0.12

0.45 0.04 1.0

0.02 0.002 8.2

Rel. conc.

0.25 0.33 0 0 0.42 0.17 1 0.25 0.25 0.33

0.23 0.30 0 0.03 0.39 0.17 1 0.21 0.22 0.32

0.22 0.29 0 0.03 0.39 0.16 1 0.21 0.23 0.32

H2 CH 4

C2H6 C2H4 C3H 8

C3H 6

n-C4Hlo i.C4H 8 n.C4H 8

i-C5H12

collected between I and 110 minutes time-on-stream. The implication of these observations is that the chemical properties of the active sites in these samples are very similar, in spite of the fact that their apparent activities may differ by orders of magnitude.

3. EXPERIMENTAL Detailed description of the experimental procedure and samples used can be found in published reports [27,28,29]. Briefly, the four zeolites used: HZSM-5, H-USY, CDHY, and HMOR were commercial samples. The n-hexane cracking activity was determined in a quartz tubular flow microreactor loaded with 0.01 to 0.07 g of 50/80 mesh zeolite pellets, supported by acid-washed quartz wool. n-Hexane (Aldrich 99% purity) was purified to reduce the level of alkene contaminant, by thoroughly mixing equal volumes of hexane and pure sulfuric acid, followed by passing the organic phase over a column of y-alumina calcined to 450~ or slurrying the acid-treated hexane with calcined LZY-62 zeolite. The resulting alkene concentration in the hexane was about 390ppm, compared to 3400 ppm in the unpurified hexane. The reactant stream, containing hexane at partial pressures ranging from 87 to 380 Pa, was prepared by passing a stream of N 2 through a saturator containing hexane at 0~ This stream was then mixed with a stream of pure N 2 of an appropriate flow rate to obtain the desired partial pressure. The reaction products were analyzed by on-line gas chromatography. Reactions were carried out between 480 and 540~ and the flow rate and catalyst weight were adjusted to yield conversions below 30%. The temperature, flow rate, and catalyst weight were changed in a random order in order to detect if the catalyst underwent deactivation. In all cases, a

27 reproducible, steady-state activity was obtained within two minutes, which was the time needed for the system to stabilize. In the absence of zeolite, the conversion of hexane was less than 0.1%. Cracking of 2-methylpentane was carried out at 300~ in the same apparatus using a feed containing about 2 kPa partial pressure of the alkane. In these experiments, 0.02 to 0.5 g of zeolite was mixed with s-alumina. Deactivation of the catalysts occurred. The initial conversion was determined by extrapolating to zero time using an empirical equation that fit the observed conversions measured at several values of time-on-stream.

4. RESULTS AND DISCUSSION

4.1 Kinetics and product distributions in hexane cracking Under the conditions employed, n-hexane cracking followed a first order kinetics up to the highest hexane conversion measured, which was about 30%. Thus, ln(1-X) was a linear function of W/F, where X is hexane conversion [27]. Fig. 1 shows the data for HZSM-5 and HUSY at 540~ Over the four zeolites tested, there was no detectable deactivation during testing over a period of three days (15 hours). The product distribution changed some with hexane conversion. The primary product distributions were obtained by extrapolating the product selectivities to zero hexane conversion. These are shown in Table 2. The data in Table 2 show that, within experimental uncertainties, the product distributions at zero hexane conversion were identical for the four zeolites. They are also very similar to those reported for HZSM-5 [30]. The similar product distributions strongly suggest that under these conditions, the same reaction mechanism prevails over these zeolites. According to the proposed mechanism for monomolecular cracking, the reaction proceeds via the formation of transition state I. Cleavage of the bonds associated with the five-coordinated carbon atom

Table 2. Molar selectivities at zero conversion of products (excluding dihydrogen) in n-hexane cracking at 500~ From reference 27. Product

HZSM-5

CH 4 C2H 6 C2H 4 C3H 8 C3H 6

5 16 6 14 36 0 2 20 0.5

i.C4H10 n.CaH10 C4H8a i-CsHl2

a) All isomers.

H-MOR 6 15 4 12 38 0 2 19 4

H-USY 8 14 5 13 38 0 4 16 3

CDHY 8 14 7 13 40 0 3 14 1

28 would lead to the formation of the primary cracking reaction products. Ideally, the ratios of H2/C6=, C~/C5, C2/C4, and C3/C3= products should be unity. However, secondary reactions, especially those of hexenes, pentenes, and other larger alkenes, would lead to higher concentrations of propene and ethene, but not to other light products, such as ethane and methane. The product distributions shown in Table 2 follow this general trend. As the hexane conversion increases, the product distributions changes, suggesting increasing extent of secondary reactions. The selectivities for ethene, propene, and isobutane increase, whereas those for ethane, propane, and butenes decrease. In view of the changing product selectivity, it is surprising that first order kinetics was observed up to 30% hexane conversion. This must imply that either these secondary reactions contribute little to hexane conversion, or the rate constants for hexane conversion by these secondary reactions are about the same as for the monomolecular reaction. The first explanation is more likely, because the increase in the selectivity for propene and ethene with conversion can be explained by the secondary reactions of the products of the primary reaction, as explained earlier. The concentration of isobutane was very low, although increasing with conversion. Thus, the product distributions and the first order kinetics strongly support the assumption that the cracking reaction in these experiments proceeds by the monomolecular mechanism. The bimolecular and oligomeric cracking reaction would contribute to hexane conversion. As can be seen in equation 1, the extent of bimolecular reaction depends on the surface coverage of carbenium ion, which is formed by adsorption of alkene on a Br~nsted acid site. Thus, it would depend on the alkene partial pressure. The variations of alkene concentration and selectivity as a function of hexane conversion are shown in Fig. 2. The data show that although the alkene concentration increased with hexane conversion, the selectivity remained quite constant. The latter suggests that the dominant mechanism did not change in these experiments. Thus, the contribution of bimolecular reactions to the kinetics of hexane cracking must be quite constant and insignificant. 4.2 Activation Energies The apparent activation energies for hexane cracking and the turnover frequencies (TOF) for both hexane and 2.methylpentane cracking are shown in Table 3. For this table, the TOF's were calculated using the number of framework A1 ions (corrected for the small amount of Na ions). It has been shown that using other methods to measure the acid site concentrations, such as based on sites that adsorb NH3 with over 90 or 120 kJ/mol yield similar qualitative trends [27,28]. The TOF's for the two Y zeolites, one steamed and one not steamed, were within a factor of two for hexane cracking, but the difference was much larger for 2-methylpentane cracking. Thus, the activities of the Bronsted acid sites in these two catalysts were quite similar for monomolecular cracking, but were apparently quite different under conditions when other mechanisms prevail. It has been shown that for the monomolecular cracking mechanism, the first order kinetics can be described by the reaction sequence that involves an adsorption preequilibrium step of the alkane (Eq. 2). The adsorbed alkane (P-s(ad)) is transformed to the product via the transition state I in the rate limiting step (Eq. 3). Ps(g) # Ps (ad) Ps(ad) ~ products

(2) (3)

29

)

120 100 ~ a . 80

5,

o.i5

"0

0 5 1015 a0254045 Fig. 1: First order kinetics plots for hexane cracking at 540~ for HZSM-5 (#) and H-USY (m).

=g

60

=.

4o 0

mm _ 9

...,;...

o"

'I

100 ~~

L.

..

10 20 I-bxane (xmv. %

30

Fig. 2: Olefin partial pressure ( . ) and selectivity (I) in hexane cracking over H-USY at 500~

Under the condition when the surface coverage of ~(ad) is small, the rate of reaction can be expressed as: Rate = kKPm

(4)

where k is the intrinsic rate constant (i.e. for step 3), K is the adsorption equilibium constant for step 2, and P is the partial pressure. Then, the observed activation energy, Eob., is related to the intrinsic activation energy, Ein, and the heat of adsorption, AH.d, by:

Eapp= AHaa + E,nt

(5)

Table 3 shows the intrinsic activation energies calculated using Eq. 5. Within experimental uncertainties, the intrinsic activation energies are the same for all the samples studied. That is, the differences in the observed activation energies can be attributed entirely to the differences in the heats of adsorption of the hexane.

4.3. Possible role of pore diffusion and effect of steaming The data in Table 3 show that, for monomolecular cracking, there is little difference between the active sites in H-USY and CDHY. The very similar activation energies and TOF's strongly suggest that the chemical properties of the active sites in these two catalysts must be very similar. However, under other conditions, primarily when other mechanisms prevail, as indicated by the fact that the catalyst deactivates, the steamed sample (H-USY) is much more active. It should be noted that the difference at the lower temperature is not due to temperature effect. It was observed in the same experiments at 673 K that under monomolecular cracking conditions, the TOF's for CDHY and H-USY remained within a factor of two.

30 This phenomenon can be explained by a reaction model described in reference 22. Briefly, the alkane cracking reaction is initiated by the monomolecular mechanism, which, at the lower temperatures and higher hydrocarbon partial pressures, has a smaller rate constant than the rate constants for the bimolecular and oligomeric cracking reactions. The monomolecular cracking reaction leads to the formation of alkene products. Adsorption of alkene on the Bronsted acid sites generates adsorbed carbenium ions and initiates bimolecular and oligomeric cracking reactions. The contributions of these two cracking reactions increase with increasing alkane conversion. Fig. 3 shows schematically the contributions of the various mechanisms to the total reaction rate as a function of the position in a plug flow reactor, which is closely related to alkane conversion.

Table 3. Activation energies and turnover frequencies for hexane and 2-methylpentane cracking. Sample

2.MePentane

HZSM-5

573 K TOF (10 -3 s"1) -

H.USY a

Hexane

Hexane tobs

AHad~

32

149+_8

(kJ/mol) -86+_6

E int (kJ/mol) 234+_14

4.2 c

3.3

177 +_9

-50+_3

227 +_12

CDHY b

0.5 c

2.3

186+_9

-50+_3

236+_12

H-MOR

-

15

157 +_9

-69+_3

226+_12

773 K TOF

( 1 0 -4 S"l

kPa "1)

(kJ/mol)

a. Y zeolite treated with high temperature steam. b. Y zeolite dealuminated with ammonium hexafluorosilicate. c. Initial rate data, taken from ref. 28.

The apparent enhanced catalytic activity by steaming can be explained as follows. The bimolecular and oligomeric cracking reactions, for which the rate constants are much larger than for the monomolecular reaction, are the dominant reactions under conditions when the large enhancement in catalytic activity by steaming is observed. Because of their large rate constants, these reactions become pore diffusion limited beyond a certain partial pressure of alkenes, which could be quite low at low temperatures, and could be achieved at relatively low alkane conversions. Steaming of the zeolite generates cracks, fissures, mesopores and other defects. These lead to a substantially larger effective external surface area for diffusion. Thus, on a per unit weight basis, the rate constants for the diffusion-limited bimolecular and oligomeric cracking reactions are much enhanced, which results in the apparently much more active catalyst. This is schematically illustrated in Fig. 3. It has been shown that a three times increase in the external surface area could result in up to 27 times increase in the rate of oligomeric cracking reaction

[22].

Some data in the literature support the possibility of pore diffusion limitation for the bimolecular and oligomeric cracking. Some of these are shown in Table 4. Haag and coworkers

31 determined the effective factor for reaction in HZSM-5 [31]. Diffusion influence on the reaction rate was observed for the rapid cracking of alkenes, which had an intrinsic rate constant about 300 times the slower monomolecular cracking of alkane. Dumesic and coworker have estimated the rate constants for reactions in the cracking of 2-methylpropane [23] and 2-methylhexane [32]. From their calculations, the ratios of the rate constants for the bimolecular/oligomeric cracking to monomolecular cracking can be estimated to be very large (Table 4). These results strongly suggest that it is quite feasible that the bimolecular and oligomeric cracking reactions are strongly pore-diffusion limited.

~lO0B

d

~ 10A 0...,

~ 1

f f

I

/

o.1-

I

Bed Length(a.u.)

t 1

Fig. 3: Relative contributions of various reaction mechanisms to the overall rate of cracking as a function of position in a plug flow reactor. Monomolecular mechanism ( ............), bimolecular and oligomeric cracking ( m . . . ) , and overall reaction (---). Case A for zeolite without steaming, and case B for steamed zeolite.

Table 4. Rate constant and pore diffusion limitation in cracking reactions.

Crystal radius, btm hexane hexene 3.methylpent-2-ene 2-methylpropane and 2.methylhexane

Intrinsic k, Mol/s-cm 3 cat.

Effectiveness factor, r I

29 7530 7420

0.025 1 1 1

B/M~= 10 4-8 (573 K) 10 3-6 (773

1.35 1 0.86 0.50

Ref. Haag [31]

Dumesic [23,32 ]

K)

a Ratio of apparent rate constants for bimolecular and oligomeric cracking to monomolecular cracking.

32 5. CONCLUSION Results are presented which strongly suggest that steaming of a zeolite does not necessarily result in the generation of active sites of unusual chemical properties or catalytic activities. The enhancement in cracking activity due to steaming depends on the reaction conditions, and can be explained by the destruction of the zeolite particles upon steaming that facilitates the pore diffusion-limited bimolecular and oligomeric cracking reactions.

6. ACKOWLEDGEMENT Support of this work by the National Science Foundation, Chemical and Thermal Systems Program is gratefully acknowledged.

7. REFERENCE

1 W.O. Haag, R.M. Lago, and P.B. Weisz, Nature 309 (1984) 589. 2 L.H. Lunsford, in: Fluid Catalytic Cracking II, ed. M.L. Occelli, American Chemical Society, Washington, D.C. 1991, p. 1. 3 R.A. Beyerlein, G.B. McVicker, L.N. Yacullo, and J.J. Ziemiak, J. Phys. Chem. 92 (1988) 1967. 4 F. Lonyi, and J.H. Lunsford, J. Catal. 136 (1992) 566. 5 R.M. Lago, W.O. Haag, R.J. Mikovsky, D.H. Olson, S.D. Hellring, K.D. Schmitt and G.T. Kerr, in New Developments in Zeolite Science and Technology: Proc. 7th Intern. Zeolite Conference, Tokyo, Marakomi and Iijima eds. Kodansha Ltd., 1986, p. 677. 6 D. Barthomeuf and R. Beaumont, J. Catal. 119 (1973) 288. 7 P.O. Fritz and J.H. Lunsford, J. Catal. 118 (1989) 85. 8 E.A. Lombardo, G.A. Sill, and W.K. Hall, J. Catal. 119 (1989) 426. 9 R.A. Beyerlein, G.B. McVicker, L.N. Yacullo and J.J. Ziemiak, Prep. Div. Petrol Chem. Amer. Chem. Soc. 31 (1986) 190. ~0 R.A. Beyerlein, C. Choi-Feng, J.B. Hall, B.J. Higgins, and G.J. Rays, Topics in Catalysis 4 (1997) 27. ~ T.F. Narbeshuber, A. Brait, K. Seshan, and J.A. Lercher, J. Catal. 172 (1997) 127. ~2 R. Carvajal, P.-J. Chu, and J.H. Lunsford, J. Catal. 125 (1990) 123. ~3 P.V. Shertukde, W.K. Hall, J.-M. Dereppe, and G. Marcelin, J. Catal. 139 (1993) 468. ~4 y. Hong, V. Gruver, and J.J. Fripiat, J. Catal. 150 (1994) 421. ~5 Q.L. Wang, G. Giannetto, and M. Guisnet, J. Catal. 130 (1991) 471. 16 A , C o r I ~ a , V~ Fomes, F.A. Mocholi, J.B. Monton, and F. Rey, in FCC II, ACS Symposium Series 452, M. Occelli ed., American Chemical Society, Washington, D.C., 1991, p. 12. 17 N.P. Rhodes, R. Rudham, and N.H.J. Stanbridge, J. Chem. Soc., Faraday trans. 92 (1996) 2817. 18 J.T. Miller, P.D. Hopkins, B.L. Meyers, G.J. Ray, R.T. Roginski, G.w. Zajac, and N.H. Rosenbaum, J. Catal. 38 (1992) 115. t9 G.R. Bamwenda, Y.X. Zhao, W.A. Groten, and B.W. Wojciechowski, J. Catal. 157 (1995) 209.

33

z0 P.V. Shertukde, G. Marcelin, G.A. Sill, and W.K. Hall, J. Catal. 136 (1992) 446. 21 W.O. Haag, and R.M. Dessau, Proc. 8th Intern. Cong. Catal., Verlag Chemie, Weinheim, 2 (1984) 305. zz B.A. Williams, S.M. Babitz, J.T. Miller, R.Q. Snurr, and H.H. Kung, Appl. Catal. A: General, in press. 23 G. Yaluris, J.E. Rekoske, L.M. Aparicio, R.J. Madon, and J.A. Dumesic, J. Catal. 153 (1995) 65. 24 S.R. Blaszkowski, M.A.C. Nascimento, and R.A. van Santen, J. Phys. Chem. 100 (1996) 3463. 25 V.B. Kazanski, and I.W. Senchenya, J. Catal. 119 (1989) 108. 26 j. Engelhard, and W.K. Hall, J. Catal. 125 (1990) 472. 27 S.M. Babitz, B.A. Williams, J.T. Miller, R.Q. Snurr, W.O. Haag, and H.H. Kung, Appl. Catal. A: General, accepted. 28 M.A. Kuehne, S.M. Babitz, H.H. Kung, and J.T. Miller, Appl. Catal. A: General, 166 (1998) 293. 29 M.A. Kuehne, H.H. Kung, and J.T. Miler, J. Catal., 171 (1997) 293. 30 T.F. Narbeshuber, H. Vinek, and J.A. Lercher, J. Catal. 157 (1995) 388. 31 W.O. Haag, R.M. Lago, and P.B. Weisz, Faraday Disc. 72 (1982) 317. 32 G. Yaluris, R.J. Madon, and J.A. Dumeisc, J. Catal. 165 (1997) 205.

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Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

35

Internal Reforming in Solid Oxide Fuel Cells R.M. Ormerod Birchall Centre for Inorganic Chemistry and Materials Science, Department of Chemistry, Keele University, Staffordshire, ST5 5BG, United Kingdom

Abstract

A test system based around a thin-walled extruded solid electrolyte tubular reactor has been developed, which enables the fuel reforming catalysis and surface chemistry occurring within solid oxide fuel cells to be studied under genuine operating conditions. It permits simultaneous monitoring of the catalytic chemistry and the fuel cell performance, allowing a direct correlation to be made between the cell output and the reforming characteristics of the anode. Using this system nickel-based/zirconia cermet anodes have been studied. The influence of the anode composition and formulation, pre-treatment method, operating temperature and methane/steam ratio on the methane reforming activity and the nature and level of carbon deposition have been investigated. Doping the anode with small quantities of molybdenum can lead to a very substantial reduction in the amount of carbon deposited. Temperature programmed oxidation has revealed that at least three types of carbon are formed on the anodes during high temperature methane reforming. Each of these species has been found to form at different rates. As current is drawn from the cell increased methane conversion occurs together with reduced carbon deposition.

1. INTRODUCTION Fuel cells are currently attracting a great deal of interest because of their tremendous potential as a more efficient and cleaner alternative method of electricity generation than heat engines, steam and gas turbines and combined cycles, since they operate electrochemically, so are not limited by the Carnot cycle, and they are catalytic with much lower emissions of NOx, in particular, and also CO2 and unreacted hydrocarbons. Solid oxide fuel cells (SOFCs) offer potential advantages in terms of efficiency and cost over other fuel cell types, largely because the high operating temperatures allow the possibility of running the cell directly on natural gas or other hydrocarbon fuels, internally reforming the fuel within the fuel cell [1-3], since at these temperatures methane and higher hydrocarbons react readily with steam or oxygen. This also leads to increased operational efficiency through utilisation of waste heat. It is generally accepted that in order for SOFCs to ever be cost effective, internal reforming of natural gas is essential, since internal reforming both increases efficiency through chemical recuperation of waste heat from the stack into the fuel supply and simplifies the balance of plant. However, several major operational problems of internal reforming remain to be solved before such cells can ever be routinely operated on hydrocarbon fuel, in

36 particular the problem of carbon deposition on the anode at the high operating temperatures involved, which leads to deactivation and a loss of cell performance and poor durability, sintering, leaching and delamination of the anode particles, the strongly endothermic nature of steam reforming and the problems of start-up and operation at low power. Other problems include obtaining optimum, durable anode formulations and the design of a suitable test system. Consequently, the majority of SOFC studies to date have been carried out using hydrogen as a fuel, and there are a lack of studies using hydrocarbons as fuels, particularly ones in which the catalysis is studied in a realistic SOFC device. The SOFC anode is generally nickel-based and can be considered to be analogous, in some respects, to supported nickel methane steam reforming catalysts, where the formation of carbon deposits on the catalysts continues to attract much interest [4-8], although the preparation method, metal content, support material and pre-treatment procedure are rather different. Although the addition of steam to the fuel is beneficial for the removal of carbon deposits, the use of large quantities of steam is generally undesirable in SOFC systems. The rate of carbon formation on nickel-based catalysts can also be reduced by the introduction of additives to the catalysts. Traditionally alkali metals, generally potassium in the form of potash, are added which serve both to accelerate the reaction of carbon with steam and to neutralise the acidity of the catalyst support, and hence reduce carbon deposition [9,10]. It is therefore possible to study the catalytic behaviour of nickel-based anodes in powder form inside a conventional catalytic reactor; this approach has been demonstrated by several research groups [2,11,12]. In many other fuel cell studies the primary focus is the electrical performance of the cell, and the influence of electrode composition, structure and other experimental parameters, are monitored by measuring current densities [13-15]. There are a genuine lack of studies in which the fuel catalysis and surface chemistry has been studied in an actual SOFC. We have developed an SOFC test system based on a small diameter, thinwalled, extruded yttria-stabilised zirconia tubular reactor which can be used to study the fuel processing catalysis in the SOFC, the chemistry occurring at the anode surface and the electrical performance of the fuel cell [16,17]. This allows a direct correlation between the cell performance and the internal reforming characteristics of the cell. In addition the test system can be readily used to study the problems of carbon deposition and poor durability in operation, as well as to develop optimised anode formulations [18,19]. The particular benefits of this test cell are that it can be rapidly assembled, heated, tested and cooled, and it has no sealing or leakage problems, which many test devices suffer from. In this paper we describe how this system has been used to study direct internal reforming in SOFCs, and in particular to obtain detailed information about the methane activation process, methane steam reforming, the nature and level of carbon deposition on the anode and the kinetics of carbon removal. The apparatus also allows the catalytic and surface chemistry of a working SOFC to be monitored enabling direct correlation of the reforming activity and surface chemistry with the cell performance [16,17]. The influence of anode composition, preparation route, reduction treatment, operating temperature, the effect of adding steam to the methane, and the effect of doping the anode with small quantities of additives, on the methane reforming activity, the surface chemistry and the carbon deposition process, as well as on the cell performance and durability, have all been investigated.

37 2. E X P E R I M E N T A L 2.1 SOFC test system All the experiments described here were carried out using the SOFC test system developed in our laboratory [16,17,20]. Briefly the apparatus consists of a microfurnace operated by a temperature controller which allows linear temperature control up to 1373 K. The test cell inlet is linked to a stainless-steel gas manifold which allows complete flexibility in gas handling, gas composition, the choice of fuel and fuel/steam ratio. The gas feed can be instantly switched between gases using a 4-way sampling valve. Thus evaluation is possible over a wide range of operating conditions and fuel compositions. The reactor outlet is linked via a heated gas sampling system to a continuously sampling on-line mass spectrometer (Leda-Mass Satellite) which permits the fuel processing reactions to be studied and the surface chemistry occurring at the anode to be investigated using temperature programmed spectroscopy. A particular advantage of the tubular SOFC design is that it can be housed in the furnace and used in the same way as a conventional stainless-steel or quartz catalytic reactor. As yttria-stabilised zirconia is a good thermal insulator, the ends of the electrolyte tube which project beyond the outer walls of the furnace remain sufficiently cool for a gas tight seal to be made, even when the inside of the furnace is at temperatures as high as 1373 K. The system has been designed so that either conventional stainless-steel or quartz reactors, or an extruded yttria-stabilised zirconia reactor can be used. 2.2 Anode preparation The SOFC anodes used in this work were prepared by physically mixing nickel oxide (Alfa Chemicals) with 8 mol% yttria-stabilised zirconia (YSZ) (Unitec-FYT11). A mixture of methanol, 1,1,1-trichloroethane and glycerol trioleate was added as a solvent and the resultant slurry was milled for three hours, with a small quantity of poly-vinyl butyral added at the end of the milling period as a binding agent. The anode sample can then be studied in the powder form, following firing, in a conventional reactor; an approach which has been used previously by ourselves and other workers [2,11,12]. However, in this case the anode slurry is coated onto the inside of the zirconia electrolyte tube prior to firing, as in an actual SOFC. Following drying at room temperature in air, the coated zirconia tubes were fired in a static air oven to 1573 K (heated to 773 K at 1 K min-1, from 773 K to 1573 K at 5 K min -1, held at 1573 K for 1 hour, and cooled to room temperature at 10 K min-1). Strontium-doped lanthanum manganite (Seattle Speciality Chemicals) was used as the cathode, and was applied to the outside of the zirconia electrolyte as an ink using a methanol/1,1,1-trichloroethane mixture as the solvent. Following drying in air at room temperature, the cells were fired to 1573 K using the same procedure that was used for the anode. Current/voltage measurements were carried out using a specially designed passive potentiostat. Various anode samples were prepared for detailed study: 50 vol% NiO/zirconia, 90 vol% NiO/zirconia and MoO3 doped 50 vol% and 90 vol% NiO/zirconia. These are referred to as 50/50, 90/10 and Mo-doped 50/50 and 90/10 Ni/zirconia anodes, respectively. The molybdenum doped samples were prepared by adding MoO3 (BDH) to the anode ink prior to milling and firing. Following firing the anodes were reduced in the reactor at 1173 K for 30 minutes in a 10% H2/He stream.

38

2.3

Catalytic experiments

Dry and steam reforming reactions were carried out by passing the fuel mixtures over the reduced anode at reaction temperature. Temperature programmed measurements were carried out using a heating rate of 10 K min -1. Temperature programmed reduction (TPR) and temperature programmed oxidation (TPO) measurements were carried out in 10% H2/He and 10% O2/He mixtures, respectively.

3. RESULTS 3.1 Temperature Programmed Reduction

TPR was used to study the reduction characteristics of the anode samples. Table 1 summarises the results for the reduction of the 50/50, 90/10 and 1% Modoped 50/50 NiO/zirconia anodes. The reduction of a simple 1:5 physical mixture of NiO and zirconia and a 10% NiO/zirconia sample prepared by wet impregnation of an aqueous solution of nickel (11) nitrate (Fluka puriss) were also studied by TPR. Table 1 Reduction temperatures of NiO/zirconia samples determined by TPR Sample 50/50 NiO/YSZ anode 90/10 NiO/YSZ anode 1% MoO3 doped 50/50 NiO/YSZ anode 1:5 NiO/YSZ physical mixture 10% NiO/YSZ (wet impregnation)

3.2 Methane Adsorption

Temperature of Peak Maxima / K 758 707 760 673 815, 894

Methane activation and decomposition was studied by carrying out temperature programmed measurements on the reduced anodes in dry methane/helium and methane/steam/He gas mixtures. Figure 1 shows a temperature programmed methane adsorption (TPMA) spectrum for the 50/50 Ni/zirconia anode, and indicates that dissociative adsorption of methane starts to occur at temperatures above 780 K. The rate of methane adsorption was found to increase with temperature, although a local maximum in the adsorption/decomposition rate was observed at 975 K. Dissociative adsorption of methane leads to the formation of surface carbon species. If the anode sample was subsequently subjected to a temperature ramp in hydrogen (temperature programmed hydrogenation (TPH)) evolution of methane was observed. However, methane was still evolved at the highest reaction temperature of 1173 K, indicating that not all the surface carbon can be removed by hydrogenation. If the TPH experiment was followed by a temperature programmed oxidation (TPO) experiment, significant high temperature evolution of CO2 was observed, confirming that a substantial quantity of carbon does remain on the anode following hydrogenation. Two maxima were observed in the rate of carbon removal indicating the presence of two distinct types of carbon species. No water

39 was produced in the TPO experiment indicating that all the adsorbed methane had completely dissociated to form purely carbonaceous species.

o N

m m

C

hydrogen

o m

300

400

500

,

,

I

600

700

800

~

i

900

Temperature I K

!

1000 1100 1200

Figure 1. Temperature Programmed Reaction spectrum of the reduced 50/50 Ni/zirconia anode carried out in dry methane.

3.3 Methane Reforming Methane was passed over the 50/50 and 90/10 Ni/zirconia anodes at different reaction temperatures to study the influence of reaction temperature and to compare the stability towards carbon formation of the two anodes. Following each methane exposure the surface carbon formed was analysed by TPO. As the reaction temperature is increased the amount of carbon deposited increased and the temperature at which the maximum rate of removal of deposited carbon by oxidation also increased, indicating that the deposited carbon formed is more strongly bound at higher methane reaction temperatures. For a given reaction temperature the TPO maxima occurred at similar temperatures for both anodes. However, the quantity of carbon adsorbed on the two anodes was different; at each temperature less carbon was formed on the 90/10 Ni/zirconia anode than on the 50/50 anode. The effect of adding steam to the methane feed has been studied over a wide range of methane/steam ratios. Figure 2 shows the exit gas compositions when a reduced 50/50 Ni/zirconia anode was exposed to a 5:1 methane/steam mixture at 1123 K. Upon introduction of the CH4/H20 mixture to the anode a large transient uptake of methane was observed together with hydrogen and CO evolution. The methane conversion fell off to a steady-state conversion of around 35% after about five minutes. Initially some CO2 was also produced though this quickly fell off to a very low level, with almost total selectivity towards CO formation. The effect of adding steam to the methane feed on the nature and quantity of carbon deposited during reforming at different reaction temperatures was evaluated using TPO. These experiments demonstrate that the addition of only small amounts of steam lead to a significant reduction in the level of carbon

40 deposition, and that this effect is most marked at the highest reaction temperatures. Figure 3 shows the two TPO spectra obtained following exposure of the 50/50 Ni/zirconia anode to dry methane and a 19:1 CH4/H20 mixture, respectively, at 823 K for thirty minutes. Even at this low reaction temperature a significant reduction in the amount of carbon deposited can be seen.

:5

..E 01

~3

0

10

20

30

40

50

60

~'0

Time I min Figure 2. Exit gas compositions following exposure of a reduced 50/50 Ni/zirconia anode to a 5"1 methane/steam gas mixture at 1123 K.

_

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=,--,,

_

,..-.

l

r 13")

methane/stea~ -~

u

300

400

500

600

700

800

900 1000 1100 1200

Temperature / K

Figure 3. TPO of the 50/50 Ni/zirconia anode following exposure to dry methane and a 191 methane/steam mixture at 823 K for 30 minutes.

4! We have studied the effect of doping the Ni/zirconia anodes with small quantities of additives, including molybdenum, lithium [12], potassium, gold and iron. The influence of molybdenum has been studied in considerable detail at different reaction temperatures and CH4/H20 ratios, over a range of molybdenum Ioadings and firing temperatures. The promotional effect of the molybdenum is complex, depending on the reaction conditions, Mo loading and firing temperature, and will be described in detail elsewhere [21]. In brief, molybdenum doping can lead to a very substantial reduction in the amount of carbon deposition. Figure 4 demonstrates this for the 90/10 Ni/zirconia anode doped with, nominally, 1 wt% Mo, following reforming in a 5:1 CH4/H20 mixture at 1123 K for 60 minutes. It can be seen that under these conditions essentially no carbon deposition occurred on the Mo doped anode, whereas significant carbon deposition occurred under the same conditions on the undoped 90/10 anode. Doping the anode with molybdenum did not result in any significant change in the methane reforming activity compared to the undoped anode under the same operating conditions. The electrical performance of tubular SOFCs coated with these anodes has been studied and show that Mo doping does not have any detrimental effect on the power output [22]. ,-~

90/10 Ni/YSZ

z= ~3

1

300

-9 I 500

I,

I

700

900

,

i 1100

1300

Temperature I K

Figure 4. TPO of 1 wt% Mo-doped 90/10 and undoped 90/10 Ni/zirconia anodes, following exposure to a 5:1 methane/steam mixture at 1123 K for 60 minutes. The effect of increased reforming time on the nature and level of carbon deposited on the anode was studied using TPO. Figure 5 shows the TPO profiles obtained following reforming of a 19:1 methane/steam mixture at 973 K over a 50/50 Ni/zirconia anode for different lengths of time. As expected the quantity of carbon deposited on the anode increases with increasing reforming time. It can also be seen that at least three types of carbon species are formed on the anode. The carbon species are referred to as; type I carbon, removed from the anode at about 870 K, type II carbon, removed just above 900 K, and type III carbon which can only be removed from the anode by oxygen at temperatures above 1000 K. As the

42 reforming time increased the relative quantities of these species altered. After 30 minutes reforming it was found that the most strongly bound carbon species (type III) was the main species formed. As the exposure time was increased to 60 minutes it was found that type I carbon became the most dominant species. Further reforming time resulted in an increase in the quantity of both type I and type II carbon, with the quantity of type II carbon eventually exceeding that of type ! carbon. The quantity of type III carbon saturated at low exposure times.

120 min

.$

i .

300

400

500

600

7OO 800

9O0 1000 1100 1200 1300

Temperature / K Figure 5. TPO of a 50/50 Ni/zirconia anode following increasing exposure to a 19:1 methane/steam mixture at 973 K.

3.4 Combined Catalytic and Electrochemical Measurements A particularly powerful feature of the SOFC test system is the capability to simultaneously monitor the catalytic chemistry occurring at the fuel reforming anode and the electrochemical performance of the SOFC under actual operating conditions, enabling direct correlation of changes in the reforming activity and surface chemistry of the anode with the fuel cell performance. Figure 6 shows the effect of drawing current from an SOFC with a 50/50 Ni/zirconia anode operating at 1123 K in a 19:1 methane/steam mixture on the reforming reaction. It can clearly be seen that as the current drawn increased there is an increase in the methane conversion, and stepwise increased production of hydrogen, and significantly increased formation of CO which parallels the increase in H2 production. In addition to increased CO and H2 production, the formation of C2 species, ethene and ethane, was also observed, the level of which show a similar increase as the current drawn increases.

43

N

03

!

o1

0.5V

L.

0.6V

E 0

\

O O1 01 o1 03

-CO

~~.,-,~'~

'

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5

- -~""H2

~

10

'

i ......

15

i

20

Time / min

'

'1

25

1

30

'

35

Figure 6. Effect of drawing current on the reforming activity of a tubular SOFC with a 50/50 Ni/zirconia anode running on a 19:1 methane/steam mixture at 1123 K. 4. D I S C U S S I O N The reduction characteristics of the anodes are intermediate between that of the physical mixture of NiO and zirconia, where reduction of the nickel oxide occurs at rather lower temperature (673 K), and those of a sample prepared by wet impregnation of the nickel component, where reduction of the nickel oxide component does not occur until significantly higher temperatures (815 K and 894 K). The detailed interpretation of these results will be reported elsewhere. Dissociative adsorption of methane starts to occur on the anode at about 780 K. The absence of any water evolution in the TPO experiments following reforming indicates that all the methane had completely decomposed to form a purely carbonaceous overlayer. The local maximum in the methane adsorption observed at 975 K is most likely due to a rate-limiting step in the process. This could be due to carbon species adsorbed on the nickel spilling over onto the support, surface carbon migrating into the nickel bulk, or the formation of carbon filaments. It is known that on nickel catalysts formation of carbon filaments and the formation of nickel carbide can both occur [5,23]. The strength of carbon adsorption following methane dissociation is demonstrated by the fact that not all the carbon can be removed from the anode by hydrogenation, even at 1173 K. The amount of carbon deposited on the anode increased with increasing reforming temperature, together with the temperature at which the maximum rate of carbon removal occurs, indicating that the carbon becomes more strongly bound with increasing reaction temperature. Our results also show that the anode composition, preparation method and pre-treatment procedure all have a considerable influence on the level of carbon deposition [21]. The addition of only a small quantity of steam to the methane feed leads to a substantial lowering in the rate of carbon formation on the anode. The methane/steam ratios used in our work are considerably higher than those typically used by other workers [2], though still demonstrate the role of steam in reducing

44 carbon deposition. The addition of steam to the methane feed alters the reactions taking place on the anode. Methane steam reforming (1) occurs in addition to methane dissociation (2). The CO produced by steam reforming can then undergo further reaction via the Water Gas Shift Reaction (3) or the Boudouard reaction (4). The steam can also potentially remove deposited carbon by reaction (5). CH4 + H20 -)' CO + 3H2

(1)

CH4 ~

(2)

Cads + 2H2

CO + H20 -~ 002 + H2

(3)

2CO --) Cads + 002

(4)

Cads + H20 -> CO + H2

(5)

Our experiments show that the Ni/zirconia anodes are active methane steam reforming catalysts, which quickly reach steady-state conversion levels. Very high selectivity towards syngas formation is observed, suggesting that both the Water Gas Shift Reaction and the Boudouard reaction are only minor reaction pathways under these methane rich conditions. The addition of small quantities of molybdenum to the Ni/zirconia anode can lead to a very substantial reduction in the level of carbon deposition compared to the undoped anode under the same reaction conditions. Doping the anodes with molybdenum does not lead to any significant change in the methane reforming activity of the anode or the cell performance compared with the undoped anode under the same reaction conditions. These findings are consistent with previous work on conventional nickel steam reforming catalysts, which show that addition of molybdenum reduced carbon deposition without resulting in significant loss of catalytic activity [24]. TPO has shown that for all the anodes studied more than one type of carbon species is formed during high temperature methane reforming. The length of the exposure time to methane significantly influences both the quantity of carbon adsorbed on the anode surface and the nature of the carbon species. These results indicate that each carbon type is either formed on different sites on the anode or that one form of carbon can be converted into another, either by higher temperature or by longer exposure. Very short exposure times result in the most strongly bound (type III) carbon being formed. The quantity of type III carbon formed saturates at a relatively low level. It seems most likely that this strongly bound carbon species is graphitic carbon formed on the surface nickel sites. With increasing methane exposure formation of the lower temperature type I carbon occurs, such that type I carbon becomes the dominant carbon species. At longer exposure times, type II carbon becomes more significant, until eventually type II carbon is the major species, though the quantity of type I carbon also continues to increase. Our results suggest that type II carbon is formed from the type I species, possibly via some sort of ageing process. The addition of molybdenum seems to inhibit the formation of type I and type II carbon but not type III carbon, suggesting that molybdenum may inhibit the transition from type III to type I and type II carbon. The total quantity of carbon deposited after prolonged exposure to methane can exceed the quantity of nickel present. A large proportion of the carbon must therefore spill over onto the zirconia or be in the form of filamentous carbon.

45 Electron microscopy indicates that the anodes suffer considerable damage when they are run under harsh coking conditions in comparison to anodes which have been run under less severe conditions, indicating that the carbon species formed actually break up the anode structure, suggesting that carbon filaments are formed. Our results on working SOFCs show clearly that drawing current from the SOFC results in increased conversion of methane, and increased production of hydrogen and CO, formation of ethene and ethane, and also reduced carbon deposition, compared to the unloaded anode under the same reaction conditions. Electron microscopy has shown that the damage to the anode is not as severe when the cell is operated under load conditions. As the current drawn from the SOFC increases, i.e. as the cell potential decreases, the flux of oxygen ions through the yttriastabilised zirconia electrolyte, from the strontium-doped lanthanum manganite cathode to the Ni/zirconia anode increases. The increased conversion of methane and the stepwise increase in CO and H2 production, with increasing current drawn, indicate that partial oxidation of methane to CO and H2 by the oxygen ions is occurring. The formation of ethane and ethene indicates that in addition to partial oxidation, in the presence of oxygen ions transported through the solid electrolyte the Ni/zirconia anode is also active towards oxidative coupling of methane. The partial oxidation and oxidative coupling pathways which occur in the working SOFC lead not only to increased methane conversion but also reduced carbon deposition. 5. S U M M A R Y

In summary, we have developed a test system for studying solid oxide fuel cells, based on a thin-walled extruded yttria-stabilised zirconia reactor, which can be used to investigate the catalytic behaviour of the fuel reforming anode, the chemistry occurring at the anode surface, and the electrochemical performance of the fuel cell, under genuine operating conditions. Catalytic measurements can be made on a working SOFC, and temperature programmed measurements can be carried out on anodes in an actual SOFC. These have been used to investigate different anode formulations, to study methane activation and methane reforming, and evaluate the nature and level of carbon deposition on the anode during high temperature operation. The system therefore permits a direct correlation to be made between the fuel cell performance and the reforming characteristics of the anode. We have shown that the deposition of carbon on Ni/zirconia anodes in SOFCs running on methane is strongly influenced by anode composition, preparation method and pre-treatment procedure, operating temperature and methane/steam ratio. Doping the anodes with small quantities of molybdenum can lead to a very substantial reduction in the level of carbon deposition, whilst having little effect on the reforming activity or the cell performance. Temperature programmed oxidation indicates that at least three different types of carbon species are formed on the anode during high temperature methane reforming, differing in their strength of binding. The relative quantities of the different types of carbon depend on the operating temperature, methane/steam ratio, the anode formulation and the presence of molybdenum. When an external load is applied to the cell, increased methane conversion occurs together with reduced carbon deposition, through reaction of the methane

46 via partial oxidation and oxidative coupling, with the flux of oxygen ions through the yttria-stabilised zirconia electrolyte. 6. A C K N O W L E D G E M E N T S

This work was supported by the UK Engineering and Physical Sciences Research Council under grant GR/K58647. BG Technology are also acknowledged for financial support. 7. R E F E R E N C E S

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24

A.L. Dicks, J. Power Sources, 61 (1996) 113. E. Achenbach and E. Riensche, J. Power Sources, 52 (1994) 283. N.Q. Minh and T. Takahashi, in: Science and Technology of Ceramic Fuel Cells (Elsevier, Amsterdam, 1995) and references therein. S.C. Tsang, J.B. Claridge and M.L.H. Green, Catal. Today, 23 (1994) 3. D. Duprez, M.C. Demichelli, P. Marecot, J. Barbier, O. Ferretti and E.N. Ponzi, J. Catal., 124 (1990) 324. O. Yamazaki, K. Tomishige and K. Fujimoto, Appl. Catai. A:, 136 (1996) 49. V.R. Choudhary, B.S. Uphade and A.S. Mamman, Catal. Lett., 32 (1995) 387. D. Qin and J. Lapszewicz, Catal. Today, 21 (1994) 551. H. Praliaud, J.A. Dalmon, C. Mirodatos and G.A. Martin, J. Catal., 97 (1986) 344. J.R. Rostrup-Nielsen and L.J. Christiansen, Appl. Catal. A:, 126 (1995) 381. R.T. Baker and I.S. Metcalfe, Appl. Catal. A:, 126 (1995) 297. N.J. Coe, R.H. Cunningham and R.M. Ormerod, Catal. Lett., 49 (1997) 189. I.P. Kilbride, J. Power Sources, 61 (1996) 167. K. Honegger, E. Batawi, C. Sprecher and R. Diethelm, Proc. 5th Int. Symp. on SOFCs, The Electrochemical Soc., 1997, 321. K. Eguchi, H. Mitsuyasu, Y. Mishima, M. Ohtaki and H. Arai, Proc. 5th Int. Symp. on SOFCs, The Electrochemical Soc., 1997, 358. C.M. Finnerty, R.H. Cunningham, K. Kendall and R.M. Ormerod, J. Chem. Soc. Chem. Commun., (1998) 915. C.M. Finnerty, R.H. Cunningham and R.M. Ormerod, Proc. 3rd European SOFC Forum, 1998, 227. R.H. Cunningham, C.M. Finnerty and R.M. Ormerod, Proc. 5th Int. Symp. on SOFCs, The Electrochemical Soc., 1997, 973. C.M. Finnerty, N.J. Coe, R.H. Cunningham and R.M. Ormerod, Catal. Today, 1998, in press. R.H. Cunningham, C.M. Finnerty, K. Kendall and R.M. Ormerod, Proc. 5th Int. Symp. on SOFCs, The Electrochemical Soc., 1997, 965. C.M. Finnerty, R.H. Cunningham and R.M. Ormerod, in preparation. C.M. Finnerty, R.H. Cunningham and R.M. Ormerod, in preparation. C.H. Bartholomew, Catal. Rev. Sci. Eng., 24 (1982) 67. T. Borowiecki and A. Golebiowski, Catal. Lett., 25 (1994) 309.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

47

Microchemical reactors for heterogeneously catalyzed reactions Dieter H 6 n i c k e D i v i s i o n o f Industrial C h e m i s t r y , Technical U n i v e r s i t y o f C h e m n i t z , 09107 Chemnitz, Germany Abstract

The application of precision engineering techniques to chemical reaction engineering has been identified as having numerous potential advantages for both the chemical process development and manufacture of chemicals. Most developed are the microfabricated components which perform the standard unit operations, i. e. micropumps, microvalves, micro heat exchanger and separation units. In comparison, the ongoing development of chemical microreactors acquires more effort for the availability of effective and reliable reactors, especially for heterogeneously catalyzed reactions. It is the purpose of this paper to briefly describe the intellectual and experimental ways of down scaling from industrially used multitube fixed bed reactor to microreactors including some illustrated applications. From the preceding information, we conclude that the extremely rapid heat and mass transport what is available through the engineered microstructures yielding reactors with high selectivities and even reaction rates. The development of microreactors for heterogeneously catalyzed reactions at present are focused on providing sufficient amounts of catalytic active material on the walls of the microstructures. This was identified as a crucial prerequisite for an extended variety of possible applications and assembling microchemical systems including the required unit operations as well as sensors, analyzers and electronic units for process managing and controlling. When the desired chemical microreactors and corresponding systems are available, they provide opportunities for point-of-use production of chemicals, for the time being in small, high purity amounts and also of inflammable or hazardous intermediates with storage and transportation restrictions.

1. INTRODUCTION Microreaction technology represents an interesting new field of chemical and electrochemical reaction engineering, analysis as well as screening procedures, and offers novel facilities for distributed chemical manufacture. Chemical microreactors are the most important parts of microchemical systems. In this paper, microreactors for heterogeneously catalyzed reactions are the focus of attention. The number of publications in the last five years concerning microreaction technology including microsensors and microdevices for unit operations and analysis increased drastically. However, only few studies were published relating to chemical microreactors for heterogeneously catalyzed reactions. Most of them report that the progress in precision technology, particularly the silicon technology for electronics was the decisive assumption for the microreaction technology, which is indeed true. But looking back in history, there was another driving force discovered for the development of more effective reactors [1,2]. Today,

48 it seems it was forgotten, for, already in the nineteen thirties, the drawbacks of the industrially used multitube reactor technology were analyzed. This led to the designing of multitube reactors having catalytic active inner walls. Additional explanation is necessary: The previously mentioned drawbacks which result from the use of multitube reactors with fixed bed catalysts were mainly i) the time consuming catalyst fill-in-operation into the thousands of tubes each of 15 to 25 mm in diameter, ii) the differences in pressure drops, which gave, iii) differences in residence times, and iv) the limited heat removal. Therefore, to overcome these, the same multitube reactors were used without fixed bed catalyst, but with tubes having inner porous layers of catalytic active components. As a result, no catalyst filling was needed. Differences in pressure drops and residence times were diminished and better heat transfer properties were expected. These so-called multitube wall reactors were tested, e.g. in the methan formation starting from syngas [3] and in the partial oxidation of naphthalene to phthalic anhydride [4]. Summarizing the principal results, two features were stated: First, the degree of conversion obtained was very low leading to insufficient product yield which was caused by a low catalyst surface area per reactor volume and inadequate radial mass and heat transfer properties. Therefore, a small space-time yield was achieved. Even the same multitube wall reactor, however, filled with catalytically inactive particles for enhanced mass and heat transfer characteristics did not result in significant higher product yield. Secondly, the catalytic active layer, irregular in thickness and catalyst distribution had a deficient adhesion strength leading to catalyst loss during the catalytic reaction. Therefore, the respective reactor development was terminated and the pursuance of the desired aim diminished into oblivion as earlier mentioned. Today, sixty years later, we are able to build reactors, viz. chemical microreactors, having high surface to volume ratios and catalytic active porous layers with high adhesion strengths. This advancement became possible by adapting and using precision techniques, surface treatment methods and deposition technologies which now possess highly developed standards.

2. R E Q U I R E M E N T S ON C H E M I C A L M I C R O R E A C T O R S Multitube fixed bed reactors actually used for heterogeneously catalyzed reactions are characterized by surface area to volume ratios of about 106 to 5 " 108 mZ/m3 (Fig. 1 - region a) dependent on specific catalyst surface areas ranging from one to several hundreds mZ/g [5]. The corresponding data for multitube wall reactors are 102 mZ/m3 or less (Fig. 1 - region c) which mirrored the low conversion degree of the feed as previously described. In considering the principle of catalytic active walls, one requirement from that arises to improve the efficiency, namely a dramatic decrease of the tubes diameter along with a tremendous increase of the number of tubes in order to achieve the region of 106 to 108 mZ/m3. This leads eventually to tube diameters in nanometer size which consequently produce a very high pressure drop during flow which would be unsuitable for practical applications. Thus, ten to several hundreds of micrometers for the tube diameter would be useful with regard to acceptable pressure drops. These measures are attainable today by numerous available precision techniques. However, not microtubes but microstructured elements having micro flow channels with equivalent diameters between 10 and 500 gm (Fig. 1 - region b) were used as described later. Consequently, about 103 to 105 m2/m 3 (Fig. 1- region b) were achieved which was calculated from the nonporous geometrical surface area of the

49 microchannels. In comparison, the human lung comprises roughly 210 4 m2/m3. Nevertheless, compared with multitube fixed bed reactors, there is still a gap of one to three order of magnitude. From that, a further requirement developed, viz. the necessity to increase the microchannel surface areas by particular treatments of or modifications to the surface. For that microchannel

~

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---

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e q u i v a l e n t diameter of the tubes and flow channels [m]

---F-

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/

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Figure 1. Surface area to volume ratio vs. equivalent diameter of the tubes and flow channels; multitube fixed bed reactors: region a: filled with fixed bed catalysts, region c: used as catalytic wall reactor; region b: microchannel reactors purpose different methods, namely electrodeposition of metals, physical and chemical deposition (PVD, CVD, grafting) of metals and metal oxides, and anodic oxidation to yield metal oxides from bulk metal were successfully used. By application of these methods, a surface area increase took place, and a simultaneous immobilization of catalytic active components was attained. As an additional requirement, the amount of catalytic active components, as well as the adhesion strength of the layer or the particles formed, have to be long term stable and highly sufficient.

3. DESIGN OF M I C R O R E A C T O R S AND M I C R O C H E M I C A L SYSTEMS Taking into account the previously described requirements, the following three technological steps are needed to build microreactors which can be completed to microchemical systems, which include specific microsensors and analyzers, micro devices for unit operations and microelectronics for process control: 1. Microstructuring of wafers to form microchannels can be done by either mechanical handling (e.g. milling, stamping, moulding, casting) or chemical and physical treatment (e.g. etching, LIGA, laser, plasma processing). A variety of wafer materials is applicable as metals, special alloys, silicon, polymers, glass and ceramics. The sizes and shapes of the microchannel's longitudinal axis and cross sections are precisely uniform in one wafer, but can also be intentionally varied from wafer to wafer. 2. Formation of catalytic active channel surfaces by using area-enlarged methods (e.g. anodic oxidation, laser treatment, plasma etching), followed by immobilization of catalytic active

50 components (e.g. impregnation, grafting) or by both of them in one step (e.g. particle or layer deposition by surface controlled reactions, sol-gel procedures or electrolytic treatment). 3. Completion of the microreactors as well as microchemical systems is accomplished by stacking the wafers and connecting them tightly by welding or gluing with specific adhesives. Eventually, a reactor cage and several adapters are needed for connecting the micro components with the macroworld to provide them with feed and energy, and also to get the formed chemical product and measured data.

Figure 2. Schema of a self-supporting microchannel reactor The state of the art of the precision technology for the three technological steps enables the production of numerous replicas, which are necessary for higher product quantities and decentralized manufacture of chemicals. Currently, two different types of microchemical devices have been developed. The first is a self-supporting microchannel reactor which is characterized by a stack of uniform microstructured and catalytic activated wafers. This is appropriate only to a definite catalyzed reaction (Fig. 2). The other type of a microchemical device is identified by a stack of several microstructured wafers useful utilizing different processes, like chemical reaction, mixing or heating with integrated pumps, valves and sensors for temperature, pressure and analysis which finally result in a microchemical system such as a chemical microplant (Fig. 3). Among the first who developed the previously described, self-supporting micro reactor, according to the first technological step, were researchers of the Forschungszentrum Karlsruhe in Germany. They applied diamonds as milling tools to form microchannels in metallic wafers [6]. By using their design, which was originally earmarked for micro heat exchangers [7], the microreactor can be easily produced in three versions: The first is a reactor which is run through in one direction using all flow-channels for the heterogeneously catalyzed reaction (Fig. 2). The heat transported from or to the channels is accomplished by using the outer reactor surface via the entire metallic matrix. The second is a reactor in which the stack is ordered in such a way that a crossflow regime is possible (Fig. 4), i.e. only the half number of the microstructured wafers is used for the reactands and the other for the heat transport medium. Fig. 4 also shows a photograph of this early developed microchannel reactor without connecting screws. This reactor is predominantly applicable for highly exothermic reactions. The third reactor version utilizes specifically designed in- and out-lets for the fluids whereby a counter-current of the reactands and the cooling or heating medium is achieved [8]. The arrangement of the channel system corresponds to that shown in Fig. 2. But, as in the second

51 version, only half the microstructured wafers are usable for the catalytic reaction, however, the counter-current results in more efficient heat

Figure 4. Photograph and scheme of a microchannel cross flow reactor Figure 3. Scheme of a microchemical system

transfer.

Meanwhile, a variety of reactors for heterogeneously catalyzed reactions were proposed, some were already produced, and their use in catalytic reactions is in progress [9a].

4. P O T E N T I A L APPLICATIONS OF CHEMICAL M I C R O R E A C T O R S Years ago the discussion of using tubes or flow channels for chemical reactions having very small diameters was declined due to the argument of prevalent laminar flow leading to insufficient radial mass and heat transfer characteristics. That is definitely true for diameters in mm-size which are used in monoliths, e.g. to treat waste gases for reasons of environmental protection. As a result, considerable long monoliths are necessary to attain high conversion degrees. However, with a decrease of the tube diameters to gm-size, the diffusion time of molecules between the narrow walls will be very short despite the laminar flow. For, the diffusion time is proportional to the square of the tube diameter. Therefore, mass and heat transfer properties are highly improved due to the short diffusion pathways. As a result, microreactors are valuable performing chemical reactions with high reaction enthalpies, especially exothermic reactions. The majority of them are fast and some are high temperature reactions which require short reaction times. The precise fabrication of the flow channels in diameter and length leads to equal residence times in all channels. These combined features allow a performance of chemical reactions under isothermal conditions which consequentially imply a uniform residence time distribution of the molecules due to the lack of thermally

52 induced irregular gas volume expansion. Therefore, higher selectivity is expected particularly for the formation of intermediates from partial reactions like oxidations and hydrogenations. Although the idea, viz. using catalytic active inner walls instead of fixed bed catalysts, was derived from conventional, industrial used multitube fixed bed reactors as described at the beginning of this paper, there is no real possibility of manufacturing microreactors which have the same product quantities. Also the frequently proposed method of "numbering up" instead of "scaling up", offered no alternative way in this respect for the replacement of industrial used fixed bed reactors by microreactors for basic and bulk products. Nevertheless, the potentials of chemical microreactors in particular for heterogeneously catalyzed reactions open new fields of applications and might be useful as follows: 1. Continuous production of fine chemicals, e.g. pharmaceuticals in moderate or low output of products. 2. Safe and reliable performance of high pressure and high temperature reactions due to the compact and encased construction of the microreactors. 3. Safe feasibility of reactions by using flammable and potentially explosive mixtures because of the small channel dimensions leading to ignition blocking. 4. Conversion and formation of hazardous and toxic substances on decentralized locations where they will be processed further into harmless final products. 5. Manufacturing of short-lived pharmaceutical drugs at the place of their immediate consumption. 6. Chemical processing within movable equipment like cars, and in manportable systems which is possible because of the compact and lightweight microreactor design. 7. Accelerated basic research for discovering novel synthesis routes, catalysts compositions and the influences of steady state or unsteady state conditions on the reaction kinetics and product formation. By reason of both the negligible amount of chemicals and materials needed for the investigations, as well as the extreme rapid heat and mass transport due to the low thermal mass of microreactors and fast thermal response times, microreactors offer opportunities for improved and fast research. 8. Catalyzed reactions for analytical purposes usable in chemical sensors for chemistry, medicine, biology, and chemistry-related industries. Besides, recent novel principles and methods are imaginable which will be the result of the progress in the development and the variety of chemical microreactors. It is envisioned that the chemical microreactors can be used to provide point-of-use and on-demand production to improve safety by eliminating storage and transportation of toxic and hazardous chemicals and reducing the potential danger due to accidents.

5. ILLUSTRATIVE EXAMPLES OF MICROREACTORS FOR HETEROGENEOUSLY CATALYZED REACTIONS The miniaturization of chemical reactors and apparatus for unit operations has faced increasing attention for nearly one decade. In particular, the latter experienced a dynamic expansion which is reflected by widespread activities and an increasing number of practical applications. Examples are miniaturized pumps, mixers, heat exchangers, and valves. On the contrary, only a few examples of tested and proven microreactors are published.

53 The results achieved in microreactors performing non-catalyzed exothermic reactions are very impressive. Examples are the gas phase phosgenation of cyclohexyl amine to yield cyclohexylisocyanate [9b], a liquid-liquid cyclization reaction using concentrated sulfuric acid [ 10a], and acylation reactions of amines with acyl- and sulphonylhalides and isocyanates as well as hydrogenation reactions [10b,11 ]. Static mixers and heat exchangers were integrated into the used microreactors. Both, higher conversion degrees and product selectivities were achieved, caused by a very fast and effective mixing and precise temperature control. The

Figure 5. Photograph of a microchannel crossflow reactor with connecting screws [6] attained results mirror the expected properties of micro reactors having flow channels in gmsize, because of their very short diffusion and thermal response times. These observations are decisive prerequisites in reaching sufficient results by performance of heterogeneously catalyzed reactions in microreactors. As earlier mentioned, one of the first developed microstructured heat exchangers [6] made of copper was used as microreactor (Fig. 5) after in situ-oxidizing the channel surfaces by oxygen [9c,12]. As a result of this process, metallic Cu ~ was mostly converted to Cu + Cu2O CH2=CH --CH 3

02 CuO

CH2=CH-CHO CO2 + H20

(Cu20) which is suitable for the partial oxidation of propene to acrolein, while Cu 2+ (CuO)

catalyzes the total oxidation: Copper-I-oxide was the first fixed bed catalyst in the corresponding industrial process for 3 acrolein production. The 1 cm -Cu-mmroreactor had channel cross sections of 80 x 80 ~m 2 and lengths of 14 mm. The oxidation results, viz. very low conversion degree and slight selectivity obtained under unsteady state conditions, gave the facts which have to follow for improving the properties of the microchannel reactor: 1. Drastic increase of the catalytic active surface area of the microchannel walls

54 2. Stabilization of the copperoxide phase composition which altered between Cu + and C u 2+ during the course of the reaction. Another, but more successful and very convincing example of heterogeneously catalyzed partial oxidation in a microreactor is the oxidehydrogenation of methanol and homologous primary alcohols to the corresponding aldehydes [9d,10a,13]" Ag CH3OH + I~2 0 2 ~ CH2=O + H20 Ag R-CH2OH + 1/2 02 .~ R-CH=O + H20 Microstructured silverfoils having channels with rectangular cross sections of 320 x 400 2 3 ~tm and lengths of 10 mm were used and crosswise arranged to a 1 cm-crossflow microreactor as depicted in Fig. 6. Furthermore, the figure also shows earlier investigated industrial-like reactor types and their corresponding results in the conversion of R-CH2OH to R-CH2=O. The first series of experiments gave the selectivities vs. conversions as shown in Fig. 7. From these results, we can conclude the more consistent the channel system is, plus more faster and effective the heat transfer is the higher the product yields. This includes lowering the reaction and hot spot temperature. The catalyzed partial oxidation of methane to form syngas was also investigated in a microreactor [10c,10d]: CH4 + ~/2 02

Rh

~ CO + 2 H 2

However, the 1500 x 254 ~tm2-channels were filled with a powdered rhodium fixed bed catalyst which gave some problems concerning high flow rates. Nevertheless, the authors conclude from their preliminary results: The use of catalysts in microchannel reactors is so promising that they have started developing engineered microstructures. This includes investigations with catalytic coatings and the development of ultrahigh surface area catalytic material.

55 An additional example of a heterogeneously catalyzed reaction carried out in a Simicroreactor is the oxidation of ammonia producing nitric oxide as a model reaction

Figure 6. Temperature data, conversions and selectivities in the dehydrogenation of a-alcohol to aldehyde using reactors having different configurations [13] 98 96 ,-.-.,

94

>, 92 >

'~

m0 3 (D

90

co 88 86 84

48

i

|

I

I

i

50

52

54

56

58

60

Conversion [%]

Figure 7. Selectivity vs. conversion in the dehydrogenation of (x-alcohol to aldehyde using a multishort tubular reactor (,,) and a microchannel reactor ( ~ [13] [ 10c, 10e]" NH3 + O2

Pt

~

NO / N20 / N2 + I420

The microreactor consisted of a 15 x 25 mm 2 silicon wafer in which a single T-shaped channel with a cross section of 0,55 x 1,3 mm 2 for gas flow was etched. The channel was capped by a SiN-Al-plate, and its inner SiN-site was deposited with thin-film-platinum as

56 catalyst. The plate carried the three gas inlet-outlet connectors. The T-shaped channel allows both radical mixing and oxidizing the reactands. In addition, the ammonia oxidation was also performed in a SiN-microreactor and a conventional reactor. Both had lower heat conductivity properties than the Si-microreactor. By the use of the Si-microreactor, results indicate no ignition phase and also a very localized reaction zone, while the channel walls and the bulk of the silicon wafer remain at room temperature. But, the other reactors reveal an ignition phase. These observations led to the conclusion that microreactors can potentially operate safer than conventional reactors. A similar investigation was carried out using a crossflow microreactor made of stainless

H 2 + !/2 0 2

Pt

~

H2

steel (Fig. 4) for the oxidation of hydrogen to water [10f]: The purpose of this study was to show the very safe performance of the exothermic water formation which was indeed successful with 100 % conversion of hydrogen to water. Platinum, as a catalyst, was immobilized by the incipient wetness method on an alumina layer which was deposited before on the channel walls by chemical vapor deposition. The cross section areas of the channels were 100 x 200 pm 2 for the reactands and 100 x 70 gm 2 for nitrogen as a coolant. It was shown that the microreactor is a safe and reasonable tool for converting and controlling combustible and potentially explosive gas mixtures. In addition to oxidation reactions, the hydrogenation is another widely used reaction type in which generally only moderate reaction heat results. In order to determine the potential of microreactors referring to the favored formation of a partially hydrogenated intermediate, formed under kinetically controlled conditions, cis, trans, trans-l,5,9-cyclododecatriene was chosen as the model compound for the hydrogenation [ 14]:

CDT CDD CDE CDA The goal of this study was developing a suitable microreactor and attaining conditions which ensure high selectivities to cyclododecene (CDE) with almost complete conversion of cyclododecatriene (CDT) and cyclododecadiene (CDD). Several beneficial steps were carried out as follows: First, different fixed bed catalysts were prepared for preliminary screening hydrogenation experiments. For this purpose aluminum wire was anodically oxidized in order to form a surface alumina layer with appropriate regular pore systems, followed by impregnation with a palladium containing precursor, as well as drying, calcining and reducing in hydrogen. The formed regular pore systems with pore diameters in m - s i z e and pore lengths in lam-size increased the geometrical surface area of the micro channels basically by one to three order of magnitude. This is sufficient to overcome the gap in the surface area to volume ratio as already described in the second chapter. Several preparation conditions were applied leading to catalyst samples with pore systems which differ in their pore diameters, pore lengths and pore densities as well as in palladium contents and distributions.

57 Second, the prepared wire was cut into extrudate-like particles which was then used as fixed bed catalysts for preliminary hydrogenation experiments. Third, as a result of that catalyst screening, the preparation conditions of the best catalysts were chosen for the corresponding treatment of microstructured wafers made of aluminium. The dimensions of the channel cross sections and lengths based on a suitable range of gas volume rates were pinpointed to 200 x 200 gm 2 and 30 mm, respectively. Fig. 8 depicted the anodically formed pore system (Fig. 8a), part of the wafer cross section containing the microchannels and the nano-pore system (Fig. 8b) as well as a stack of the anodized and catalytic activated wafers (Fig. 8c) [15]. The latter was caged in a stainless steel housing equipped with diffusors and connectors for educt inlet and product outlet, finally resulting in a

Figure 9. Microchannel reactor with caged wafer stack, diffusors and connecting screws [5]

58 Fourth, as the last step, the hydrogenation of cyclododecatriene was carried out using the lnicrochannel reactors and for comparison also reactors with fixed bed catalysts, like the best one made of aluminum wire, pieces of microstructured wafers (same as in the best microreactor) as well as an industrially used but not optimized Pd hydrogenation catalyst. Fig. 10 shows the achieved yields of cyclododecene vs. conversion of CDT and CDD. In conformity with the expectation the hydrogenation carried out on one of the microchannel reactors gave the highest yield to cyclododecene of 90 % at a conversion of 98 %, whereas the yields on the fixed bed catalysts were considerably lower. The excellent product yield attained with the microchannel reactor on the basis of these comparable results is caused by the lack of temperature and velocity gradients over the whole passage through the microchannels. These emerge mainly from the gm-dimensions of the uniform flow channels.

Figure 10. Yields of cyclododecene in the partial hydrogenation of cyclododecatriene vs. overall conversion of CDT and CDD on a microchannel reactor and reactors with different fixed bed catalysts; 393 K, Ptotai 110 kPa, PCDT= 110 Pa, PH2 = 330 Pa [5,10g] =

6. EXAMPLES OF ONGOING RESEARCH The public activities, workshops, national and international conferences generally held in the United States of America and Germany have shown a dynamic expansion of and growing interest on microreaction technology which today is widespread throughout continents. Most of the ongoing research projects are supported by funding programs, sponsored for instance in the USA by the Department of Energy, the Automotive Manufacturers, the Ministry of Defense and NASA [1 Oh] in England by the British Nuclear Fuels Ltd. and CRL laboratories and in Germany by the Federal Ministry of Science and Technology and Ministry of Economy. The German Society for Chemical Apparatus, Chemical Engineering and Biotechnology (DECHEMA) piloted the corresponding activities which initiated to meetings and funding programs [10i]. Table 1 shows a list of some current research projects dealing with microreactors for heterogeneously catalyzed reactions. The examples represent a wide field of chemical reaction engineering for production, screening, energy supply and pollution control.

59

Project 1. Partial oxidations of n-butenes to maleic anydride and ethene to ethylene oxide in lnicroreactors 2. Periodic processing in microreactors used in the deamination of amines and alcohol dehydration 3. Microreactors for fuel processing/fuel cell systems which include reactors for partial oxidation to produce syngas, shift reaction to convert CO to hydrogen and preferential CO oxidation for diminishing the CO level to less than 10 ppm. 4. Parallel preparation and testing of catalysts in microreactor arrays having different degrees of miniaturization. The realization will be demonstrated in oxidation an hydrogenation reactions. 5. Heterogeneously catalyzed liquid phase hydrogenation of organics to partially hydrogenated products in microchannel reactors

Project goal Supported by Determination of the extreme limit GE regarding partial and total pressures Demonstrating the potential advantages, viz. higher selectivities and production rates than in stationary processing Realizing the concept of hydrogenbased new vehicle generation via fuel cells as well as distributed power generation units

Demonstration of rapid catalyst preparation and testing as a high throughput technique to find novel catalysts.

GS

UE UA

GS

Estimation of the technical and GSM economical potentials of microreactors for continuous liquid phase reactions related to those attained in batch reactors GE: German Federal Ministry of Economy; GS: German Federal Ministry of Science and Technology; UE: U.S. Department of Energy; UA: U.S. Automotive Manufacturers, GSM: Saxon Ministry of Science and Art (Germany)

Table 1. Examples of some current research projects dealing with heterogeneously catalyzed reactions and funding programs [ 10i].

microreactors

for

60 7.

APPROACHES FOR MINIATURIZEI) CHEMICAL PLANTS

For decades, several approaches have been made concerning down-scaled chemical plants especially for portable purposes and decentralized manufacture. However, some new proposals dealing with concepts of miniaturized chemical plants in the last five years, appeared during a period of increased progress in microtechnology and microreaction engineering [16]. Examples for miniaturized plants refer to distributed HCN- and ethanolproduction, either as continuous, batch or semi batch processes [17]. Unfortunately, these proposals do not shown the order of dimination down to microsize. However, the author illustrates and describes convincingly that the model for the respective approach is the way in which nature works. Nature provides many examples of sophisticated chemical manufacture carried out in equipment that, by the standards of the chemical industry, would be regarded highly unlikely and unsafe. Organisms and animals can be thought of as chemical plants for manufacturing a highly complex product, namely more animals. These "plants" have a number of characteristics and operate typically batchwise. They have a limited life time and are composed of disposable and recyclable materials. These materials are invariably nonmetallic and flexible, even their instrumentation. In general, they do not require internal cleaning, repair or maintenance. From this model we receive a strong impact to reflect all our targets and courses for developing such sophisticated microchemical plants. That includes refocus on reaction technology, simplify separation, avoid recycle where possible, devise robust equipment, which will avoid maintenance; furthermore, application of high tech devices and materials as well as methods, e.g. cooling units using semiconductor Peltier-effect, high resistant polymers, even those with limited life time, radio frequency heating, gravity as driving force.

8. FUTURE TARGETS In the near future it is anticipated that the assembly of chemical processing and energy conversion systems ranging in size from smaller than one cubic centimeter to several cubic meters will be available. Most of them include chemical reactors for heterogeneously catalyzed reactions. However, the extent of industrial used applications will be increased with the variety of successful proven catalytic reactions in microreactors and microchemical systems. At that point, the today's somewhat perceivable retention of the industry will become more interested and involved in microreaction technology. One of a very impressive future target is the NASA's plan to install compact chemical processing plants on the surface of the Mars [1 Oh]. The interest for that is diminishing the costs associated with robotic (starting 2001) and human missions (starting 2011) to Mars by reducing the required launch mass from Earth. The plan as part of the Mars program includes lightweight, chemical processing systems able to produce the needed chemicals and propellants from indigenous space materials. The basic idea is to use stored hydrogen from Earth and atmospheric CO2 from Mars for producing propellants and oxygen. A microtechnology-based system could consists of adsorption units (for collecting CO2), electrochemical units (for CO2 reduction to 02 and CO) and catalytic microreactors (for CH3OH formation from CO and H2). Several other reactions are also of interest, e.g. Fischer-

61 Tropsch-syntheses (for higher hydrocarbons), reverse-water-gas shift (to form water and CO) an Sabatier reaction (to form C H 4 a s heating gas). Finally, another proposal was announced, viz. the development of distributed processing systems based on chemical microsystems for global carbon management which would lead to CO2 emission reductions [1 Oh]. Such systems may have significant advantages, especially if combined with modular fuel cells. The latter are not Carnot-cycle limited whereby the mentioned combination can be inherently more efficient than conventional combustion or steam power plants.

9. C O N C L U S I O N S The state of the art in precision technology including the semiconductor fabrication for electronic devices, in recent years, continues to give the best presupposition developing chemical microreactors and systems. The application of microfabrication concepts to microreactors has been identified as having a number of advantages. Since the majority of industrial used chemical reactions are catalytic, the respective reactors in particular for heterogeneously catalyzed reactions were the topic of this paper. In addition to the demands on microreactors for homogeneous reactions two important aspects has to fulfill for microreactors suited for heterogeneously catalyzed reactions, viz. high compactness, i.e. sufficient high surface area to volume ratio, and high adhesion strengths of catalytic active depositions formed as layers or particles on the surface of reactor microchannels. Both aspects are the focus of attention in the ongoing developments. The small number of really successful operating microreactors probably mirror the challenges of attaining the requirements. Nevertheless, the perceptible growing interest, the widespread activities and the increasing variety of proposals for practical applications demonstrate the progress in microreactor evolution particularly for heterogeneously catalyzed reactions. In closing, not only the development of microreactors, but also that of complete microchemical systems with the needed unit operations, integrated sensors, analyzers and electronics for control, promises to be one of the powerful supplies for diversified technologies in the near future. 10. R E F E R E N C E S

F. Paneth, K.F. Herzfeld, Z. Elektrochem. 37 (1931) 577 G. Damk6hler, Chem. Ing. 3 ( part 1) (1937) 402 M. M. Gilkeson, R. R. White, C. M. Sliepcevich, Ind. Eng. Chem. 45 (2) (1953) 460 T. G. Smith, J. J. Carberry, Chem. Eng. Sci. 30 (1975) 221 G. Wiel3meier, Doctorate Thesis, Universit~it Karlsruhe (TH), Karlsruhe, Germany 1996, Shaker Verlag Aachen 1997, ISBN 3-8265-2183-8 K. Schubert, W. Bier, G. Linder, D. Seidel, Chem. Ing. Tech. 61 (1989) 2, 172 W. Bier, W. Keller, G. Linder, K. Schubert, D. Seidel, D. H. Martin, Chem. Eng. Process. 32 (1993) 1, 33 K. Schubert, Forschungszentrum Karlsruhe, Germany, Hauptabteilung Versuchstechnik, personal communication DECHEMA-Monographien Vol. 132, Micro-System-Technology for Chemical and Biological Microreactors, Workshop, DECHEMA, Frankfurt/Main 1996, ISBN 3-527-

62

10

11 12 13 14 15 16 17

10226-4, a) H. L6we, W. Ehrfeld, K. Gebauer, K. Golbig, O. Hausner, V. Haferkamp, V. Hessel, Th. Richter, pp. 63-74, b) J. J. Lerou, M. P. Harold, J. Ryley, J. Ashmead, T. C. O'Brien, M. Johnson, J. Perotto, C. T. Blaisdell, T. A. Rensi, J. Nyquist, pp. 51-69, c) D. H6nicke, G. WieBmeier, pp, 93-107, d) K.-P. J~ickel, pp. 29-50, Proceedings of the 2 nd International Conference on Microreaction Technology, March 1998, New Orleans/LA, USA, a) O. W6rz, K.-P. J~ickel, Th. Richter, A. Wolf, pp. 183185, b) N. Schwesinger, O. Marufke, F. Qiao, R. Devant, H. Wurziger, p. 124, c) D. L. Brenchley, R. S. Wegeng, pp. 18-23, d) A. L. Y. Tonkovich, J. L. Zilka, M. R. Powell, C. J. Call, pp. 45-33, e) A. J. Franz, D. Quiram, R. Srinivasan, I.-M. Hsing, S. L. Firebaugh, K. F. Jensen, M. A. Schmidt, pp. 33-38, f) U. Hagendorf, M. Janicke, F. Schfith, K. Schubert, M. Fichtner, pp. 81-87, g) G. WieBmeier, D. H6nicke, pp. 24-32, h) R. S. Wegeng, M. K. Drost, pp. 3-9, i) J. P. Baselt, A. F6rster, J. Herrmann, D. Tiebes, pp. 13-17 N. Schwesinger, O. Marufke, F. Qiao, Technische Universit~it Ilmenau, Germany, R. Devant, and H. Wurziger, Merck KgaA Darmstadt, Germany, personal communication G. WieBmeier, Diploma Thesis, Universit~it Karlsruhe (TH), Karlsruhe, Germany, 1991. O. W6rz, K.-P. J~ickel, BASF AG Ludwigshafen, Germany, personal communication G. WieBmeier, D. H6nicke, Ind. Eng. Chem. Res. 35 (1996) 4412-4416 G. WieBmeier, D. H6nicke, J. Micromech. Microeng. 6 (1996) 285-289 R.S. Benson, J. W. Ponton, CHERD 71, A2, 160-168, 1993 J. W. Ponton, Microreaction Technology, Proceedings of the First International Conference on Microreaction Technology, Febr. 1997, Frankfurt/M., Germany, pp. 1019, Ed. E. Ehrfeld, Springer Verlag 1998, ISBN 3-540-63883-0

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Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

65

Novel frequency response techniques for the study o f kinetics in heterogeneous catalysis M. Cavers, J.M. Davidson~ , I.R. Harkness, G.S. McDougaU and L.V.C. Rees Department of Chemistry, The University of Edinburgh, King's Buildings, West Mains Road, ~ b u r g h , EH9 3JJ, Scotland, U.K., Department of Chemical Engineering, The University of Edinburgh, King's Buildings, Mayfield Road, Edinburgh EH9 3JL

Abstract Frequency response (FR) techniques are applied to determine the dynamics of gas surface interactions in 2 systems. By studying the response of propane adsorption on silicalite-1 it is demonstrated that FR is an accurate and rapid method for simultaneous determination of diffusivities and isotherms of microporous sorbents. Frequency response analysis of in-situ IR spectra of CO oxidation on Rh/A1203 shows that the predominant adsorbed species under reaction conditions, geminal dicarbonyl, plays no role in CO~ production and that the minority species CO on Rh8 § and carbonate are more important.

1. Introduction Frequency response methods are based on the manner in which a system, at equilibrium, responds to a periodic change in some external parameter influencing the equilibrium and, in particular, the dependence of this response on the frequency of perturbation. They are an established means of analysis in process systems [1], however, they are equally applicable to any system which can be described by a set of differential equations linear in the perturbation. In heterogeneous catalysis, although application of FR methods is as yet relatively rare, the potential of the technique has been demonstrated by a number of studies. Sinfelt has considered the advantages of FR in flow systems over either static chemisorption experiments or conventional differential flow reactors [2]. In experimental studies, most notably those by Yusada, complex kinetics have been resolved in the conversion of dimethylether over HZSM-5 [3]. Work on supported metal catalysts includes studies of alkali promotion of Rh/TiO2 by HE adsorption [4l and CO adsorption/resorption kinetics employing in situ transmission infrared (IR) analysis of the adsorbed phase [5]. Until recently, our own FR studies have largely concerned measurement of diffusion [6] and sorption [7] in zeolites in batch reactors, however, we now describe the development of a novel FR flow reactor system employing Mass Spectrometry (MS) and Diffuse Reflectance Infrared Fourier Transform Spectroscopy (DRIFTS). Results from a study of the adsorption and diffusion of propane in silicalite-1 are described first. This system has been well characterised both by FR [8] and by other methods [9] and so was chosen as a suitable trial system to test the MS detection, flow properties of the FR system

66 and modelling of flow reactor data. A combined MS/DRIFTS FR study of CO oxidation over a Rh/A1203 catalyst illustrates the ability to correlate the gas phase FR with the rates of elementary reaction steps on the catalyst surface.

2. Experimental The experimental apparatus is depicted in Figures 1 and 2 and consisted of a gas delivery system, a reactor, which took the form of either a simple tube reactor or a DRIFTS cell, and a mass spectrometer.

Figure 1. Schematic of the apparatus

Figure 2. Cross section of the DRIFTS cell

In the simplest scenario, three mass flow controllers (MFC's) supplied the sample, reference and diluent gases. A small modulation in the flow of the combined sample and reference gas stream was produced by driving the solenoid valve with a signal generator. This was then injected into a much larger constant flow of diluent gas. The resultant gas stream had a near constant volumetric flow rate but an oscillating composition at a total pressure of around 1 atmosphere. Appropriate settings of the signal generator allowed the small amplitude (< 10%) changes necessary to ensure that the reponse remains linear with respect to the perturbation. The maximum frequency usefully delivered was principally dependent on the flow characteristics downstream from the injection point but could be in the order of 3-5 Hz if care was taken to reduce tubing lengths and cross sections. The minimum frequency is determined by the long term stability of the flow rates and was of the order of 0.02 Hz. More than 2 orders of magnitude of frequency could therefore be investigated. The reference gas signal was used both as a phase and amplitude reference and results were calculated as an amplitude ratio and phase lag relative to this reference. Control experiments have shown that in the absence of any sample there is good agreement between the 2 signals throughout the frequency range studied.

67 The tube reactor used for the propane/silicalite-1 studies was simply constructed from a suitable length of 1~ inch outside diameter stainless steel tubing with -30mg of the zeolite sample held between glass wool plugs. Electron microscopy studies of the sample showed it to be made up of roughly spherical crystallites of approximately 10ktm radius. For the DRIFTS cell depicted in Figure 2, ~30mg of the finely powdered Rh catalyst was placed in a depression on a silica sample post. The use of powdered samples in DRIFTS reduces the potential for artefacts due to macropore diffusion associated with pressed discs required in transmission IR experiments. Gas enters the cell at the base and passes up a slot in the Macor TM gas guide next to the sample post and then flows across the sample and down a similar slot to exit. The total dead volume of the cell is minimal and, under the gas flow rates typically employed (60 sccm), gives residence times in the cell of the order of ~150ms, far better than commercially available cells. Sample heating was carded out by a heater coil inserted in the silica sample post. This could produce sample temperatures, measured by a thermocouple placed in the sample bed, in excess of 450~ without degradation of the seal. The low path length of the infra red radiation through the gas phase means no interfering gas phase signals are observed for species that adsorb (i.e. CO). Gas phase CO2 is however observed due to the absence of an adsorbed species in that region of the spectrum. The FTIR spectrometer used in association with the in house constructed DRIFTS cell was a Biorad FTS-6000 equipped with a narrow band Mercury Cadmium TeUuride detector. The mass spectrometer system employed with both the tube reactor and DRIFTS cell was a Leda Mass quadrupole fitted with a by-pass pumped fused silica capillary inlet system which could be inserted directly into the exit of the tube or DRIFTS reactors.

3. ADSORPTION AND DIFFUSION OF PROPANE IN SILICALITE-1

The adsorption and diffusion of propane in silicalite-1 was studied at 348K and at a range of propane concentrations. The average propane flow rate was varied from 0.6 sccm to 5 sccm while those for the argon reference gas and the helium diluent were 5 and 55 sccm, respectively. The propane partial pressure was thus in the range 7-56 Torr. The frequency of perturbation was varied from 0.02 Hz to 2 Hz. Typical time domain data from the mass spectrometer is shown in Figure 3. The influence of adsorption/desorption processes on the propane waveform is immediately evident. Whereas the argon retains it square wave character after passing through the sample bed, the propane wave becomes more rounded. The relative amplitude of the propane modulation is also reduced and there is a clear phase lag relative to the argon wave. Since any gas hoM-up effects will be present in the argon signal the amplitude attenuation and phase lag are characteristic of the dynamics of the sorption processes in the sample. Data such as those presented in Figure 3 are then Fourier transformed to yield an amplitude ratio and phase lag relative to the argon. The variation of these parameters with frequency of perturbation is shown in Figure 4, as is the dependence of this variation on propane partial pressure. In all case the amplitude ratio (argon/propane) increases with frequency and the phase lag

68

increases with frequency before passing through a maximum and dropping sharply. 1.8 q Argon 1.4 The increase of amplitude ratio with frequency 1.6-1 l (i.e. a preferential damping of the propane 1.4-1] 1,0 wave at high frequency) is in conflict with the 0.8~. 1.2 0.6 model for continuous flow frequency response 1.0 . . . . . 0.4 ~" published recently I101. Also in conflict with this model is the fact that the phase lag 0.8 o.0 O ~0 100 exceeds n/2. This is impossible if the sample "nine/s is considered to be a uniform discrete adsorber Figure 3. Partial pressure oscillations after since it would imply net adsorption occurring at times when the pressure of sorbent is below passage through the sample ,9

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Figure 4. Dependence of phase lag and amplitude ratio on frequency at various partial pressures T-- 348 K average. This is clearly non-physical. A more satisfactory treatment of the data results if the sample is considered a column of finite length of particles in which Fickian diffusion for an isotropic sphere is assumed. The gas flows through the sample as dispersed plug flow. Thus

-

t.JJ

=

a

where Dz = axial dispersion coefficient, C = propane gas concentration, q = adsorbed propane concentration, v = gas velocity, z = column coordinate, e = fractional void space, t = time. Taking Laplace transforms

. s = . Laplace operator, q - GC - D z 0z 2 ) + vr176 l,~z) + ~ s q ~ . = sC. where

(2)

69 G being the Transfer Function relating the gas and adsorbed concentrations. This has been derived previously by Do [ 10]. Therefore

(3)

-Dz 0z 2 ) + ~,-~z)+"e s G C = s C

This is an ordinary differential equation and thus can be solved. The perturbation is harmonic therefore the Laplace variable s can be replaced by Io3, yielding the transfer function relating the argon and propane partial pressures after the column

4'~176 ,)i)

= Ar C3H8

-Io3L}'~

e('v

(4)

Where the first exponential term is to allow for the transportation lag and

cot G = 9DxK 4

io) ,3

~ Dx

R-

(5)

/I;Io~R 2

where o3 = angular frequency of perturbation, L -- length of sample bed, R = size of particles, D r = micropore diffusivity and K = gradient of the isotherm at the equilibrium conditions. The amplitude ratio is equal to the magnitude of the complex transfer function whereas the phase lag is equal to the argument. The lines in Figure 4 show the theoretical best fits to the data, agreement between the model and the data is clearly good with the fit parameters presented in Table 1. Table 1 Parameters Propane Pressure (Torr) 7.7 29 56

from the fitted Transport Diffusivity (m2/s) 2.9 x 10-11 3.9 x 10-1~ 4.8 x l0 -11

curves in Figure 4 K Axial Dispersion 3800 1670 1130

1.3 x 10-3 1.7 X 1 0 -3 1.3 x 10.3

Void Fraction 0.89 0.95 0.93

Calculated Loading (mol/uc) 4.2 13 19

Corrected Diffusivity (m2/s) 2.5 x 10-~1 2.4 x 10-~1 2.1 x 10-11

As can be seen from Table 1 the calculated void fraction is of the order of 90%. This is consistent with the sample acting as a column despite the small amount of sample used. Were the sample to be packed perfectly it would be expected to be a few millimetres deep and thus much shorter than the wavelength of even the fastest perturbation used. Conditions

70 in the bed would then be uniform and the model of Do et al 1101 would be expected to be valid. With a void fraction of 0.9 the bed is an order of magnitude longer and thus acts as an extended column. The K values show the expected trend with partial pressure and were used to construct an isotherm shown in Figure 5 from the appropriate 8 o FittedIsotherm ~ ~ straight line segments. This isotherm is a good fit to a Langmuir isotherm and the calculated loadings are in reasonable agreement with traditional isotherm measurements [8]. These 12 loadings were used to calculate the corrected _@ io diffusivities according to the Darken equation (Dc=D(~lnq/~lnP). These values are in agreement with previous macroscopic 6 o7 measurements but are lower by 2-3 orders of magnitude than PFG NMR values I91 which are themselves in good agreement with the most r recent FR measurements under batch conditions [8]. The discrepancies between various measures Pressure/Ton" of diffusivities in zeolites is the subject of an ongoing discussion in the literature I911111 but it Figure 5. An isotherm based on values interesting to note that the values obtained here of K from the fitting of the FR data agree most closely with the values obtained in and a fitted Langmuir isotherm. Zero Length Chromatography, a technique in which the experimental setup is the most similar to that used here [9].

4. CO OXIDATION OVER Rh/AI203 The CO oxidation experiments a fourth MFC was used to blend a 1:1 mixture of CO and 02 which was then mixed with the He diluent as before. The tube reactor was replaced with the DRIFTS cell illustrated in Figure 2. The DRIFTS spectra were first used to follow the equilibration of the catalyst under the gas stream and Figure 6 shows the spectrum that results after exposure of the initially calcined catalyst to a flow of CO:O2:He (1"1:35 sccm) at one atmosphere and 489 K for 15 minutes. The spectrum is the ratio of a single beam spectrum prior to exposure of the sample to the gas stream against a second single beam spectrum recorded at equilibrium with the intensity of the peaks expressed arbitrarily in absorbance units. It shows gas phase CO2 (2363 and 2330 cm-~) indicating a steady state production of CO2, adsorbed CO (2087 and 2008 cm 1) and formation of carbonate species on the oxide (1645 and 1436 cm-~). The adsorbed CO is partly in the form of a geminal dicarbonyl species on the reduced Rh. CO on partially oxidised Rh (Rh~+) also contributes

71 to the intensity of the higher wavenumber side of the 2087 2087 cm"1 o.15

"~i/

; 2008 em "1 /

t

]6~5

0 O45 2363 em "1

.

•

11436 cm "1

i

0.030

/

§

cm

-1

i ~ - - 2109 cm "I ii ~ Jl

feature [121.

1436 em "1 ]645 orn 4

\

0.015

-~ 0.10

2i63

tJ

c m "1

0.05

2 1400

i

I 2200

i

0.000

?U I

2000

,

1

1800

i

,i

1600

i

-0 015 ,i

1400

Wavenumber / cm "l

Figure 6. The DRIFT spectrum after equilibration of the sample under the reactant gas stream at 489 K.

Wavenumber / cm "l

Figure 7. Representative spectra during modulation of the gas composition.

When modulation of the composition of the gas stream was introduced, the FTIR spectrometer was used to acquire DRIFTS spectra at a rate of 5 spectra per second for 4 minutes at each modulation frequency. The quality of the DRIFTS data was such that good spectral sensitivity could be retained at modulation frequencies up to 1.2 Hz. The spectra shown in Figure 7 are again absorbance spectra but now they are the ratio of a single beam spectrum of the sample at equilibrium under the CO/O2 stream against representative single beam spectra from one modulation cycle. The intensity of each of the spectral features can be quantified either on the basis of peak height or integrated area and plotted against time to generate a 'functional group chromatogram' (FGC) analogous to the MS time domain data shown in Figure 3. The combined MS and FGC time domain data for modulation at 0.21 Hz are shown in Figure 8. In this figure the relative amplitudes of each trace have been scaled to highlight differences in phase and the CO2 MS synchronised with the gas phase CO2 IR. At this point our modelling of the data is not as advanced s Spec. ' '~ as for the propane example discussed above, however, a number of important features are immediately evident from the combined MS/DRIFTS FR data. Firstly, .~ 12109cmTD/ "~ /" ! ~ I the carbonate regions of Figures 6 and 7 -1 J : are similar and show the proportion of the total surface carbonate which is 1!998cm 1 . _ lV ; ; 7 i ; 77 ; ' , , 6 8 10 12 14 16 reversibly adsorbed on the time scale of Time / seconds the modulation. The relative phase of the CO2 MS signal or the CO2 IR FGC and Figure 8. MS and IR time domain data for the carbonate peaks are near identical and CO oxidation over Rh/AI203 Vertical line retarded with respect to either the CO or shows maximum CO IR signal 02 suggesting a common rate determining

72 step controlling the surface population of the carlmnate and the prt~luction of gas pha~ CO2. The CO section of the spectrum is more complex with the main peak in Figure 7 appearing at 2109 cm t. This corresponds to CO on the Rh~+ sites and the IR FGC shows the feature to be reversible and oscillate in phase with the CO MS trace. The negative peak at 2008 cm-t agrees well in wavenumber with the lower of the two bands associated with the geminal dicarbonyl species and displays a totally flat but gradually dropping FGC. The higher mode of the dicarbonyl is largely obscured by the reversible feature at 2109 cmq but contributes to the slight negative tail of that peak. Together, this would suggest that during the experiment there has been a slight net oxidation of the sample with a gradual decrease in the CO present on Rh and concomitant increase in CO on Rh~+, however it is the CO on Rh~+ that is transformed to CO2 and that the CO associated with the Rh metal as geminal dicarbonyl species is irreversibly adsorbed at the temperature and gas compositions employed. 6. ACKNOWLEDGEMENTS This work was funded by the EPSRC under the Managed Programme in Catalysis. We are also grateful to Dr Ron Brown for isotherm measurements through the SHEFC funded Catalyst Evaluation and Optimisation Service (CATS) 7. REFERENCES

.

5 6 7 8 9 10

11 12

D.R. Coughanowr, Process Systems Analysis and Control, (McGraw-Hill, N.Y., 1991) J.R. Schrieffer and J.H. Sinfelt, J.Phys.Chem., 94 1047-1050 (1990) Yusada, in Studies in Surface Science and Catalysis 84 - Zeolites and related Microporous materials, Elsevier, Amsterdam, 1994 G. Marcelin, J.E. Lester and S.F. Mitchell, J. Catalysis,102, 240, (1986) Y.-E. Li, D. Willcox, R.D. Gonzalez, A.I.Ch.E.J. 35(3),. 423-429 (1989) L.V.C. Rees and D. Shen, Gas Sep. and Purif., 7(2) 83-89 (1993) Gy. Oneystak, D. Shen and L.V.C. Rees, J.Chem.Soc. Farad. Trans., 92(2) 307-315 (1996) L. Song and L.V.C. Rees, Microporous Materials, 6 363-374 (1996) J. K~ger and D.M. Ruthven, Zeolites, 9 267-281 (1989) I.S. Park, M. Petkovska and D.D. Do, Chem.Eng.Sci., 53(4) 833-843 (1998) T.A. Nijhuis, L.J.P. van den Broeke, J.M. van de Graaf, F. Kapteijn, M. Makee and J. A. Moulijn, Chem. Eng. Sci., 52(19), 3401-3404 S. Trautmann and M. Baerns, J. Catalysis, 150, 335-344, (1994)

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science Ltd. All rights reserved.

73

R e a c t i v i t y of novel m e t a l s u b s t i t u t e d h e t e r o p o l y a c i d c a t a l y s t s u s i n g s t e a d y - s t a t e a n d t r a n s i e n t r e s p o n s e kinetics H. T. Randall a, P. L. Mills b and K. Kourtakis b aFirmenich SA, CH-1283 La Plaine, Geneva, Switzerland bDuPont Company, Central Research and Development, Experimental Station, E304/A204, Wilmington, Delaware 19880-0304 USA Abstract

The reactivity of molybdenum-based heteropolyacid (HPA) catalysts was investigated using both steady-state and transient response methods. Using the non-substituted catalyst H4PMollVO4o as the basis, other HPA's were synthesized where either Mo cations or protons were substituted with other metal cations, such as Mn, Co, Ni, Cu, Zn, Ca and Cs. It is shown that the transient reoxidation of reduced HPA's can be quantitatively described by combining a surface oxidation reaction that is first order in gas phase oxygen and second order in surface oxygen vacancies with subsurface oxygen diffusion. The rate of diffusion of oxygen in the lattice of the HPA catalysts during their reoxidation is found to correlate with the activity at steady-state. However, the rate constant for surface reoxidation does not correlate with activity at steady-state. It is suggested that the rate of reoxidation of the catalyst is not the only property that affects the activity at steady-state, but that surface acidity might also play an important role. Finally, the amount of both CO2 and maleic anhydride formed during the transient reduction of reoxidized HPA's by n-butane increases with the oxo-capacity, whereas no correlation is observed between the oxo-capacity and the activity under steady-state conditions. 1. I N T R O D U C T I O N

When compared to most mixed metal oxides, heteropolyacid (H4PMonVO4o) catalysts are unusual since they have both strong acidity and redox properties that make them potentially useful for the selective oxidation of alkanes [1, 2]. HPA catalysts containing molybdenum have been found to be selective for the conversion of alkanes to the corresponding organic acids. Previous work has shown that the partial oxidation of either n-butane, n-pentane, or n-hexane over Mo-based HPA catalysts produced maleic anhydride as the primary reaction product [3]. However, the ability of HPA catalysts to supply lattice oxygen and the effect of various metal substitutions on the lattice oxygen chemistry has not been reported. Synthesis of modified HPA's by either substituting protons in the

74 cationic position, or by substituting the metals in the ionic positions with other metal cations, would conceivably result in new materials with potentially useful catalytic properties. The primary objective of this work is to study the effect of substituting a neutralizing counter-cation, as well as substitution of molybdenum with another transition metal, on both the catalytic structure and reactivity. Particular HPA catalyst systems that are examined where molybdenum substitution is performed include H4PMo11V040 , HxPMoloVMnO40 , HxPMOloVCoO40, HxPMOloVNiO40, HxPMoloVCuO40 , and HxPMOloVZnO40. Typical HPA catalysts with countercation substitution include BaxH4.xPMOloVMn040 , CaxH4.xPMOllVO40, and CsxH4.x PMOllVO40 where x = 0.5 and 2.5 for Cs. The experimental protocols used to evaluate the reactivity are based on both steady-state and transient response m e a s u r e m e n t s of catalyst performance using a conventional fixed-bed microreactor and the TAP@ (Temporal Analysis of Products) reactor.

2. EXPERIMENTAL 2.1 Catalyst preparation The H4PMo11VO4o catalyst was prepared by refluxing stoichiometric amounts of

MoO 3 (Cerac Chemicals, Ronkonkoma, New York, USA) and V205 (Alfa Chemicals, Ward Hills, USA) in water. Aqueous phosphoric acid (85%, J.T. Baker, Phillipsburg, New Jersey, USA) was then slowly added over the course of 20 minutes until the desired stoichiometry was achieved. The reflux was continued for at least 16 hours. The resulting sample was dried, pelletized, crushed and sieved to produce catalyst microparticles with diameters between 250 to 425 ~m. The HsPMOloMV04o compounds, where M = Mn, Co, Ni, Cu or Zn, were prepared as described above, but a stoichiometric amount of the metal dichloride was used in addition during reflux. The CS2.sH1.5PMOllVO4o and CuH2PMOllVO4o compounds were prepared by adding an aqueous solution of either cesium or copper II carbonate to an aqueous solution of H4PMOllV04o. The (VO)2P2O 7 (VPO) catalyst was a commercial catalyst (Chevron, USA). Prior to the kinetic runs, it was activated in a fixed-bed microreactor for at least 100 hours in a continuous flow of 1.5 % n-butane in air at 380~

2.2 Catalyst characterization For the HsPMOloCUVO40 and HsPMoloZnVO40 compounds, single crystals were grown by concentration of their aqueous solutions. Single crystal structure data for both compounds showed that stoichiometric substitution of Mo with Cu or Zn cations had occurred, and that the Zn and Cu cations were in the Keggin unit. The oxidation states of the metals were determined by X-ray photoelectron spectroscopy (XPS), which verified the presence of Mo 6§ V 5§ Cu 2§ and Zn 2§ The BET method showed that the HPA's all had surface areas of approximately 6 m2/g, except for Cs2.~H1.sPMol~VO4o, which has a surface area of 125 m2/g.

75

2.3 Steady-state experiments

The steady-state performance of the HPA and VPO catalysts was evaluated using n-butane oxidation as a test reaction. The experiments were conducted in a fixed-bed stainless steel tubular reactor with an internal diameter of 4.572 mm. The reactor tube was heated by placing it in an isothermal fluidized sand bath in which silicon carbide was used as the heat transfer medium. The reaction temperature was controlled by monitoring the external reactor wall temperature at the midpoint of the catalyst bed. The reaction feed gas was 2 % n-butane in air, which was blended upstream using pre-mixed 10 % butane in nitrogen, pure oxygen, and pure nitrogen. The mixed gas was then introduced to the reactor using mass flow controllers. It was varied from 10 to 100 ml/min so that a wide range of n-butane conversion could be obtained. The amount of catalyst used was about 1 gram. The reactor temperature was varied from 320 to 370~ in 15 to 20~ increments. The reaction feed and product gas compositions were measured on-line using a Hewlett-Packard Model 5890 Series II gas chromatograph equipped with both a flame ionization detector (FID) and a thermal conductivity detector (TCD). A J&W DB-1 capillary column was used for resolution and quantification of the organic reaction products. The activity of the HPA catalysts usually decreased with time-on-stream. This was attributed to a slow decomposition of the HPA that occurs as the temperature approaches 370~ for most Mo-based HPA's [4]. Therefore, the initial activity and selectivity were used to compare the performance of the various HPA catalysts. The initial pseudo first-order rate constant for butane disappearance was taken as a measure of the catalyst activity, since the oxygen was present in a large excess.

2.4 Transient step-response experiments

The kinetics for both the reduction and reoxidation of the HPA and VPO catalysts were studied using a step-up transient response technique with the DuPont T A t ~ reactor system. The experimental setup is described elsewhere [5]. The particular catalyst pretreatment procedures and the operating conditions are summarized below.

Reoxidation of reduced catalysts.

The catalyst was first reduced for 15 minutes in 10 % n-butane in N2 at 340~ using a flow rate of 11.3 ml/min. This was assumed to be a consistent degree of reduction since no CO2 or maleic anhydride was observed. The flow of butane was then stopped, and N2 was then pulsed into the reactor for approximately 2 minutes before a step-up input of 02 was introduced. The step input of 02 was carried out between 320 and 340~ using an inlet mole fraction of 02 of 0.02 and at a total flow rate of 25 ml/min. The feed also contained 20 % of Ne and 78 % of N2 in addition to 02. The dynamic response of 02 was monitored by quadrupole mass spectrometry. The response of CO2 was also monitored during the catalyst reoxidation. Reduction of reoxidized catalysts. The catalysts were first reoxidized for 15 minutes in 2 % 02 in N2 at 340~ using a flow rate of 25.0 ml/min. As in the case of the catalyst reoxidation, it was assumed that the catalysts achieved a consistent degree of oxidation since no more oxygen uptake was detected after this time period. The flow of 02 was then stopped, and N2 was pulsed into the reactor for approximately 2 minutes before a step-up input of n-butane was introduced. Concentration steps of butane were carried out at 340~ using an inlet mole

75 fraction of butane of 0.10 and a total flow rate of 11.3 ml/min. The dynamic responses of n-butane, CO2 and maleic anhydride were monitored by mass spectrometry. Because only a single mass could be monitored during the course of a transient, the reduction and reoxidation experiments were carried out in a periodic fashion until the responses during reduction and reoxidation were reproducible and exhibited cycle invariance. Cycle invariance was generally attained after 3 to 4 periods.

3. RESULTS and DISCUSSION 3.1 Performance of HPA's under steady-state conditions The pseudo first-order rate constants for the disappearance of n-butane (k) and the selectivities for maleic anhydride (SMA)at 40% butane conversion are given in Table 1 for different HPA's and as well as the standard VPO catalyst. The substitution of one Mo cation with either Mn, Co or Ni leads to a considerable decrease in the activity. However, the substitution of either protons or a Mo cation with Cu cations leads to a more than two-fold increase in activity, and to an increase in selectivity in the case of CuH2PMollVO4o. Upon introducing two Cu cations into the anionic position, both the activity and selectivity decrease substantially. The Zn-substituted compound has about the same activity as the non-substituted catalyst, but a slightly higher selectivity. The activity of the Cssubstituted compound is more than two-fold higher than for the non-substituted HPA. The selectivity of the Cs-compound is, however, slightly lower than for the non-substituted catalyst. Both the activity and the selectivity of the VPO catalyst remain substantially higher than for any of the HPA's that were studied. In the next section, the transient kinetics of reduction and reoxidation of the HPA's t h a t have the best performance u n d e r steady-state conditions are examined, i.e., H4PMOllVO40, HsPMOloVZnO40, HsPMOloVCuO40, CuH2PMOllVO40, and Cs2.sH1.sPMOllVO40. 3.2 Catalyst reoxidation step response experiments Figure 1 compares the normalized dynamic responses of the reduced HPA and VPO catalysts at 340~ to a step-up input of oxygen. The normalized response of an inert gas (Ne) is also shown. The difference in area between the response of the inert gas and the oxygen response for a given catalyst is proportional to the total amount of 02 that is transferred to the catalyst during reoxidation. The total amount of oxygen transferred per unit of catalyst weight is referred to here as the oxo-capacity, n% and was calculated according to Equation 1. Q'Y02,0 no2 = ~ [ Finert(t) - Foz(t) ] dt meat 0

(1)

where Finert and Fo2 are the normalized responses of Ne and 02, respectively. The oxo-capacities of the different catalysts are compared in Figure 2 where VPO is seen to have the highest oxo-capacity. The oxo-capacities of the HPA's increase by

77 substituting Mo with either Zn or Cu. Substitution of protons leads to an increase in the oxo-capacity in the case of Cu-substitution, but to a dramatic decrease in the case of Cs-substitution. Table 1. Performance of various heteropolyacid catalysts at steady-state Catalyst H4PMoIIVO4o

S.S.

0.63

41

370

HxPMOloVMnO4o

Mn(Mo)

0.17

-

370

HxPMoloVCo04o HxPMOloVNi04o HsPMOloVCu04o CuH2PMollV04o

Co(Mo)

0.21

-

370

Ni(Mo)

0.36

26

370

Cu(Mo)

1.41

45

370

Cu(H)

1.32

40

370

HxPMo9VCu204o

Cu2(Mo)

0.26

14

370

HsPMOloVZnO4o

Zn(Mo)

0.60

43

370

Cs2.sH1.sPMOllVO4o

Cs(H)

1.57

38

360

(VO)2P~O~

VPO

3.30

75

360

*Pseudo first-order rate constant for butane disappearance In Figure 3, the pseudo first-order rate constants for butane disappearance from Table 1 are plotted against the inverse of the characteristic diffusion time of oxygen in the solid, 1/td. This latter parameter, as well as the surface re-oxidation rate constant kox, was obtained by nonlinear parameter estimation in which the experimental step-up responses at 340~ were matched to the predictions of a detailed model t h a t accounts for t r a n s p o r t and reaction in the fixed-bed microreactor [6]. Figure 3 suggests that a correlation exists between these two parameters, i.e., the rate of subsurface oxygen diffusion is proportional to the steady-state catalyst activity. In Figure 4, the pseudo first-order rate constants for butane disappearance from Table 1 are plotted against the rate constant for surface reoxidation kox. The situation is analogous to the one above, since a correlation exists between these two parameters for all catalysts except for CuH2PMOllVO4o. The latter HPA has about the same activity at steady-state as HsPMOloCUVO4o, but the rate constant for surface reoxidation is greater by about a factor of 30. This suggests that the rate of reoxidation of a HPA is not the only property that determines the activity under steady-state conditions. Surface acidity is another important property that might have an important role during butane oxidation. The importance of Bronsted and Lewis acidity for the partial oxidation of n-butane to maleic anhydride has been previously emphasized by Busca et al. [7].

78

1.2 o

w

c

Ne

1

,O

co 0.8 2

~r

(VO)2P20 7

CS2.sH1.sMo11VO40

9 0.6

N

H4PMo11VO40

m m m

E 0.4 O

HsPMoloZnV04o

z 0.2

HsPMoloCuVO~

CuH2Mo11VO4o 0

50

100 Time, s

150

200

Figure 1. Responses of 02 to a concentration step of 02 over reduced HPA and VPO catalysts. Conditions: T = 340~ P = 1.01.10 ~ Pa, Yo2,o= 0.02, Q = 25 ml/min.

3.3 Catalyst reduction step r e s p o n s e e x p e r i m e n t s Typical responses for n-butane, CO2 and maleic anhydride to a step-up input of n-butane over reoxidized CuH2PMollVO4oat 340 ~ are shown in Figure 5. Traces of acetic acid (m/e = 60), furan (m/e = 68) and acrylic acid (m/e = 70) were also detected but are not shown. The formation of CO (m/e = 28), which is a major product under steady-state conditions, could not be monitored since N 2 ( m / e = 2 8 ) was used as a carrier gas. The responses of both CO2 and maleic anhydride have a maximum and subsequently decrease as lattice oxygen from the catalyst is depleted. The concentration of butane is zero at the onset of the transient and reaches a steady-state exit concentration after about 15 minutes at which point the catalyst is fully reduced. The total amounts of maleic anhydride and CO2 formed per kg of catalyst during the reduction were determined by integrating the experimental responses of maleic anhydride and CO2. These values are plotted against the oxo-capacity of the corresponding HPA's in Figure 6. A correlation exists between the oxocapacity and the amounts of products during the reduction of reoxidized HPA's. It should also be noted that this is not true under steady-state conditions, for which there is no correlation between oxo-capacity and activity. The Cs-substituted compound, for instance, has the lowest oxo-capacity, but also has the highest activity at steady-state.

79

(VO)2P=07

Cs2.sH1.sPM~ HsPMOloZnVO4o

CuH2PMo11VO4o HsPMOloCUVO4o H4PMo11VO4o 0

0.01

I

I

i

0.02

0.03

0.04

0.05

Oxo-capacity, mol 021kgcat Figure 2. Oxo-capacity of HPA and VPO catalysts.

c

o

2

em

E

=

~'~

1.6-

Cu (Mo)

|C gm l . 2 =

T 370 0 C

#

,"

'

S

9

#

#

"

0.00

T Z 360~

e

Zn (Mo)

0.40

9

#

9

~

,,,

C

= 8o.8- NS.' ~,

Cs (H)

Cu (H)

I

I

I

I

0.01

0.02

0.03

0.04

0.05

Rate of oxygen diffusion, 1/td (S"1) Figure 3. Steady-state pseudo first-order rate constant for butane disappearance versus the rate of subsurface oxygen diffusion 1/td at 340~

80

Catalyst oxidation rate constant (m Z/kg s) 0.45 0.5 0.4

C

o

i m

a.

2'

E

~e~' r , m ~ o ~

1. 6 .

em9 ~ =

1.2-

~9

T=370oC NS

o0.8-

9

er162

,

Cs (H)

Cu (Mo)

8

9

".~-"

,,

w

Cu (H)

" ~ T=360~

Zn (Mo)

|

>l

"0

m

0

r

0.000

I

I

0.010

0.020

0.030

Catalyst oxidation rate constant (m Zlkg s) Figure 4. Steady-state pseudo first-order rate constant for butane disappearance versus the catalyst reoxidation rate constant at 340~ 4. CONCLUSIONS The reactivity of molybdenum-based heteropolyacid (HPA) catalysts was investigated using both steady-state and transient response methods. The oxidation of n-butane was used as a test reaction so that a standard vanadiumphosphorus oxide catalyst could be used as a basis for comparison. It was shown that the rate of oxygen diffusion in the HPA catalyst lattice during their reoxidation correlates with the steady-state activity. However, the rate constant for surface reoxidation does not correlate with activity at steady-state. This suggests that the rate of reoxidation of the catalyst is not the only property that affects the activity at steady-state, but that surface acidity might also play an important role. Finally, the amount of both COs and maleic anhydride formed during the transient reduction of reoxidized HPA's by butane increases with the oxo-capacity, whereas no correlation is observed between the oxo-capacity and the activity under steady-state conditions.

81

0.12 C 0 0

Ne

0.1

4~

Butane

0.08

o

3 - ~==h

"L_ - 0.06

o

~

#=.

C02

m

O.O4

26"

X==~

1

MAN

0.02

- I J - - - - - - - , ~ - - - - - , ~ T ~ - !

0

50

100 150 Time, s

200

o

- 0

250

Figure 5. Responses of Ne, n-butane, CO2 and maleic anhydride to a concentration step of n-butane over reoxidized CuH2PMollVO4o. Conditions: T = 340~ P = 1.01.105 Pa, Y%o = 0.10, Q = 11.3 ml/min. 0.014

/

ca 0 . 0 1 2 m .

0

E E i...

4,,,I

qD

C02 ./

0.01 0.008-

/-

o.oo6 0.004 -

,. 0 . 0 0 2 o

0

.m

9

i

" ...."""

i

i

0

...." .~176

|::..:::.:::~ .....

...o.

..""

.... 9

..-.O""....---"

MAN

I

I

0.01

0.02

0.03

Oxo-capacity, mol 02/kgcat

Figure 6. Total amounts of maleic anhydride and CO2 formed during the reduction of reoxidized HPA catalysts. Conditions: T = 340~ P = 1.01.105 Pa, Yc4,o= 0.10, Q = 11.3 ml/min.

82 5. R E F E R E N C E S

1. F. Cavani and F. Trifiro, Catalysis, Royal Society of Chemistry, Vol. 11, 223, 1994. M. Misono, Catal. Rev.-Sci. Eng., 29 (1987) 269. 3. K. Bruckman, J. Haber and E. M. Serwicka, Faraday Discuss. Chem. Soc., 87 (1989) 173. G. A. Tsigdinos, "Heteropolycompounds of Molybdenum and Tungsten", in Topics in Current Chemistry, Vol. 76, Springler-Verlag, Berlin, 1978. H. T. Randall, P. L. Mills and J. S. McCracken, presentation at the 3rd World Congress on Oxidation Catalysis, San Diego, CA, September 21-26, 1997. P. L. Mills and H. T. Randall, Multiregion distributed parameter dynamic model of a fixed-bed microreactor, paper presented at the SIAM Annual Meeting, Toronto, July 13, 1998. 7. G. E. Busca, G. Finocchio, G. Ramis and G. Ricchiardi, Catal. Today., 32 (1996) 133. 0

0

0

0

6. ACKNOWLEDGMENTS The authors wish to thank J. Scott McCracken for his expert assistance and advice in conducting the T A t ~ reactor experiments. Dr. John D. Sullivan and Mr. W. E. (Mike) Guise provided valuable assistance in execution of the steadystate HPA catalyst testing experiments. Dr. Richard Harlow performed the various characterization experiments used to verify the HPA metal substitution.

Reaction Kinetics and the Developmentof Catalytic Processes G.F. Froment and K.C. Waugh(Editors) 9 1999Elsevier Science B.V. All rights reserved.

83

Use of isotopic transient methods for mechanistic analysis of ethylene hydroformylation over 4 wt% Rh/SiO2 catalyst Steven S. C. Chuang 1, Scott A. Hedrick, and Mark A. Brundage Department of Chemical Engineering, The University of Akron, Akron, OH, 44325-3906 Abstract The transient responses of deuterated ethylene, ethane, and propionaldehyde to a deuterium pulse during ethylene hydroformylation have been studied on 4 wt% Rh/SiO2 at 0.1 MPa and 483 to 573 K. This study shows the complexity of the deuterated product formation from the deuterium pulse. The deuterated ethylene/ethane responses overlapped at all temperatures, indicating that rapid H/D exchange and alkyl hydrogenation take place on the catalyst surface consistent with the Horiuti-Polanyi mechanism. The d l- and d2-propionaldehyde responses exhibited a two-hump response to the D2 pulse, indicating two distinguishable deuteration pathways contributing to their formation. Activation energy for the production of deuterated propionaldehyde was found to be directly proportional to the number of deuterium atoms in the final product. This isotope effect supports a previous study which suggests acyl hydrogenation to be the ratelimiting step.

1. INTRODUCTION The stimulus-response technique is an experimental method which has found wide use for studying various systems in science and engineering [1]. One particularly useful stimulus-response technique for study of heterogeneous catalytic reactions is steady-state isotopic transient kinetic analysis (SSITKA). This technique involves replacement of a reactant by its isotopically labeled counterpart, typically in the form of a step or pulse input function. Producing an input function with isotope-labeled reactant permits the monitoring of isotopic transient responses while keeping the total concentration of labeled plus nonlabeled reactants, adsorbates, and products at steady state conditions. These transient responses, which are monitored by mass spectrometry, carry kinetic and mechanistic information about the steady-state reaction. This fundamental information can be used as a guide for the development of more active and selective catalysts [2]. Ethylene hydroformylation is the reaction of ethylene with syngas (CO/H2) to form propionaldehyde as the desired product and ethane as the byproduct. The reaction of ethylene and syngas has been shown to be a useful probe reaction to study oxygenate activity and selectivity of supported metal catalysts in the FischerTropsch (F-T) synthesis [3]. The reaction has also served as an excellent model reaction for SSITKA studies [4-9]. Use of 13CO as a tracer has allowed 1Corresponding author

84 d e t e r m i n a t i o n of the reaction p a t h w a y s involving carbon-containing intermediates and identification of the rate-limiting step [4,5]. The hydrogenation step was found to be a critical step in determining the rate and selectivity for producing propionaldehyde [5]. Thus, tracing the hydrogen pathway would help develop a f u n d a m e n t a l u n d e r s t a n d i n g of the role of hydrogen in the hydroformylation reaction. Deuterium pulse tracing was utilized in this study to trace the hydrogen reaction pathway for the hydroformylation reaction on 4 wt% Rh/SiO2 at 0.1 MPa and temperatures ranging from 483 to 573 K. The proposed reaction pathway for hydrogen/deuterium is shown in Figure 1.

C2H4(g)

CO (g)

~

C2H~De.,CO (g)

aeyl hydrogenation

H2 / D2 (g)

*~ *CO *H/*D *C2I-~D5.x-'~ ,r *C2H~Ds.~CO r~~ D~ CO adsorbed acyl spillove *H/* insertion ethyl *Si-OH/*Si-OD hydrogenation *Si-OH/*Si-OD C2HxD6.x (g) acyl ~ C2I-I~De.xCO(g) hydrogenation

Figure 1. Reaction pathway for incorporation of D2 into final product. The first step in Figure I is adsorption of D2 onto the metal site. The deuterium can then migrate, i.e. spillover, onto the silica support or deuterate adsorbed ethylene to form an adsorbed ethyl. This ethyl species can subsequently be deuterated or undergo CO insertion to form adsorbed acyl. Finally, to form propionaldehyde, this acyl group undergoes deuteration by either metalchemisorbed deuterium or spillover deuterium. Though not noted specifically in the figure, hydrogen and deuterium can undergo rapid exchange with adsorbed ethylene, ethyl, and acyl groups [8-13]. 2. EXPERIMENTAL A 4 wt% RtdSiO2 catalyst was prepared by the incipient wetness impregnation method using an aqueous solution of RhC13e3H20 (Alfa Products) onto a large pore SiO2 support (Strem Chemicals, surface area of 350 m2/g). After impregnation, the catalyst powder was dried in air overnight at room temperature and then was reduced in flowing H2 at 673 K for 16 h. Hydrogen uptake at room temperature was determined by the flow chemisorption method to be 122 ~tmol/gcat, corresponding to a Rh particle size of 1.5 nm and a dispersion of

85 0.62, a s s u m i n g an a d s o r p t i o n s t o i c h i o m e t r y of *H:Rhsite=l. (* denotes a chemisorbed species.) Approximately 35 mg of catalyst powder, which was pressed into a thin disk, was placed in the IR reactor where it was supported by two CaF2 rods t h a t allow the i n f r a r e d b e a m to pass t h r o u g h the sample. Specific details of the experimental a p p a r a t u s and IR reactor cell have been reported elsewhere [14]. The reaction was carried out under differential conversion at 483, 513, 543, and 573 K at 0.1 MPa with a total flow rate of 120 cm3/min. The catalyst was reduced a minimum of two hours in flowing hydrogen at 673 K between runs. The reactant mixture consisted of equimolar amounts of H2, CO, C2H4, and He. The H2 stream consisted of 2 vol% Ar, which acted as an inert tracer t o determine the flow pattern in the reactor system without interacting with the catalyst surface. Upon r e a c h i n g s t e a d y - s t a t e flow at the desired conditions, an HP-5980A gas chromatograph (GC) equipped with a flame ionization detector (FID) was used to determine the steady-state concentrations of reactants and products. A six-port valve was then used to pulse 10 cm 3 of D2 into the H2 stream. After the pulse, IR spectra were recorded by a Nicolet 5SXC spectrometer with a DTGS detector in order to monitor t r a n s i e n t responses of adsorbates on the catalyst surface. The t r a n s i e n t responses of the gaseous effluent were simultaneously recorded by a Balzers QMG 112 mass spectrometer (MS) interfaced to a computer t h a t allows m e a s u r e m e n t of eight m/e (i.e., amu) as a function of time. The m/e ratios monitored were 2 (H2), 3 (HD), 4 (D2), 30 ( d 2 - e t h y l e n e / d o - e t h a n e ) , 31 (d3ethylene/dl-ethane), 32 (d4-ethylene/d2-ethane), 33 (d3-ethane), 40 (Ar), 58 (dopropionaldehyde), 59 (dl-propionaldehyde), 60 (d2-propionaldehyde), and 61 (d3propionaldehyde). The di- prefix indicates the number of deuterium atoms in the molecule. Each experiment was performed four times at each condition. 3. RESULTS AND DISCUSSION 3.1. Effect of temperature on product activity and selectivity. Table 1 lists the steady-state rate of product formation and selectivity from the CO/I-I2/C2H4~e reaction (1/1/1/1) as a function of temperature calculated from GC data. Activity is characterized by turnover frequency (TOF), or the n u m b e r of m o l e c u l e s r e a c t i n g p e r site p e r second. S e l e c t i v i t y is d e f i n e d as TOFc~HsCHO / TOFc~He. Table 1 Activity and selectivity of products from function of temperature. Total flow rate Temp. (K) TOF (s -1) X 103 CH4 C2H6 483 0.18 4.35 513 0.17 12.4 543 0.19 30.4 573 0.24 97.7

the CO/H2/C2H4/He reaction (1/1/1/1) as a = 120 cm3/min and pressure = 0.1 MPa. Selectivity C2H5CHO C3-C5 HC 0.82 0.03 0.189 1.37 0.08 0.112 1.95 0.34 0.064 2.94 1.11 0.030

86 The two major products formed from this reaction are ethane and propionaldehyde, whose formation rates continue to increase through 573 K. Selectivity t o w a r d p r o p i o n a l d e h y d e declines m a r k e d l y with i n c r e a s i n g temperature. The C1 and C3+ hydrocarbon (HC) formation rates increase substantially with temperature, suggesting that CO dissociation is enhanced at high temperature. The TOF's of ethane and propionaldehyde for four series of rate measurements are plotted in Arrhenius form in Figure 2a. Activation energy of the reaction remains relatively constant, indicating that the intrinsic properties of the catalyst were not significantly altered during these repeated runs. The activation energy for ethane was found to be 19.4_+1.1 kcal/mole and the activation energy for propionaldehyde formation was found to be 8.7_+0.5 kcal/mole, agreeing well with previous studies on Rh/SiO2 [4,6].

10 o

a

b

I

i,

" _ 10-1 - O r O

0

Ethane Propionaldehyde

rl

I

10 0

I

mol 18.3 kcal/mol 10-1

!

9

10-2 10-3

.1

" 8.7•

22.1 k~cal/m

kcaU .l-Tg'"

!

10.4 1.7

1.8

1.9

1000fr

2

"

.~ r ~0_~ro~iona]~e~e :---~,~-- - ro lona e e ," q~,-i~roplonalqepyqe a~-t'ropmnalaenyae I~ I I 1.8 1.9 2.O

(K-1)

1000fr

X~ -

10-2

10-3

(K -1)

Figure 2: (a) Arrhenius plot for ethane and propionaldehyde formation (repeated symbol at a specific temperature is the result of repeated runs) and (b) Arrhenius plot for deuterated propionaldehyde formation. The low activation energy for propionaldehyde formation falls into the range of diffusion limitation effects. The effect of diffusion limitations on the propionaldehyde response can be estimated by calculation of the generalized Weisz-Prater criterion [15,16]. The equation is 2

s

(rC~HsCHO)obsPsLg(C ) = ,c << 1, 2D eAJo' g(C) dC where (rc~nsCHO)obs is the observed rate of propionaldehyde formation,p s is the bulk density of the catalyst, L is the thickness of the disk, DeA is the effective diffusivity of CO at the surface of the pellet, and g(C)=[CO]-0.39[H2]O.90[C2H4] 1.15 [5]. Since all

87 E(t) (s -1) 0.03

h

i

I

I

I

I

I

I

I

I(t) i . . . .

0.02

E(t)

I

-

3 1071

~

"

2 10.7t

/ ~

.

[(t) 0

--

10-1] 0

-

i0-10 10-1o

o 0.03

b[

[ [ I [ :

:

I(t) I . . . .

']

'

I(t)

-

4 10l~

0.02 0.01 0 0.04

IC

I I

I I

I I

I I

I I l(t)

I'

I I

I I

'r

I I

'

[

9

'

I(t)

1 I08

3 10t~

0.03 5 10"9

0.02 0

0.01

0

30

60

90

0 0

60 120 - - o - - - 483 K ~--513 K

180 240 S ~ 543 K ~ - - 573 K

[} o

60 120 180 240 s 483 K o--- 543 K 513 K ~ 573 K

F i g u r e 3: T r a n s i e n t r e s p o n s e s of (a) H D , (b) d l - e t h y l e n e , (c) d 2 - e t h a n e , (d) dop r o p i o n a l d e h y d e , (e) d l - p r o p i o n a l d e h y d e , a n d (f) d 2 - p r o p i o n a l d e h y d e to a D2 into H2 pulse.

88 the concentrations are the same, g(C) simplifies to [H2] 1.66. The smallest value of the binary diffusivity was that of the CO-C2H5CHO system at 0.035 cm2/s. This value was used with tortuosity and porosity of the pellet in the calculation of the Weisz-Prater criterion to obtain an upper limit of ~=0.003 at 573 K. Thus, the effect of m a s s t r a n s f e r l i m i t a t i o n s on p r o p i o n a l d e h y d e f o r m a t i o n was insignificant. The transient responses of HD, dl-ethylene, d2-ethane, do-propionaldehyde, dlpropionaldehyde, and d2-propionaldehyde to the D2 pulse at various temperatures are shown in Figure 3. d3-, d4-ethane and d3-, d4-propionaldehyde are not included in the figure due to their low formation rate. Inset figures are the I(t), or MS intensity, curves which show clearly an increase in product formation as temperature is increased. The E(t) curves were normalized by I(t) E(t) = ~ , ~ ( t ) dt where E(t) is the normalized curve. E(t) allows calculation of average residence time of each species and comparison of their residence time distributions. There is no distinguishable difference in residence time between the D2 response curve and the di-ethane/di-ethylene responses. This indicates that the residence times of the i n t e r m e d i a t e s is too short for d e u t e r a t e d ethylene and e t h a n e to be distinguished from the transient response curves. A steady-state GC-MS analysis of the CO/D2/C2H4 reaction was conducted at 513 K and 0.1 MPa to determine d e u t e r i u m distributions in ethylene and ethane. These results are shown in Table 2. Table 2 Deuterated ethylene and ethane deuterium distributions. Ethylene Ethane do --2.45% dl 87.00% 1.37% d2 7.34% 0.46% d3 1.12% 0.09% d4 0.15% 0.02% d5 --0.01% This analysis revealed t h a t the positive m]e=30 response consists of 75% C2H2D2 and 25% C2H6. The formation of do-ethane in the presence of D2 can be explained by the classical Horiuti-Polanyi m e c h a n i s m [17,18], which proposes a rapid equilibrium between adsorbed ethylene and adsorbed alkyl species. *C2H4 +*H e-> *C2H5 + * (* indicates active site.) *C2H5 +*H ~ C2H6(g)+ 2* This rapid equilibrium allows release of adsorbed hydrogen from the adsorbed alkyl species to react with other adsorbed alkyl species to form do-ethane. Both m/e=31, which represents d3-ethylene and d l - e t h a n e responses, and m/e=32, which represents d4-ethylene and d2-ethane, were also resolved with the GC-MS experiment. It was determined that 55% of m/e=31 can be assigned to dl-ethane

89 with the remainder d3-ethylene and t h a t 75% of nde=32 can be assigned to d2ethane, with the remainder being d4-ethylene. The transient di-propionaldehyde responses are shown in Figure 3d, e, and f. The dl- and d2-propionaldehyde responses exhibited two humps, or two maxima at temperatures below 543 K. The first-hump is due to H/D exchange on the adsorbed acyl species while the second-hump is due to H/D exchange on the adsorbed alkyl species which undergoes CO insertion [8]. The decay portion of the second hump response for d l-propionaldehyde has been suggested to be due to the reaction involving Si-OD, the d e u t e r i u m on the silanol [8,9]. The dlpropionaldehyde may be in the form of either C2H5CDO or C2H4DCHO, which can not easily be distinguished by mass spectroscopy due to the overlapping of their fragmentation patterns with other deuterated products. Increasing the t e m p e r a t u r e increased the formation rate of deut erat ed propionaldehyde and resulted in smoother response curves as compared to those at 483 K. As temperature increased, the second hump nearly merged with the first hump, m a k i n g the two hum ps almost indistinguishable. The higher sensitivity of the first hump to temperature indicates it has a higher activation energy. Increase in t e m p e r a t u r e to 573 K caused the dl-, d2-, and d3propionaldehyde responses to completely merge together and overlap the D2 response, indicating t hat the rates of reaction for the formation of these products are high and the difference in residence time of intermediates leading to these products is too small to give any significant differences. Table 3 shows the relative amount of deuterated propionaldehyde produced during the deuterium pulse which were estimated by utilizing the MS response areas. At each temperature, the amount of lower deuterated products is larger than those of higher deuterated products indicating deuterium is not randomly distributed in the products, but is according to the deuteration kinetics. An isotope effect is clearly observed for propionaldehyde formation as evidenced by the increase in the fraction of high deuterated propionaldehyde production with temperature. Table 3 ....Deuterium distribution in propionaldehyde. 513K 543K

dl d2 d3 d4 d5 d5

First ,,, H u m p 3.86% 5.66% 15.55% 0.37% 0.25% 0.03%

Second ,, H u m p 47.61% 26.67% ---------

First , Hum p 6.20% 7.96% 14.56% 6.21% 3.64% 2.53%

Second Hump 38.45% 20.46% ---------

573 K First .... H u m p .... 7.91% 14.62% 14.52% 5.68% 1.79% 0.98%

Second Hump 31.15% 23.35% ---------

This trend is clearly demonstrated by the Arrhenius plot of various deuterated propionaldehydes in Figure lb. The activation energy is directly proportional to the number of deuterium atoms in the final product. The presence of a deuterium

90 isotope effect supports a previous study which suggests acyl hydrogenation to be the rate-limiting step for propionaldehyde formation [5]. 4. CONCLUSIONS The rate of production of all products increased with temperature; however, selectivity toward propionaldehyde decreased markedly. This study shows the complexity of the deuterated product formation from the deuterium pulse. The dland d2-propionaldehyde responses exhibited a two-hump response to the D2 pulse, indicating two distinguishable deuteration pathways contributing to their formation. The first-hump is due to H/D exchange on the adsorbed acyl species while the second-hump is due to H/D exchange on the adsorbed alkyl species which undergoes CO insertion. The decay portion of the second hump response for d l-propionaldehyde has been suggested to be due to the reaction involving SiOD, the deuterium on the silanol. Increasing the temperature increased the formation rate of deuterated propionaldehyde and resulted in smoother response curves as compared to those at 483 K. As temperature increased, the second hump nearly merged with the first hump, making the two h u m p s almost indistinguishable. The higher sensitivity of the first hump to t e m p e r a t u r e indicates it has a higher activation energy. Increase in temperature to 573 K caused the dl-, d2-, and d3-propionaldehyde responses to completely merge together and lead the D2 response, indicating that the rates of reaction for the formation of these products are high and the difference in residence time of intermediates leading to these products is too small to give any significant differences. The d e u t e r a t e d ethylene/ethane responses overlapped at all temperatures, indicating that rapid H/D exchange and alkyl hydrogenation is taking place on the catalyst surface consistent with the Horiuti-Polanyi mechanism. This is further supported by the formation of C2H6 during the steady-state CO/D2/C2H4 reaction. Activation energy for the production of deuterated propionaldehyde was found to be directly proportional to the number of deuterium atoms in the final product. This isotope effect supports a previous study which suggests acyl hydrogenation to be the rate-limiting step. The effect of mass transfer limitations on the production of propionaldehyde was found to be insignificant. The transient responses of the deuterated products can be further analyzed by use of compartment modeling, which can provide the framework for estimation of intermediate residence times and kinetic parameters [6,8,9]. 5. ACKNOWLEDMENT The authors gratefully acknowledge the partial support of this research by the National Science Foundation under grant CTS-9421119996 and the Ohio Board of Regents Research Challenge Grant.

91 6. R E F E R E N C E S

10 11 12 13 14 15 16 17 18

O. Levenspiel, Chemical Reaction Engineering, 3rd ed., New York, 1999. Rabo, J., in "Proc. 10th Int. Cong. Catal., Part A" (L. Guczi, F. Solymosi, and P. Tetenyi, Eds.). Akad. Kiadd, Budapest, Hungary and Elsevier Science, Amsterdam. 1993. P. R. Watson G. A. and Somorjai, J. Catal., 74 (1982) 282. M. W. Balakos and S. S. C. Chuang, J. Catal., 151 (1995) 253. M. W. Balakos and S. S. C. Chuang, J. Catal., 151 (1995) 266. M. A. Brundage and S. S. C. Chuang, J. Catal., 164 (1996) 94. M. A. Brundage, M. W. Balakos, and S. S. C. Chuang, J. Catal., 173 (1998) 122. M. A. Brundage and S. S. C. Chuang, J. Catal., 174 (1998) 164. M. A. Brundage, S. S. C. Chuang, and S. A. Hedrick, Catalysis Today (in press). Ponec, V., and Bond, G. C., in "Catalysis by Metals and Alloys" (B. Delmon and J. T. Yates, Eds.), Vol. 95, p. 449, New York. 1995. S. A. Goddard, R. D. Cortright, and J. A. Dumesic, J. Catal., 137 (1992) 186. C. Kemball, J. Chem. Soc. 735 (1956). G. C. Bond, J. J. Philipson, P. B. Wells, and J. M. Winterbottom, Trans. Faraday Soc. 62 (1966) 443. S. S. C. Chuang, M. A. Brundage, M. W. Balakos, and G. Srinivas, Appl. Spectrosc., 49(8) (1995) 1151. P. Weisz and C. Prater, Adv. Catal., 6 (1954) 143. G. F. Froment and K. B. Bischoff, Chemical Reactor Analysis and Design, New York, 1990. A. Ozaki, Isotopic Studies of Heterogeneous Catalysis, New York, 1977. J. Horiuti and M. Polanyi, Trans. Faraday Soc., 30 (1934) 1164.

This Page Intentionally Left Blank

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

93

CO oxidation over a supported Pt catalyst" transient kinetics using temporal analysis of products (TAP) A.H.J. Colaris, J.H.B.J. Hoebink, and J.C. Schouten Laboratory of Chemical Reactor Engineering, Eindhoven University Technology, P.O. Box 513, 5600 MB Eindhoven, The Netherlands

of

Abstract Pulse type transients were introduced in a TAP micro reactor filled with an inert quartz bed, with a porous SiOe support, and with Pt/SiOe catalyst. Under Knudsen conditions, rate parameters concerning diffusion of several noble gases and diffusion, adsorption, and desorption of reactants and products were determined by nonlinear regression of the pulse responses. Only part of the intrapellet catalyst volume interferes with the admitted pulses. Bed scale diffusivities might be approached with values for non porous pellets.

1. I N T R O D U C T I O N Temporal analysis of products (TAP) is an elegant experimental method to obtain qualitative and quantitative kinetic information from transient experiments. Transients or perturbations are introduced in a TAP [1] micro reactor by admitting small and narrow pulses of reactants. No carrier gas is used. Transport through the catalyst bed takes places by Knudsen diffusion only. Time resolved pulse responses are measured via quadrupole mass spectrometry (QMS) at the outlet of the micro reactor. Quantitative kinetic data, e.g. rate parameters like activation energies and pre-exponential factors from TAP experiments were reported for CO oxidation over Pt sponge catalysts [2-4]. This use of non porous Pt pellets allowed to avoid non uniform concentration profiles as possibly met in supported catalysts. Such profiles might be due either to a limited diffusion rate in combination with a fast reaction or to accumulation effects induced by the transient. CO oxidation on a supported catalyst has been studied with TAP as well, but only qualitatively [5]. The current work addresses CO oxidation kinetics over supported catalysts with TAP and reports about the interactions between the catalyst and the individual gaseous components involved.

94

2. E X P E R I M E N T A L 2.1 Material c h a r a c t e r i s a t i o n In order to achieve isothermal conditions, the fixed bed in the micro reactor consisted of two inert quartz beds, one at the front and one at the back of the micro reactor, and a catalyst bed in between. Each bed had a length of about 12 ram. The catalyst, 2 wt% Pt/SiO2, was supplied by Engelhard. The inertness of the non porous quartz (Merck) was checked in separate experiments. The effects of the porous silica support (Engelhard) were studied with a reactor completely filled with support. Some properties of the used materials are shown in Table 1. The pellet diameter is determined from sieving, the pore diameter and intrapellet porosity from N2 physisorption. The interpellet bed porosity was calculated from the intrapellet bed porosity and the total porosity (~total = ~bed + [1--~bed]'~penet). The catalyst showed a 12% dispersion of Pt, determined with CO chemisorption at 323 K (Micromeretics ASAP). Before the experiments, the beds were p r e t r e a t e d at 773 K: O2 oxidation followed by H2 reduction, each lasting between 1.0 and 1.5 hours. TAP experiments were performed between 323 and 473 K u n d e r ultra high vacuum (UHV) conditions, ranging from 10 .9 m b a r near the QMS in the analytical chamber up to 10 .6 mbar near the micro reactor in the reaction chamber. Table 1. Properties of used bed dpellet Packing [pm] Quartz 315 - 400 150 - 212 SiO2 2 wt% Pt/SiO2 150 - 212

packings. Etotal

E pellet I

~: bed

dpore I

[--] 0.41 0.85 0.85

[--] 0 0.74 0.76

[--] 0.41 0.42 0.38

[nm] -37 + 7 32 + 7

surf. area 1 BET [m2/g] 0.02 2 114 120

Obtained by N2 fysisorption; 2 Krypton BET.

2.2 E x p e r i m e n t a l p r o c e d u r e s The experiments were performed according to the following procedure: first a multipulse consisting of several h u n d r e d s (typically 600) of pulses is admitted to the inlet of the reactor at a frequency of one pulse per second. Then, after approximately 300 seconds, the multipulse was repeated in order to estimate the degree of surface coverage qualitatively. Finally, single pulses were admitted at time intervals of about 3 seconds in order to improve the signal to noise ratio by signal averaging. Pulse repetition also ensures a quasi equilibrium of the surface coverage in the bed. Each pulse contains typically 2"1015 molecules consisting of 90 vol% gaseous component and of 10 vol% Ar as reference gas.

95

2.3 Modelling Experimentally obtained pulse responses were simulated by solving Equation (1), representing non steady state Knudsen diffusion with, whenever necessary, adsorption/desorption of component i on the active sites (Ric), and on the support (Ris): ~

OCi

Kn ~2 C i -1- (1

- De c~t --

aX2

-

Ebed ) a g

i

i

i

i

(1)

[Lt, sR, + Lt, cRc]

This equation corresponds with an interpretation that the catalyst pellets are non porous being the case if the time scale of pellet diffusion exceeds largely the time scale of bed diffusion [6]. The boundary conditions were taken as reported [3,6]. Diffusivity and rate parameters were obtained by non linear regression of the full TAP response curves with the model solution by the application of NAG Fortran Library Routine D03PGF [7].

Table 2. Elementary sorption steps and corresponding source terms. Elementary steps Source term R i CO

+

t

~,

~

CO

t

RC~

--

- kC~

C02

+

t

-" ~.

~"

C02

t

RCO2

-

_ ~" ~co2a,8 (1-Ocoe) Cc% + k c~ d,s 0co2

CO

+

*

~

"

CO

*

RC~

-

- kC~

C02

+

*

~

C02

,

RC~

-

- k c~ a,c (1-Ocoe) Cco 2 + I~ ~c~ d,c 0co2

02

+

*

~

02

*

R~

-

- k~

+

*

~

20

*

Ro~%

-

02*

~

(1- Oco) Cco + kC~

(1-0co) Cco + k C ~

c (1-Ooe) Cos + k~

Oco

Oco

c Oos

oo [8]

Diffusivities in the inert quartz were determined without involvement of the source terms. As pulse responses of CO and CO2 from the support were somewhat retarded compared to Ar, sorption on the support was taken into account for these components. All adsorption/desorption rates on catalyst and support were expressed in elementary steps, corresponding with the schema of Table 2. The diffusivities and rate parameters achieved with the support, were applied when assessing the kinetic parameters for the Pt/SiO2 catalyst. Separate continuity equations with appropriate boundary conditions were added to Equation (1) as to deal with the combination of inert quartz beds and a catalyst bed in between [3].

96 3. R E S U L T S

AND DISCUSSION

3.1. T h e i n e r t b e d 3x10 -8

0.0025

~"

N , , , ..~,

o.oo2o

2x10 -8 ~'

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

I I

\

lx10.8

0.1

0.2

0.3

0.4

0.5

t[s]

Figure 1. Regression of a typical TAP pulse response compared with the model: 90 wt% CO (10 wt% Ar) over quartz.

....

, ....

, ....

, ....

T = 323 K

0005 . . . . . . . . . . . . . . . . . . . . . . . . 0.05 0.10 0.15

0.20

0.25

M -lr2 [(moi/g) "1/2]

Figure 2. Diffusivities (DKne) in the quartz bed of several gases versus the inverted square root of the molar weight (M-1/2) at T = 323 K.

0.0010 . . . . . . . . . . . . . . . . . . . . . . .

, ....

Ne

0.0008

0.002 ~

o.ool

02

0.0010 ~ X e

00

0.003

co

Kr f

0.0015

CO pulse response

21//Ne

c%~

CO2.__~

0.0006

Dp(Kr)

0.0004 Xe i

,

0.0002 ,

I

18

,

,

,

,

i

,

,

,

,

i

20

,

,

,

,

2'1

,

,

,

,

T 1/2 [K 1/2]

Figure 3. Diffusivities (DK%) in the quartz bed of Ne, CO, CO2, and Xe versus the square root of the temperature (T1/2).

35'

..

t

,

,

,

i

,

17

,

,

,

i

19

,

L

i

i

i

<
,

,

.

,

,

,

,

,

25

T lr2 [K lr2]

Figure 4. Diffusivities of Kr versus the square root of the temperature. The solid line represents non porous pellets (Dnp: tpellet >> tbed). The dotted line concerns intrapellet concentrations that equal the concentration at the outer pellet surface (Dp: tpeuet << tbed) at any time (very fast intrapellet diffusion).

To e n s u r e t h a t e x p e r i m e n t s w e r e p e r f o r m e d in t h e K n u d s e n r e g i m e , t h e K n u d s e n r e l a t i o n for t h e effective d i f f u s i v i t y w a s c h e c k e d w i t h v a r i o u s gases. F i g u r e 1 p r e s e n t s a t y p i c a l e x a m p l e of t h e r e g r e s s i o n of a m e a s u r e d o u t l e t pulse. Effective d i f f u s i v i t i e s d e d u c e d f r o m t h e m o d e l a r e s h o w n as a f u n c t i o n of M -1/2 in F i g u r e 2 a n d of T 1/e in F i g u r e 3. T h e r e s u l t s a r e in line w i t h t h e K n u d s e n relation.

97

3.2. The support bed For a reactor completely filled with the porous silica support, diffusivities of Ne, Ar, Kr, and Xe were determined, and Figure 4 presents typical data for Kr. The solid line in the figure represents the diffusivity for non porous silica pellets (Dnp) obtained from similar experiments with inert quartz but corrected for the pellet diameter via the Knudsen relation. The dotted line shows the diffusivity (Dp) when corrected for the internal pellet porosity [6], which situation refers to very fast intrapellet diffusion compared to diffusion on the scale of the bed. Since the data points approach the solid line, it is concluded t h a t only part of the intrapellet volume interferes with the pulses fed, and t h a t the porous pellets can be approximated roughly as being non porous for the interpretation of TAP pulse responses. This conclusion is supported by calculating the ratio of the characteristic diffusion time in the pellets tpellet to the characteristic diffusion time in the bed tbed, a s estimated with the following expressions [6]: tbed -"

l b2ed ~ bed

Kn

(2)

2 De, bed

tP ellet "-

2 ~ pellet )~l pellet Kn 4 De, pore

(3)

The ratio between intrapellet (DKne,pellet) and interpellet (DKne, bed) diffusivities is represented by [6]" D e,r~ore Tbed d pore _ 3 (1 - 6bea )6pellet

D e,bKned

Z,

T pellet

d pellet

(4)

Taking lbed -- 12 mm, the ratio of the tortuosities Tbed/Tpellet ~ 1, F. bed -- 0.41, dpellet : 180 pm, and dpore- 37 nm we obtain: t bed

-0.21

t~e,,~,

(5)

which means t h a t the time scale of pellet diffusion is larger t h a n the time scale of diffusion on the scale of the bed. In Figure 5 CO and N2 diffusivities are given as a function of the square root of the t e m p e r a t u r e when the responses were interpreted without involvement of source terms. The behaviour of the inert N2 is similar as t h a t of Kr. However, the CO diffusivities deviate from the predictions for non porous pellets notably at lower temperatures. The lower values obtained indicate t h a t CO is adsorbed/desorbed by the support, as also could be concluded directly from the observed pulse responses, which were somewhat retarded compared to noble gas pulse responses. It means that an interpretation of the responses with Equation (1) without involvement of a source term is not allowed. CO2 pulsed over support material experienced a similar effect, as is shown in Figure 6. However, the 02 pulse responses behaved like an inert gas.

98

0.0010

t"~ t ' v ~ 4 I::

Dnp(CO/ N 0.0010 ~

{

2

1

....

0.0008

~

z~

~

A

0.0006

A

0.0004

0.0005

A

9

35 . . . . . . 17 . . . . . . . . .19 ....

COoverpo,~ S~ N~over porous SiO2 CD/ N2(t~ >>t~) 21

.

.

.

.

A

0.0002 C~5 ....

25

CO,overporousSiO.l (t~w . > > t ~ )

i , , , , I , , , , I ....

17

19

21

I ....

23

"

25

TI/2 [KI/2]

TI/2 [K1/2] F i g u r e 5. D i f f u s i v i t y of N2 a n d CO as a f u n c t i o n of t h e s q u a r e root of t h e t e m p e r a t u r e . T h e line r e p r e s e n t s t h e d i f f u s i v i t y (Dnp) in case of t h e n o n p o r o u s s i t u a t i o n (tpellet > > tbed).

F i g u r e 6. D i f f u s i v i t y of CO2 as a f u n c t i o n of t h e s q u a r e root of t h e t e m p e r a t u r e . T h e line r e p r e s e n t t h e d i f f u s i v i t y in c a s e of t h e n o n p o r o u s (tpellet > > tbed) s i t u a t i o n .

A reinterpretation of the responses for CO and CO2 was carried out with source terms taken into account, while the diffusivities were taken as for non porous pellets. The corresponding model parameters are shown in Table 3. The surface coverage shown concerns the value in between subsequent pulses and is almost equal to zero. The adsorption rate coefficient of CO, and the ad- and desorption rate coefficients of CO2 are more or less not activated, while desorption of CO is an activated process. Multipulses indicate t h a t the total surface concentration of active sites on the support is less t h a n one pulse size, i.e. avLit,c < 0.01 mol/m3c, which is confirmed by the values of avLt. Table 3. P a r a m e t e r s resulted from regressed pulse responses over the support. W

R co8

[K]

0co

kc~

323 373 423 473

15"10 .5 4.6"10 .5 9.2"10 .5 7.8"10 .5

2.9"10 1.6"10 6.5"10 9.0"10

4 4 3 3

k c~ 9.4 3.4 0.14 0.021

RC~ avLC~ 5.0"10 3 2.8"10 .3 3.9"10 .3 9.5"10 .4

0co2 4.9"10 .5 2.7"10 .4 1.1"10 .6 1.3"10 .6

kC~ 3.8"10 1.6"10 7.9"10 8.7"10

4 4 4 4

kc~ 60 1.3 75 96

avLC~ 3.6"10 .3 1.6"10 .3 2.0"10 .6 7.6"10 .6

3.3. The catalytic bed CO pulse responses from single pulse experiment over the catalyst were not detectable by the QMS, most likely due to an extremely low CO desorption rate. Figure 7 shows the effect of CO2 sorption on the support to be small compared to the sorption on the catalyst. The solid lines represent the results from using the model parameters. For 02 pulse responses two models were checked by regression: reversible dissociative adsorption and reversible molecular adsorption followed by an

99 Table 4. Parameters resulted from regressed pulse responses over the catalyst. T Ro% RCO% [K]

0o2

k~

323 345 373 423 473

3.9-10 .5 1.4"10 .5 3.3"10 .5 2.0"10 -5 . .

1.6.10 4.0"10 3.5"10 4.5"10 . .

k~ 2 2 2 2 .

avL~

6.2.101 1.5"10 2 1.3"10 2 5.2"10 2 . . .

0.51 0.26 0.17 0.15

0co2 6.4"10 .9 . . 1.3"10 .7 6.8"10 s 6.4"10 "s

kC~ .

24 . 24 12 8.6

kC~ .

1.8-10 . . 3.0"10 3.9"10 4.4"10

avLC~ 2 . 2 2 2

21 12 12 8.7

instantaneous irreversible dissociation. The last model (also represented in Table 2) appeared to be the best choice for predicting the 02 pulse response by the model parameters shown in Table 4. The irreversibility of the instantaneous dissociation is confirmed by the pulse response of the second multipulse (5 minutes after the first) indicating that the active sites stayed almost completely filled. Nevertheless, an additional (non diffusional) delay in the 02 single pulse response shape was observed, which has been interpreted as reversible adsorption of molecular oxygen. The parameters obtained from regression of the sorption models, implemented in Equation (1) over the catalyst are shown in Table 4 for CO2 and 02. The diffusivities used were based upon the situation of non porous inert pellets. 5 x l O ' 9 F.....-.x. . . . . . . . . . . . .

~[~k

LII\ k\

4x10"9 FI I\

--

3x10-9flI \ ~ 2x10-9~ ~

~ ---i~,~pp~,~.(~)~ - - modelling -~ __,.~~,~,.,j C02 pulseresponse ~ ~

lx10-9 O0 . . . . 015 . . .

1.0

1.5

2.0

t[s] F i g u r e 7. O u t l e t p u l s e s r e s p o n s e s of CO2 o v e r inert/support/inert and over inert/catalyst/inert represented by the dotted lines. The solid lines are

modelled

predictions

parameters from Table 3.

with

use

of

the

The surface coverage of both CO2 and 02 in between two succeeding single pulses was almost equal to zero. CO multipulses indicate a total surface concentration of active sites on the catalystwithavalue~176176176 18 mol/mc 3 was obtained at 323 K. This is in line with the value avLC~ obtained by regressing the CO2 pulse response at 323 K. This value is 1/3 of

the total surface concentration obtained with ASAP, being another indication of slow pellet diffusion. The avLC~ decreases, thus the effective overall number of active sites diminishes, with increasing T. This effect is caused by higher desorption rate coefficients at higher T. The avL~

values show a

similar temperature effect, but the absolute value is much smaller compared to avLC~ as almost all sites are covered with O adatoms due to the irreversibility of the instantaneous dissociation of adsorbed 02 as mentioned before. The obtained

100 adsorption rate coefficients are rather independent from T, which is in line with the expected absence of any activation. The desorption rate coefficients in Table 4 tend to increase with increasing temperature. The activation energies were determined as 27 and 25 kJ/mol for 02 and CO2, respectively. It is expected t h a t more accurate results would be obtained if diffusion on both the bed and pellet scale is incorporated in the interpretation of the responses.

4. C O N C L U D I N G REMARKS From TAP pulse responses with inert gases over porous catalyst it is concluded that only part of the intrapellet catalyst volume interferes with the pulse. Bed diffusivities are close to the values for non porous pellets. Based upon this result estimates were determined concerning adsorption/desorption rate coefficients for 02 and CO2 on Pt/SiO2, and for CO and CO2 on the silica support.

Acknowledgements The Netherlands Foundation for Chemical Research (SON) is acknowledged for financial support. The authors are grateful to Engelhard De Meern BV for providing the catalyst and support material. Raymond P.J. Pelsers is acknowledged for his contribution to the TAP measurements.

5. R E F E R E N C E S 1 2 3 4 5 6 7 8

J.T. Gleaves, J.R. Ebner, and T.C. Kuechler, Catal. Rev. - Sci. Eng., 30(1), pp. 49-116, 1988. J.P. Huinink, Ph.D. Thesis, Eindhoven University of Technology, Eindhoven, 1995. J.H.B.J. Hoebink, J.P. Huinink, and G.B. Marin, Appl. Catal. A, 160, pp. 139-151, 1997. T.A. Nijhuis, Ph.D. Thesis, Delft University of Technology, Delft, 1998. Y.J. Mergler, J. Hoebink, and B.E. Nieuwenhuys, J. Catal., 167, pp. 305-313, 1997. J.P. Huinink, J.H.B.J. Hoebink, and G.B. Marin, Canad. J. Chem. Eng., 74, pp. 580-585, 1997. NAG Fortran Library Manual, Mark 16, Numerical Algorithms Group Limited, Oxford, 1995. T. Engel and G. Ertl, Adv. Catal., 28, pp. 1-78, 1979.

6. MAJOR NOTATION av

Greek symbols

external surface catalyst / support area per m 3 catalyst / support

mc2/mra

concentration of gaseous component i diameter effective Knudsen diffusivity

tool/rag3 m

mg3/mrS

~f

vacant site on Si02

F

molar flow rate

vacant site on Pt

maximal molar concentration of

moYs mol]mr 2

*

Lt

a

adsorption

c

catalyst

C, d DK%

0i

surface sites

porosity fraction surface sites of component i

Sub- and superscripts

Rij

source term of component i for phase j

l/s

d

desorption

T

temperature

K

t

time

s

np p

non porous porous

x

axial reactor coordinate

m,

s

support

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

101

Transient kinetics of e t h y l e n e and carbon m o n o x i d e oxidation for a u t o m o t i v e exhaust gas catalysis: experiments and modelling. J.M.A. Harmsen, J.H.B.J. Hoebink, and J.C. Schouten Eindhoven University of Technology, Laboratory of Chemical Reactor Engineering, P.O. Box 513, 5600 MB Eindhoven, The Netherlands. phone +31 40 2472850; fax +31 40 2446653; E-marl [email protected]

Abstract The t r a n s i e n t oxidation of ethylene by oxygen over a commercially available Pt/Rh/CeO2/7-A1203 three-way catalyst is described. Experiments were carried out in a fixed-bed micro reactor with two separate inlets, enabling alternate feeding of ethylene and oxygen. The experimental conditions were chosen as to obtain intrinsic kinetics, i.e. in the absence of external and internal mass and heat transport limitations. Also, these conditions resemble the cold-start period of an Otto engine in a car. For these conditions, only total oxidation of ethylene to carbon dioxide and steam was observed. From the shape of the CO2 peaks, two types of adsorbed ethylene species seem to exist. The presence of ceria in the catalyst did not influence the ethylene oxidation significantly at the investigated conditions. A kinetic model, based on elementary reaction steps, is presented, which can be combined with the model for transient carbon monoxide oxidation on the same catalyst [1].

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

Most cars are nowadays supplied with a monolithic catalytic converter to remove pollutants in the exhaust gas as less harmful components. The socalled three-way catalyst allows simultaneous reactions of carbon monoxide, uncombusted hydrocarbons, and oxides of nitrogen towards carbon dioxide, water, and nitrogen. Especially at cold start conditions, the currently applied converter cannot eliminate all pollutants. Therefore the everexpanding road traffic causes an on-going demand for less pollution from automotive engines, as expressed by more severe legislation. For a further optimisation of the converter, a detailed understanding of the relevant kinetic processes, taking place simultaneously and interactively, is necessary. It requires modelling of intrinsic kinetics on the basis of elementary reaction steps, because, especially during catalyst warmup, the operating conditions change continuously, resulting in possible

102 alterations of an assumed rate-determining step. Also oscillations of the reactor feed composition are induced by the lambda controller, which may affect the time-averaged conversion, due to the non-linear kinetics. Research at Eindhoven University of Technology aims to produce an elementary step kinetic model, based on transient experiments, which will be able to predict the behaviour of complex exhaust gas mixtures on one type of commercial catalyst, Pt/Rh/CeO2/7-A1203, under transient conditions. Previously, the steady state kinetics of carbon monoxide oxidation [2] and ethylene oxidation [3] as well as the transient kinetics of carbon monoxide oxidation [1] have been reported for this catalyst. In this study, ethylene has been chosen as a representative of the m a n y different hydrocarbons in an automotive exhaust gas. Especially ethylene is used as a model compound for the hydrocarbons, which are easily oxidised. According to Impens [4], most abundant hydrocarbons are typically ethylene (25 mol% of the total hydrocarbons), acetylene and methane (both 20 mol%), and aromatics (toluene: 15 mol% and benzene 5 mol%). Moreover, ethylene has been shown to display similar oxidation behaviour as toluene and benzene [5,6]. Many authors studied the steady state oxidation of alkenes. Around the stoichiometric point usually a negative partial reaction order for the alkene and a positive order for oxygen is found with Pt and Pd catalysts [711], while only under a large excess of oxygen these orders are reversed [12,13]. For Rh catalysts a negative partial reaction order for oxygen and a positive order for ethylene was found [7,10,11]. Kinetic modelling based on elementary steps has been performed by Burch et al. [14] on a Pt/A1203 catalyst, for the steady state oxidation of propane and propene in the presence of nitrogen oxide and excess oxygen. For propene they concluded t h a t the catalytic surface was predominantly covered with hydrocarbon species, while for propane, the surface was mainly covered by oxygen adatoms. Steady state modelling, not based on elementary reaction steps, has been performed by Voltz et al. [8] for carbon monoxide and propene oxidation by oxygen and nitrogen oxide on a Pt/A1203 catalyst, and by Montreuil et al. [15] for propane and propene oxidation over a Pt/Rh catalyst. Both studies show inhibition by hydrocarbons and CO. The transient oxidation of hydrocarbons has also received broad attention. Muraki et al. [16] found t h a t the oscillation frequencies for m a x i m u m conversion of propene over Pt, Pd and Rh increased with increasing oscillation amplitude and temperature. According to Shinjoh et al. [9], the optimum frequency for the oxidation of propene and propane decreases with increasing temperature. Amon-Meziere et al. [17] investigated light-off behaviour of several hydrocarbons from C1 to C9. The general conclusion is that feed cycling decreases the light-off temperatures, because of less hydrocarbon inhibition [18]. Sant et al. [19] constructed a kinetic model, based on elementary reaction steps, for the transient oxidation of ethylene over a Pt/SiO2 catalyst. It contains two pathways, one leading directly to CO2 under oxygen rich conditions and one with CO as an intermediate product for oxygen poor conditions.

103 The influence of ceria for the oxidation of hydrocarbons is said to be negligible or negative below 573K [20-22]. Above 573K ceria can provide oxygen to enhance hydrocarbon oxidation [23]. This paper presents experimental results and a kinetic model, based on elementary reaction steps, for the transient oxidation of ethylene in the framework of automotive catalysis. Moreover, experimental data for the simultaneous oxidation of ethylene and carbon monoxide are reported in detail.

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

The experimental setup, used for the transient oxidation of ethylene and carbon monoxide via cyclic feeding, has been described in detail by Campman [24]. It consists of feed, reactor, and analysis sections. Feed section The feed section consists of two duplicate gas blending systems to generate two feed streams with different compositions. For each component, a blending system contains an electromagnetic valve and a thermal gas mass flow controller. By means of four magnetic valves, the two gas feeds are alternated over the reactor with an adjustable frequency up to 10 Hz. Reactor section The reactor section consists of a tubular preheater and a stainless steel fLxed bed reactor (type 316). A cross section of the reactor is depicted in figure 1. The two feed lines remain separated in the preheater upto the catalyst bed and are sealed off by sapphire beads, unless the beads are pressurised. This allows only one gas flow to be led over the reactor. The catalyst bed has a length of Figure 1. Detailed view of the fixed 15 mm and a diameter of 13 mm. Sample bed reactor. chambers, immediately above and below the catalyst bed are connected via capillaries to the on-line mass spectrometer for analysis. A thermocouple tube allows monitoring the axial temperature profile in the catalyst bed, in order to assure isothermal experiments. Radial temperature profiles are monitored by a thermocouple located on the outside reactor wall. This latter thermocouple is used as input for a PID controller, which sets the reactor temperature by controlling two infrared radiators. The reactor pressure is manually controlled by a spring loaded backpressure valve and measured downstream the reactor.

104

Analysis section

Table 1 Range of experimental conditions

/K 393 - 443 The analysis section consists of Temperature total pressure /kPa 110 an online gas chromatograph (Carlo p0 C2H4 /kPa 0.0 - 0.15 Erba I n s t r u m e n t s GC 8340) for pO o2 /kPa 0.0 - 0.55 steady state m e a s u r e m e n t s and an pO co /kPa 0.0 - 0.50 online quadrupole mass spectrofrequency /Hz 0.1 - 0.25 meter (VG Sensor lab 200D) for duty cycle /% 10 90 steady state and transient meaWcat /10 "3 kgcat 1.01 surements. The mass spectro- meter cat dilution /m3inert 0.48 m e a s u r e d the He, C2H4, CO, 02, and m'3inert+cat CO2 concentrations both at the total flow /mol s -1 5.6 10 .3 reactor inlet and outlet during the cyclic feeding experiments. Analysis can be done with a frequency of 120/n Hz, where n is the n u m b e r of masses analysed. To distinguish between CO and C2H4 also the ethylene fragmentations at m/e=27 and m/e=26 were measured. Daffy calibration of the mass spectrometer was required for obtaining quantitative data.

Catalyst The Pt/Rh/CeO2/y-A1203 three-way catalyst as used for coating monoliths, was provided by Degussa A.G. as a powder with a mean particle diameter of 12 ~m. The powder was pressed into pellets, crushed and sieved until the particle diameter was between 0.11 and 0.15 mm. The reactor was filled with 1.01 g of catalyst and 1.46 g of inert c~-A1203 dilution with a pellet diameter range 0.15 - 0.21 mm. The catalyst p r e t r e a t m e n t at 773 K consisted of oxidation by oxygen during one hour followed by reduction with hydrogen for half an hour. More details about the catalyst properties were reported elsewhere [1-3]. Reproducible experimental data could be obtained by keeping the catalyst continuously under t r a n s i e n t conditions, alternating between a rich feed and a lean feed, hence avoiding the slow deactivating oxidation of the noble metal particles [25-27].

3.

E X P E R I M E N T A L R E S U L T S AND D I S C U S S I O N

The ranges of the experimental conditions, which were chosen to approximate an automotive e x h a u s t gas, are depicted in Table 1. Two types of cyclic feeding experiments have been carried out: 1) experiments with ethylene in helium in one feed and oxygen in helium in the other feed, and 2) experiments, like type 1), with an a m o u n t of carbon monoxide added to the ethylene in helium feed. A typical inlet signal is shown in figure 2a. Figure 3a shows an inlet signal with 0.1 kPa carbon monoxide added to the ethylene in helium flow. As can be seen, a step change is completed within 0.5 s, by using the fast valves for switching between the feeds. The corresponding outlet signals at a t e m p e r a t u r e of 423 K can be seen in figures 2b and 3b. When oxidising ethylene only (fig. 2b), CO was not observed as an intermediate product. If it is formed at all, as reported by Sant and Wolff [19], it does not desorb or is very rapidly converted to carbon dioxide.

105

oxygen

0.5

>o

-~ 0.4

~" 0.4

_o

0.3

8 o.3

_

0.2

~

ethylen:

0.2

ethylene

0.1

0

0

2

4

6

time (s)

8

10

12

•q•

0.6

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

0.4

2

4

2

ii

0.5

-x 12

0.45

oxy]e~

0.35 =~0.3

"~ o~'O. 4

0"120~ ~

carb~ ~

0.04 E

0.25 t0~

'

0.2

(3b)

o.o8~ ~ t ~o.2

A 0

10

0.16~,

0.3

0.1

timeIs)

-

8

0.6 ...................................................................................................

o0"3

--~0.2 E

6

0.2

oxygen

0.5

0

0.15 0.1

o.1

0.05

4

6

time (s)

8

lO

12

0

0

o

2

4

6

time (s)

e

io

12

Figure 2a,b: reactor inlet (a) and outlet (b) signals for the transient oxidation of ethylene at a frequency of 1/10 Hz, a temperature of 423 K, and a total pressure of 110 kPa. Partial inlet pressures for oxygen and ethylene are respectively 0.55 kPa and 0.15 kPa. Figure 3a,b: the same as figure 2, with 0.1 kPa carbon monoxide in the same feed as ethylene. Note that the carbon dioxide fractions are depicted on the right hand y-axis in figures 2b and 3b.

When switching from ethylene to oxygen (e.g. at time = 6 s) a peak of CO2 arises in the outlet signal, corresponding to the amount of ethylene adsorbed on the catalyst surface, right before switching. This peak consists of a very steep ascent, and a slower descend, that abruptly retards, at about 8 s, before decreasing to zero. This indicates the presence of two types of ethylene surface species, one which is oxidised very fast, and a second, which is oxidised much slower. It is reasonable to assume that the second type of species can be formed from the first, using vacant sites that arise at the catalytic surface. The fast reacting species is ascribed to n-ethylene, adsorbed on one single active site, while the slow reacting, more stable species would be di-~-ethylene, adsorbed on two noble metal sites [28-32]. This is in line with the steady state results of Nibbelke et al. [3], who found a C2H4 adsorption rate that is first order in the fraction of vacant sites (n-ethylene) and an activation energy for ethylene desorption corresponding with di-~-ethylene. As a large number of vacant sites are available in a steady state, u-ethylene can be very rapidly converted into the more stable, thus less reactive di-c-ethylene. The CO2 peak at the switch from oxygen to ethylene (e.g. at time = 1.5 s) corresponds with adsorbed O adatoms on the catalyst. The peak has a much smaller surface area, as six O adatoms are required for the oxidation of one ethylene molecule. Therefore, all oxygen on the surface will be consumed very rapidly by the adsorbing ethylene. As the height of the CO2 peak is much smaller t h a n shown by

106 the peak at 6 s, it is assumed that the adsorbing u-ethylene is almost immediately converted to the slowly oxidising di-(~-ethylene. This will be possible, because many vacant sites become available due to the reaction stoichiometry. For the investigated conditions, the time-averaged ethylene conversion increases with increasing temperature and oscillation frequency. An optimal duty cycle was found at 50%. The same experiments, performed with a Pt/7-A1203 catalyst, gave similar results as in figure 2, thus excluding any significant influence from ceria, present in the currently used catalyst. This is in line with reported results [20-22]. When carbon monoxide is added to the ethylene containing feed, the oxidation of ethylene becomes inhibited, resulting in a faster relaxation of the ethylene outlet signal and a larger CO2 peak, see figure 3b. Applying higher CO partial pressures t h a n 0.1 kPa, leads to a strong increase of the inhibition, until hardly any ethylene is converted at a CO partial pressure of 0.3 kPa. Clearly, CO adsorbs faster onto the noble metal, blocking sites for ethylene. Comparing figure 2b with 3b, a number of differences can be seen. After switching from rich to lean conditions, the oxygen signal in figure 3b displays a break-through peak followed by a dip, which is typical for the CO oxidation at these t e m p e r a t u r e s [1]. Also the ability of CO to adsorb on an oxygen covered noble metal site, can be seen (e.g. time = 3-4 s), as the CO conversion equals 100% during 1.5 seconds after the rich feed is led over the reactor. Clearly, ethylene lacks this ability; its signal is similar to figure 2b. The carbon dioxide peak at time = 8 s displays the same type of behaviour as in the ethylene only oxidation, however the peak is much higher and more narrow. Especially the long tail, which was characteristic for the ethylene oxidation, is much smaller now. This is probably due to a large inhibition of the formation of the slow reacting di-c~-ethylene. Hence the surface is covered mostly with the fast reacting species CO and u-ethylene, which are each adsorbed on one active site only. The result is t h a t more molecules are converted in less time. The CO2 peak (at 2-3 s) consists, as in the ethylene only experiment, only of a very small ascent followed by a slow descent to zero. The main reasons for its increased peak area compared to figure 2b are stoichiometry and the availability of the bifunctional ceria path for CO oxidation [1,2]. The shape of this peak is however completely different from the experiments with CO only, which could indicate t h a t the CO oxidation becomes inhibited by refractory di-(~-ethylene. Like in the CO only oxidation [1], the direct adsorption of CO on an oxygen-covered site is the main path here, resulting in a long period of 100% CO conversion. The vacant sites, which are formed t h a t way, will partly become occupied by ethylene, which retards the reaction between adsorbed CO and adsorbed oxygen. This, and reaction stoichiometry, results in a smaller CO2 peak then in CO only oxidation. From the slow relaxation of the CO signal towards the inlet level, it can be concluded t h a t the tail of the CO2 peak is mainly due to CO oxidation, resulting in a surface, which is primarily covered with CO.

4. K I N E T I C M O D E L L I N G

From the discussion above a reaction scheme is suggested for the oxidation of ethylene. This scheme is displayed below.

107 Reaction steps" 1. 02 (g) + 2*

20*

2.

C2H4(g) +

*

Cell4*

3.

C2H4"

+

*

C2H4"*

4.

C2H4"

+

60*

2CO2 + 2H20 + 7*

5.

C2H4"* +

60*

6.

CO2

t

+

~,

~

2COe+2H20+8*

~

CO2t

The adsorption of oxygen (step 1) involves molecular adsorption followed by i n s t a n t a n e o u s dissociation in line with results by Nibbelke et al. on the same catalyst [1]. Ethylene is assumed to adsorb reversibly as n-ethylene (step 2). The ~ethylene on the surface can be converted to the di-(~ form when vacant active sites are available (step 3). Both adsorbed ethylene species can be oxidised directly to carbon dioxide and water, using six oxygen adatoms (steps 4 and 5). Steps 4 and 5 cannot be considered as single elementary reaction steps, but are assumed to consist of a n u m b e r of subsequent elementary reaction steps, all being very fast, except an oxygen assisted dehydrogenation of ethylene. The reversible adsorption of carbon dioxide onto the catalytic support (step 6) has to be considered as well under t r a n s i e n t conditions, as shown in a previous study of t r a n s i e n t carbon monoxide oxidation [1].

5. C O N C L U S I O N S

The t r a n s i e n t oxidation of ethylene was studied. Experiments, under conditions representing the cold start of an Otto motor, indicate the presence of two types of adsorbed species, which are ascribed to ~- and di-(~-ethylene. The former species can be oxidised rapidly, while the latter one is r a t h e r refractory. No intermediate products were detected. Ceria does not appear to have a significant influence on the ethylene oxidation at the investigated conditions. Addition of carbon monoxide to ethylene inhibits the oxidation of ethylene, especially the formation of the refractory. Formation of di-(~ ethylene retards the reaction between adsorbed CO and oxygen adatoms.

Acknowledgements Financial support for the study was given by the Dutch Technology Foundation (STW). The authors are grateful to Degussa A.G. for providing the catalyst. Jeroen K l u y t m a n s is gratefully acknowledged for his contribution to the experimental work.

108 6. R E F E R E N C E S

1. R.H. Nibbelke, A.J.L. Nievergeld, J.H.B.J. Hoebink, G.B. Marin, Appl. Catal. B Environmental, accepted for publication 2. R.H. Nibbelke, M.A.J. Campman, J.H.B.J. Hoebink, G.B. Marin, J. Catal. 171, (1997) 358 3. R.H. Nibbelke, R.J.M. Kreijveld, J.H.B.J. Hoebink, G.B. Marin, Proc. 4rd Int. CAPOC 1,(1997) 139 4. R. Impens, in "CAPOC" (A. Crucq, A. Frennet, Eds.), Elsevier, Amsterdam (1987) 11. 5. G. Mabilon, D. Durant, Ph. Courty, in "CAPOC III" (A. Frennet, J.-M Bastin, Eds.) Elsevier, Amsterdam (1995) 775 6. J.M. Bart, A. Pentenero, M.F. Prigent, ACS Symposium Series, 495, (1992) 42 7. N.W. Cant, W.H. Hall, J. Catal. 16 (1970) 220 8. S.E. Voltz, C.R. Morgan, R. Morgan, D. Liederman, S.M. Jacob, Ind. Eng. Chem. Res. Dev., 12 (1973) 294 9. J.R. Hawkins, S.E. Wanke, Can. J. Chem. Eng., 57 (1979) 621 10. Y.-F.Y. Yao, J. Catal. 87 (1984) 152 11. H. Shinjoh, H. Muraki, Y. Fujitani, Appl. Catal., 49 (1989) 195 12. L. van de Beld, M.P.G. Bijl, A. Reinders, B. van der Werf, K.R. Westerterp, Chem. Eng. Sci., 49 (1994) 4361 13. C.G. Vayenas, B. Lee, J. Michaels, J. Catal., 66 (1980) 36 14. R. Burch, T.C. Watling, Proc. 4rd Int. CAPOC 1, (1997) 69 15. C.N. Montreuil, S.C. Williams, A.A. Adamczyk, SAE Technical Paper Series 920096 (1992) 16. H. Muraki, H. Shinjoh, H. Sobukawa, K. Yokota, Y. Fujitani, Ind. Eng. Chem. Prod. Res. Dev. 24 (1985) 43 17. I. Amon-Meziere, F. Castagna, M. Prigent, A. Pentenero, SAE Technical Paper Series 950932 (1995) 18. P.L. Silveston, Cat. Today 25 (1995) 175 19. R. Sant, D.J. Kaul, E.E. Wolf, J. AIChE, 35 (1989) 267 20. L. Padeste, A. Baiker, Ind. Eng. Chem. Res. 33, (1994) 1113 21. A.F. Diwell, R.R. Ramaram, H.A. Shaw, T.J. Truex, in CAPOC II,(A. Crucq, Ed.) (1991) 139 22. B.I. Whittington, C.J. Jiang, D.L. Trimm, Cat. Today 26 (1995) 41 23. A. Trovarelli, Cat. Rev.-Sci. Eng. 38 (1996) 439 24. M.A.J. Campman, Ph.D. dissertation, Eindhoven University of Technology, (1996) 25. L.M. Carbello, E.E. Wolf, J. Catal., 53 (1978) 366 26. R. Burch, M.J. Hayes, Journal of molecular catalysis: A Chemical, 100 (1995) 13 27. L. Hiam, H. Wise, S. Chaikin, J. Catal. 10 (1968) 272 28. J.-F. Paul, P. Sautet, J. Phys. Chem., 98 (1994) 10906 29. S. B. Mohsin, M. Trenary, H.J. Robota, J. Phys. Chem., 92 (1988) 5229 30. P. Berlowitz, C. Merigis, J.B. Butt, H.H. Kung, Langmuir, 1 (1985) 206-212 31. D. Velic, R.J. Levis, J. Chem. Phys. Vol. 104 No. 23 (1996) 9629-9639. 32. G. Szulczewski, R.J. Levis, J. Am. Chem. Soc., 118 (1996) 3521

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

109

D y n a m i c s o f N O adsorption and t r a n s f o r m a t i o n over s u p p o r t e d Pt catalysts for t h e t r e a t m e n t o f l e a n b u r n e n g i n e e m i s s i o n s G.E. Arena, 1 A. Bianchini, 2 G.Centi,1, 2. F. Vazzana, 1 and P. Vitali 2 J Dip. Chimica Industriale, Univ. of Messina, Italy. 2 Dip. Chimica Industriale e dei Materiali, Univ. of Bologna, Italy. Abstract

The role of the NO adsorption and transformation over a model Pt(1%)/A1203 catalyst for the reduction of NO by propene/O 2 is discussed on the basis of NO thermodesorption and transient reactivity data. The experimental results indicate that NO strongly chemisorbs over the Pt oxidized surface and this inhibits the Pt reactivity towards hydrocarbon oxidation, but the NO then progressively migrates over the alumina surface from which it can desorb in the form of NO2. At the temperature of maximum activity in NO reduction, the rate of migration is slow and thus the hydrocarbon reacts with the adsorbed nitrogen oxide species over the metal surface, whereas at higher temperature the rate of migration is faster leaving the oxidized metal surface free for hydrocarbon combustion. INTRODUCTION Lean burn or diesel engines are characterized by lower fuel consumption and C O 2 emissions than current engines operating at a stoichiometric air/fuel ratio, but the presence of O2 in their emissions prevents the use of current "three way" type catalysts. It is thus necessary to develop catalysts able to reduce NOx to N 2 in engine emissions containing oxygen. In recent years, considerable research activity has been focused on the development of transition metal containing microporous materials for the reduction of NO in the presence of oxygen [ 1,2], but due to their low activity and stability, interest has now been focused on the possibility of modifing conventional supported noble metal catalysts in order to improve their activity and temperature operative window in NO reduction to N2 in the presence of 02 and hydrocarbons. These catalysts show a sharp volcano-type temperature range of activity centred at about 200-250~ in the conversion of NO in the presence of O2/hydrocarbons (propene, for example) [3]. Above this temperature, NO conversion to N 2 decreases and NO oxidation to NO2 increases in coincidence with the sharp increase in the conversion of the hydrocarbon to CO2. Widening of the operative window requires limiting the direct oxidation of the hydrocarbon by oxygen in favour of its role as a selective reductant of nitrogen oxides. This in turn requires understanding the catalytic phenomena and the dynamics of adsorbed species which determine the catalytic behavior. Recent evidence indicates that the catalytic behavior is not only associated with a localized reaction at the active sites, but the mobility of adspecies over Salita Sperone 31,98166 Messina, Italy. Fax: +39-090+391518, e-mail: [email protected]

110 the surface also plays a significant role. Researchers at Toyota [4] have recently developed a catalytic system based on the concept of "NOx storage-reduction". The engine is forced to oscillate in the air/fuel ratio between lean and rich conditions. During the lean period (presence of O2) the NO is oxidized to NO 2 on Pt particles and then is assumed to move over the alumina surface (the support) to sites (usually BaO) where it is stored in the form of nitrates. During the rich period back-spillover of the nitrate species occurs which again migrates to the Pt surface to be reduced by the hydrocarbons, CO and H2 present in the exhaust stream under rich conditions. Although the details of the reaction mechanism over Toyota-type catalysts are not well understood, the data already available [4] suggest the fundamental role of the mobility and dynamics of adspecies over these catalysts. However, very limited data are available in the literature. The scope of this work is to analyze the question of the role of the dynamics of surface processes on the catalytic behavior of a Pt(1%)/A1203 catalyst used as a model reference for conventional supported noble metal catalysts. In subsequent papers, the influence on the surface processes of the presence of additional elements with nitrate storage ability will be analyzed. EXPERIMENTAL The Pt(l% wt.)/A1203 catalyst was prepared using commercial ~/-A1203 (Rhone Poulenc 531; surface area 108 m2/g) to which platinum was added by incipient wet impregnation using an aqueous solution of H2PtC16. After drying and calcination up to 550~ the catalyst was reduced at 400~ for 12 h with a flow of pure hydrogen and then reoxidized in mild conditions. Steady-state, thermodesorption and transient reactivity data were obtained using a quartz fixed bed reactor loaded with 0.1 g catalyst in the form of micropellets (40-60 mesh) and a flow rate of 12 L/h. A thermocouple in contact with the catalytic bed allowed control of the reactor temperature. The inlet and outlet reactor compositions were monitored on line by a mass quadrupole apparatus. The mass intensity data were corrected to consider mass overlap. RESULTS

Steady-state Activity The steady-state activity in NO conversion of model Pt(I%)/A1203 samples was studied using propene as the reductant (0.05% NO, C3H6/NO = 2, 5% 02, remainder He). The catalyst shows in agreement with literature data [3] a sharp maximum in the conversion of NO (48%) centred at about 240~ (selectivity to N2 of about 60%), in correspondence with the sharp increase in propene conversion which reaches 100% at about 250~ At temperatures above 240~ the conversion of NO decreases (20% at 300~ and 13% at 350~ and the formation of NO2 increases significanty.

Thermodesorption Tests The effect of oxygen on the strength of the nitrogen oxides bound with catalyst surface sites was studied by thermodesoption experiments. In these tests, the catalyst is first usually conditioned at 500~ (2h) in a helium flow to clean the surface and then after cooling down to 100~ still in the He flow, pretreated with a flow of 02 (1% or 2% in He), H2 (5% in He) or pure helium for 30 rain. After the pretreatment, NO is chemisorbed at 100~ (0.1% in He) for 30 rain and then after switching the flow to pure helium or 02 (1%) in helium, the catalyst

111 temperature is increased linearly (20~ other temperature ramps were also evaluated, but the results were not significantly different) up to 500~ monitoring the desorption of NO or the formation of N2, N20 , NO2, 02 and H20 by on line mass quadrupole spectrometer. The results are summarized in Figure 1. In all cases, the products of reaction were below the detection limit and thus only the thermodesorption curves for NO are reported. If not specifically indicated, the thermodesorption curves refer to data obtained using helium as the carier gas. When NO is chemisorbed over the catalyst pretreated only in He, two clear distinct peaks occur centred at about 180~ and about 350~ The second peak is also associated with weak oxygen desorption suggesting the decomposition of a nitrate-like species. Worth noting is that this maximum catalyst activity in the reduction of NO by propene/O2 occurs at temperatures when desorption of the NO species associated with the lower temperature peak is nearly complete. When the catalyst is pretreated with O2/He, the intensity of the second high temperature peak does not change consideraby, but the intensity of the peak at the lower temperatures varies. Its intensity depends on the concentration of 02 in the pretreatment step. When instead the pretreatment in made with H2/He, the low temperature peak is not markedly influenced, but the high temperature peak disappears. Using a reactive flow during the thermodesorption (1% O2 in He), independenly of the type of pretreatment, a single type of thermodesorption curve is observed characterized by the absence of the low temperature peak and the weakening and shifting to higher temperatures (400~ of the higher temperature peak. However, the thermodesorption of nitrogen oxides is not complete, because when pulses of Hz/He at 500~ are sent after the end of the thermodesorption tests, the presence of residual nitrogen oxides over the catalyst surface is found. Thus when 02 is present in the carrier gas during the thermodsorption runs, the adsorbed nitrogen oxides convert to very strongly bound species which do not desorb even at 500~

300

250

CO

9

d.

o.. 200

dZ ~- 150 o ..~ ~ 100 o t-r

o

o

50

vo

V

v

~~

1O0

9

d.-~ 9 ~

v,,

200

o ,

9 9

o ~

~D

.,~r ~

300 Temperature, ~

~

"4~,..,'o 9

400

I

500

Fig. 1 NO thermodesorption curves in a helium flow (a -~ d) and in 1% 02 in a helium flow (e). NO adsorption (30 min): (a,e) 0.1% NO/He, (b) (0.1% NO + 1% O2)/He, (c) (0.1% NO + 2% O2)/He, and (d) 0.1% NO/He on the catalyst pretreated at 100~ with a flow of 5% H2/He for 30 min.

112

Fig. 2 O2-step change experiments at 240~ in a flow of (a) 0.1% NO in He and (b) about 0.1% C3H 6 in He in the presence or absence of about 0.1% NO.

Transient Oxygen Step-mode Tests The effect of oxygen on the transient reactivity of the catalyst was studied in the 240-350~ temperature region (range where catalyst activity in NO reduction by propene/O2 starts to decline) using rectangular step-type changes in oxygen concentration (from 0 to about 2%02 and then, after from 2 to 10 min, back to 0% O2) in a main feed of (i) 0.1% NO in He, (ii) about 0.1% C3H 6 in He and (iii) about 0.1% NO + 0.1% C3H6 in He. The results obtained at 240~ are reported in Figure 2. The catalyst pretreatment was desorption in a He flow up to 500~ followed by cooling the catalyst up to the reaction temperature still in the He flow. Using a main feed of 0.1% NO in He (Fig. 2a) immediately after addition of oxygen a rapid decrease in NO outlet concentration occurs which, however, does not correspond to the formation of other detectable nitrogen-containing species and thus indicates a fast strong adsorption on the catalyst. The adsorption is complete in about 1-2 min, although the NO outlet concentration remains slighly lower than the inlet value due to the formation of oxidized nitrogen oxide species which remain adsorbed over the catalyst surface. After switching off the O2 additions, a very small desorption of NO occurs for about 1 min and then the NO concentration returns to the starting value. Using a main feed of 0.1% C3H 6 in He (Fig. 2b) the hydrocarbon outlet concentration rapidly decreases to nearly zero, ie. complete hydrocarbon conversion. After the oxygen is switched off, the hydrocarbon conversion similarly returns quickly back to the starting value. When about 0.1% NO is present in the main feed together with about 0.1% propene (Fig. 2b), the hydrocarbon outlet concentration initially decreases quickly, but then after about 1 min slightly increases again reaching a constant value after about 6 min. Propene conversion in these conditions is about 40% with respect to 100% in the absence of cofed NO, clearly indicating the inhibition of the surface reactivity in the presence of NO. Interestingly, the profile of the NO outlet concentration during cofed experiments very closely follows that of the hydrocarbon. With respect to the shape in the absence of cofed propene (Fig. 2a), the negative peak is less accentuated and the return to a constant concentration value is much slower. Worth noting is that the NO conversion is higher (around 30%) than in the absence of propene,

113

Fig. 3 Sequence of O2-step change experiments at 240~ (a) and 300~ (b) in a flow of 0.1% NO in He. Note: the values of the O2 conc. are divided by 20 and the value of NO 2 in Fig. 3a is multiplied by 5. although due to the overlap of CO and C O 2 peaks on those of N2 and N20, respectively, it was not possible to evidence to which products NO converts. NO2 formation', however, was negligible. The analysis of the thermodesorption of NO (using He as carrier gas) after the transient 02step change tests of Fig. 2 indicates that different types of species remain adsorbed on the

114 catalyst depending on the presence or not of cofed propene. While in the absence of C3H6, a broad peak centred between 350-400~ is observed, when propene is cofed together with NO a weaker peak centred between 250-300~ is detected. This indicates that the NOx species desorbing at the higher temperature is probably consumed by the reaction with propene. A sequence of O2-step change experiments in a main feed of 0.1% NO in He also provided usefull information. The results for the reaction temperatures of 240~ and 300~ are summarized in Fig.s 3a and 3b, respectively. In this sequence of rectangular changes in 02 concentration, oxygen was added for about 2 min followed by about 5 min without oxygen. At 240~ (Fig. 3a), the shape of the NO change in concentration during the first O2-step change follows the trend already described for Fig. 2a. In this case the formation of NO2 (only product of conversion of NO detected) is also reported. Note, that for clarity the NO2 outlet concentration was multiplied by a factor of 5. The shape of N O 2 response is the opposite of that detected for NO adsorption. Initially, immediately after addition of O2 (where NO conversion is maximum) the NO2 formation is nearly zero and then it gradually increases even when the NO adsorption has nearly stopped. After the oxygen is switched off, the NO2 concentration decreases to zero, but very slowly. The shape of its peak therefore is largely asymmetrical. In the consecutive rectangular O2 changes, the shape of the N O 2 peak remains similar, but its formation increases progressively. On the contrary, the intensity of the initial adsorption of NO decreases in magnitude, although this peak is still present in the consecutive rectangular 02 changes. The magnitude of the initial adsorption of NO depends on the time between the two consecutive rectangular O2 changes. As the time passed increased, decrease in the initial adsorption of NO is less accentuated. At higher temperatures (Fig. 3b), ie. in conditions after the maximum in NO conversion to N2 in steady-state conditions, there are relevant differences in the responses, although the general features remain similar. The following differences can be noted: i) N O 2 formation is much greater (the multiplying factor is no longer necessary) and the form of the peak also changes. After the initial increase in its concentration, there is a second further increase in its concentration near the end of the rectangular 02 change. The slow desorption during the period of absence of O2 is still present. ii ) The initial adsorpion of NO after adding O2 is still present, but does not change considerably between the various rectangular 02 changes. Furthermore, after the initial adsorption, the outlet concentration of NO decreases, but with a rate much slower than at the lower reaction temperature. iii ) After the 02 is switched off, the slow desorption of NO no longer occurs, but instead NO continues to be adsorbed/converted for about 1-2 min.

Transient Hydrogen Step-mode Tests Reported in Figure 4 is the effect at 240~ of consecutive rectangular H 2 changes in a main pure He feed, but after the tests of consecutive 02 changes in a main NO/He feed at the same temperature (Fig. 3a), and after having removed gas and weakly adsorbed species with a helium feed for about l h. The reactivity towards H2 of adsorbed species formed during the sequence of consecutive 02 changes is thus analyzed. The addition of hydrogen causes the partial desorption of NO from the catalyst, reasonably due to the reduction of oxidized and strongly adsorbed species over the catalyst surface. The oarallel formation of water (not reported for clarity) was also detected. The shape of the

115

Fig. 4 Sequence of H2-step change experiments at 240~ in a flow of pure He after the sequence of O2-step change experiments at the same temperature (Fig. 3a). Note: the values of H2 conc. are divided by 200. peak of NO during the rectangular H 2 change is very close to that of N O 2 formation during rectangular O2 changes (Fig. 3). Of particular interest is the formation of N 2 during these H2-step change experiments (Fig. 4). During the first rectangular change, N2 initially forms very quickly and then rapidly decreases to a nearly constant value followed by another decrease when the H 2 additions are switched off. However, after some induction time (around 1 min), N 2 formation starts to increase again forming a large broad peak when H 2 is no more present in the feed. During the second rectangular change in H 2 concentration, N2 does not form, but instead N 2 forms again during the period in the absence of H2. The intensity of the broad peak decreases along the consecutive changes in H 2 concentration. DISCUSSION The key feature of supported noble metal catalysts in the conversion of NO by C3H6/O 2 is the presence of a sharp maximum in the conversion of NO to N2 centred at about 200-250~ The results of the present experiments allow an explanation of this effect to be suggested. O2 step-change experiments feeding C3H6/He or (C3H6+NO)/He (Fig. 2) clearly showed that the catalyst activity in propene oxidation at 240~ [temperature of the maximum in NO conversion to N2] is deactivated by the cofeeding of NO. Reasonably, the effect can be attributed to the formation of strongly bound NOx species over the Pt surface. In agreement, thermodesorption experiments (Fig. 1) showed the presence of a nitrate-like species which desorbs only at high temperature. This species can be eliminated when the catalyst is pretreated in H 2 in mild conditions (100~ (Fig. 1, curve d), reasonably due to the reduction of the oxidized surface of Pt which cannot be cleaned up during the initial treatment in He flow at 500~

116 At the same temperature (240~ 02 step-change experiments feeding NO/He (Fig.s 2a and 3) indicate the presence of a very fast initial adsorption followed by a second slower process. Using a sequence of rectangular changes in 02 concentration (Fig. 3) it is possible to suggest that the two step process reasonably derives from an initial strong adsorption of NO on the oxidized metal surface, followed by a slower process of migration of the oxidized nitrogen species towards the alumina surface where it remains adsorbed and partially desorbs in the form of NO2 above a certain surface concentration. In fact, NO2 does not form immediately after addition of O2, but gradually increases. In particular its formation continuously increases in the sequence of rectangulat changes in O2 concentration (Fig. 3a). Furthermore, in H2 step change experiments (Fig. 4) after the O2-step change experiments of Fig. 3a, it is possible to evidence that during the first rectangular change H 2 quickly reduces the nitrogen oxides species present over the metal, thus cleaning its surface. In these conditions, spillover hydrogen species form which migrate to the alumina surface. Otherwise it is not possible to explain why a second large peak of N 2 is detected after some induction time after the switching off of H2 in the main He feed. During consecutive rectangular H2 changes, the nitrogen species over the metal surface are no longer present (thus N 2 does not form), but instead spillover hydrogen forms again which takes some time to accumulate over the surface to be able to reduce the nitrogen oxide adspecies to N2. Reduction of N O 2 like adspecies (present over the alumina surface) to NO is instead easier and thus can occur within the time of the rectangular H 2 change, but due to the necessity for the formation and migration of spillover hydrogen species, the NO formation increases during the rectangular H2 change. Note that the NO does not decrease immediately to zero after the H2 is switched off, but still forms especially in correspondence to the induction time before the start of N2 formation. Being similar the formation of NO2 during the 02 step-change experiments (Fig. 3a), ie. NOx species migrate from the metal to alumina surface before desorption, the shape of NO 2 (Fig. 3a) and NO (Fig. 4) responses is similar. At higher temperature (300~ ie. above the maximum in NO conversion to N 2 by propene/O2, the process of surface migration of oxidized nitrogen oxides from the metal to the alumina surface is faster. This explains the various differences in the responses between the two reaction temperatures (Fig. 3) as noted before. However, the oxidized nitrogen oxide species also inhibit the reactivity of chemisorbed oxygen over the platinum metal surface towards total combustion of the hydrocarbon, favouring, on the other hand, the reaction of the hydrocarbon (reasonably, activated by the metal surface) with the oxidized nitrogen oxides, first step of their selective reduction to N2. NO has thus an effect of self-moderation of the platinum activity towards total hydrocarbon oxidation, thus enhancing its role as a selective reductant, but due to the tendency of oxidized nitrogen oxides to migrate towards the support surface the effective role is determined by the rate of surface diffusion. This indicates that control of this factor can be a key to improving performances of supported noble metals and especially widening their operative windows in the reduction of NO. REFERENCES 1.

(a) M. Iwamoto, H. Hamada, Catal. Today, 10 (1991) 57. (b) M.Shelef, Catal. Rev.-Sci. Eng., 95 (1995) 209. 2. G. Centi, S. Perathoner S., Appl. Catal. A, 132(1995) 179. 3. (a) R. Burch, P.J. Millington, Catal. Today, 26 (1995) 185. (b) R. Burch, P.J. Millington, A.P. Walker, Appl. Catal. B, 4 (1994) 65. 4. N. Miyoshi, S. Matsumoto, K. Katoh, T. Tanaka, J. Harada, N. Takahashi, K. Yokota, M. Sugiura, K. Kasahara, SAE Paper 950809 (1995).

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

117

Effect of C a t a l y s t D e a c t i v a t i o n on t h e D y n a m i c s o f C y c l i c Reactive Processes Daniel O. Borio*, Noemi S. Schbib* and Jorge E. Gatica § * Planta Piloto de Ingenieria Quimica (UNS-CONICET), Camino La Carrindanga, Km. 7, Casilla de Correo 717. 8000 Bahia Blanca, ARGENTINA. (FAX: 54-91-861600- E-mail: [email protected]) § Department of Chemical Engineering, Cleveland ~tate Umversity, 1960 East 24 th Street- SH455, Cleveland, Ohio 44115-2425, USA. (FAX: 216-687-9220 - E-mail: [email protected])

Abstract The transient behaviour of a set of industrial fixed-bed catalytic reactors is presented in this paper. These reactors, used for the dehydrogenation of 1-butene into 1-3 butadiene, operate in reaction-regeneration cycles. Unlike previous contributions, the influence of catalyst sintering on the reactor operation is analysed in this work. The dynamic model presented includes a mechanism for catalyst deactivation due to fouling and the loss of activity due to thermal degradation. 1. I N T R O D U C T I O N The cyclic operation of fixed-bed reactors is common practice in many of today's industrial processes. Thus, for instance, the dynamic operation of autothermal reactors where mildly exothermic reactions are carried out has shown to enhance the process performance significantly. This paper examines the operation of an industrial process under cyclic conditions as an alternative to overcome catalyst deactivation effects.

Figure 1: Schematic of the process

The catalytic dehydrogenation of 1-butene into 1-3-butadiene is commonly carried out under adiabatic conditions, at high temperatures, and low reactants partial pressures. The operation is cyclic, with the feed being periodically switched between beds. For continuous effluent flow and concentration, at least three reactors in parallel must be used (cf. Figure 1).

118 The process is started-up with the butene being fed to the first reactor. A fast catalyst deactivation by coke deposition occurs. Thus, after a short operation time, the feed is switched to a second catalytic bed. The first bed is then purged with steam, and the coke deposited on the catalyst is burned off by combustion in air or mixtures with low oxygen concentration. When the regeneration has been completed, the air stream is shut off and a fuel gas mixture is admitted to burn off the residual oxygen and pre-treat the catalyst under reducing conditions. After a purge stage, restarting the dehydrogenation stage completes the cycle. Typical industrial operation conditions require 15-20 minutes cycles (Rielly, 1977, Craig and Dufallo, 1979). This paper is aimed at the analysis of the transient behaviour of a threereactor train in reaction-regeneration cycles. The model will include the coke deposition and sintering effects on the process performance.

2. MATHEMATICAL MODEL The mathematical model is based on previous contributions (Borio et al., 1992; Borio and Schbib, 1995). Unsteady-state, plug flow and adiabatic conditions are considered for the four stages. The dehydrogenation stage (DS) is simulated by means of a pseudo-homogeneous model, using the kinetic data reported by Dumez and Froment (1976). The formation of coke occurs via a parallel and series mechanism with the main reaction. Since the intermediate stages last only short periods of time, heat transfer between gas and solid is considered during the purge (PS) and evacuation (ES) stages, with changes in catalyst activity by coke deposition being neglected. During the regeneration stage a sharp interface model is assumed for the catalyst particle, whereas the external mass and heat transfer resistances are considered (Borio and Schbib, 1995). The kinetics parameters for coke combustion over a Cr203/A1203 catalyst are taken from the literature (Mickley, 1965). The loss of catalytic activity under coke-free conditions is accounted for through the kinetic model proposed by Blasco et al. (1992). The dynamic model is solved numerically by an iterative process, the details of the numerical algorithm can be found elsewhere (Borio and Schbib, 1995).

Table 1: Kinetic models used in the simulation d e h y d r o g e n a t i o n reaction 1-C4H8 --->C4H6 + H2

with rD = TI kD f(pB, pD, pH) ac1 as

coke formation m e c h a n i s m

and,

1-Calls -~ coke

rc = [kcl fl(ps, pD, pH) + kc2 f2(ps, pD, pH)] ac2 as

C4H6 --->coke

following Dumez and Froment (1976)

119

sintering kinetics (coke-free conditions)

regeneration stage (coke burning)

dCc _ kcCca s dt

d a s _ ksa2s 2 dt following Blasco et al. (1992)

following Mickley et al. (1965) 3. R E S U L T S AND DISCUSSION

Figures 2-5 show the temperature, coke and activity axial profiles during the first 1000 cycles (i.e., 250 hours) of operation. The operating conditions for the four stages remain unchanged as detailed in Table 2.

Table 2: Operating Conditions of the different stages Stage time (min.) To [K] P [atm] Feed

Dehydrogenation

Purge

Regeneration

Evacuation

5 823 0.25 1-butene

2 823 1 steam

6 823 1 2.8% 02 97.2% N2

2 823 0.25 fuel-gas

For comparison purposes, the axial profiles corresponding to the "steady-state cycle" (SSC) are also shown. These results were obtained by neglecting catalyst sintering in the model (i.e., the activity coefficient remains constant at as = 1). Once this condition is reached, each cycle becomes equivalent to the previous one, provided that the operating conditions of the different stages remains unchanged (Borio and Schbib, 1995). 920

0.12 . cycle 40 SSC "~ 0.08

Z

-d 840

800 0.00

'-' 0.04

, 0.20

0'40 z[m]

0.'60

i 0.80

F i g u r e 2: Temperature profiles after the regeneration stage.

0 0.00

cvcle~ 0.20

0.40 z[m]

0.60

0.80

F i g u r e 3: Coke profiles after the regeneration stage.

120 The initial condition for cycle 1 (start-up) is an isothermal catalyst bed (at Ts = 823 K), packed with a fresh catalyst load (i.e., coke-free conditions and no sintering). The fast initial reaction promoted by the fresh catalyst results in an accumulation of heat in the solid phase during the early cycles (cf. Figure 2). Due to the significant catalyst sintering occurring during this early stage, the average axial solid temperature shows a continuous decrease from cycle 40 onwards. A non-monotonic temperature is observed with the reactor exhibiting a hot spot. During the subsequent cycles, catalyst sintering results in a quenching of the first reactor section, and a shift of the hot spot towards the end of the reactor. Catalyst sintering is also responsible for lowering the coke deposition (during the DS) and regeneration rates (Borio et al., 1996a, b). Thus, the concentration of residual coke shows an initial increase (up to around cycle 80) and then decreases towards negligible levels (lower than 1%, cf. Figure 3). 1.2

1 cycle

1

SSC

0.8 cycle 4(I__

0.8

~ ) 0.00

5OO

0.'20

0.'40

z[m]

0.60

0.6 0.4

200

0.4

~

SC

c~cle 1

t

le 40 O0

0.2

0.80

F i g u r e 4: (coke-free) Catalyst activity after the regeneration stage.

0 0.00

i

i

1

0.20

0.40 z[m]

0.60

0.80

F i g u r e 5: (combined) Catalyst activity after the regeneration stage.

The loss of catalyst activity by sintering is shown in Figure 4. After a short initial period, the rate of sintering shows a continuous decreasing trend with increasing operating time. After 250 hours of operation, the average catalyst activity has been reduced to approximately 20% of that of the initial fresh catalyst load. The non-uniform sintering observed along the catalyst bed is a direct result of the non-linear temperature profiles with the second half of the reactor showing the highest catalyst deactivation levels. Figure 5 presents the catalyst activity when both coke deposition and sintering are taken into account. This combined activity coefficient (a) includes the activity coefficient for coke-free conditions (as) and the loss of activity by coking (acl, see Table 1), i.e. "a" = as acl. At the beginning of the process, when only negligible sintering has occurred, the activity profile shows a trend opposite to that of the coke profile (cf., "cycle 40" in Figures 3 and 5). Conversely, once significant catalyst sintering occurs, the residual coke decreases to negligible levels and the shape of the combined activity profile mimics that of the coke-free activity (cf. "cycle 1000" in Figures 4 and 5). As a result, a gradual increase in the catalyst activity in the last zone of the reactor is observed (cf. cycles 500 and 1000 in Figure 5).

121 Figures 6-13 show the dynamics of the temperature and coke profiles in the two main stages of the process (DS and RS). Figures 6-9 correspond to cycle 40, when the catalyst has not been strongly sintered. Figures 10-13 correspond to cycle 1000, i.e. after 250 hours of operation. 920

0.09

0.06

880

e~ =5 rain

840

800

L)

0.00

0.03

~

0.00 O.iO

0'40

z[m]

0.60

0.80

F i g u r e 6: Evolution of the temperature during the DS (cycle 40).

=0

020

0.00

0~0

z[m]

060

080

F i g u r e 7: Evolution of the coke profiles during the DS (cycle 40).

Figure 6 shows that the temperature decrease, and therefore the extent of the desired reaction, occurs only in the first part of the reactor. Conversely, in the second zone the rate of dehydrogenation is almost null. This is due to the very low catalyst activity resulting from the accumulation of residual coke near the reactor outlet (cf. "cycle 40" in Figure 5). Since the rate of coke formation is also affected by this residual coke accumulation, a similar trend is observed for the coke deposition (cf. Figure 7). 920

0.09

880 -

0.06 e~

840

~

0.03

//t=O

800 0.00

0.'20

0"40 ztm]

0.60

0.80

F i g u r e 8- Evolution of the temperature during the RS (cycle 40).

~

0.00 0.00

'

=6 rain

0.'20

0.'40 z [m]

0.'60

0.80

F i g u r e 9" Evolution of the coke profiles during the RS (cycle 40).

The temperature increase during the regeneration stage seems significant in the first 50 cm. of the catalyst bed only (cf. Figure 8). The temperature in the outlet zone, on the other hand, remains almost constant due to oxygen depletion (cf. Borio and Schbib, 1995). The low regeneration rates for z > 60 cm. are apparent in Figure 9.

122 0.012 880 -

~, 0.008 ,---, 8 6 0 ,__., [--

t=5 min

0.004

840 -

820 0.00

, 0.20

, 0.40

z[m]

, 0.60

0.000 0.80

F i g u r e 10: Evolution of the temperature during the DS (cycle 1000).

0.00

0.'20

0"40

z[m]

0.60

0.80

F i g u r e 11: Evolution of the coke profiles during the DS (cycle 1000).

When the amount of coke deposited on the catalyst is negligible, for instance at cycle 1000, the oxygen reaches the last reactor section and therefore the regeneration takes place at all the axial positions (cf. Figures 12 and 13). The axial profiles corresponding to the DS of cycle 1000 (Figures 10 and 11) are much smoother than those shown in Figs. 6 and 7. Since the coke levels are considerable lower than those corresponding to cycle 40, the dehydrogenation and coke deposition reactions occur at all the axial positions, being the catalyst activity slightly higher in the first reactor section because of the milder thermal degradation (cf. cycle 1000 in Figure 5). 900

0.012

880 0.008 ,-.-,

:~

860

[-..

0.004 840 ~ ~ . , , / ~ 820

t=O i

0.00

0.'20

0.40 z[m]

0.000

i

0.60

0.80

F i g u r e 12: Evolution of the temperature during the RS (cycle 1000).

0.00

040

z[m]

060

0s0

F i g u r e 13: Evolution of the coke profiles during the RS (cycle 1000).

The influence of the inlet temperature at the different stages on the cumulative production (per bed) for 2400 cycles (600 hr. in operation) is shown in Figure 14. Curves for To = 823 K correspond to the same operating conditions given in Table 2. When the catalyst sintering is not considered in the model (steady-state cycle,

123

SSC) a straight line is obtained, because the production of butadiene per cycle becomes constant after a few cycles. When catalyst sintering is taken into account, the curve for To= 823 K shows a decreasing slope (i.e., a decreasing production of butadiene per cycle). The associated gradual cooling of the catalyst bed in time (cf. Figure 2) contributes to this phenomenon. The curves for To = 723, 773 and 873 K in Figure 14 correspond to different values of the inlet temperatures (the same for the four stages, kept constant on time). The remainder operating conditions are given in Table 2. The maximum amount of butadiene after 600 hr. is obtained for the intermediate inlet temperature, i.e. To = 773 K. For higher temperatures (e.g., To = 873 K), the resulting higher rate of catalyst sintering tends to offset the higher initial production per cycle, yielding a lower cumulative production after 600 hours of operation. Figure 15 shows the effect of the oxygen content at the reactor inlet (during RS) on the cumulative production (per bed), for the operating conditions given in Table 2. The higher the value of yA0 the higher is the production rate, particularly in the first 300 hr. in operation. The cumulative production after 600 hr. increases 6.6% when the oxygen content is increased from 2 % to 4.5%. The most important factor for these changes is the strong dependence of the residual coke concentration on the selected yA0 value that, in turn, affects the catalyst activity and consequently the production rate of butadiene. 140

140 To=823 K"

120 ~'

To=7 7 3 ~ ' ~

~

60

o

40

~

yao=~

,.--,

100

"~ 100

80 ...,

YAo=2.8%,'

120

~ .o

80

o

40

~

0

i

|

200

400

600

Time [hr]

F i g u r e 14: Cumulative production for different values of To

0

200

400

600

Time [hr]

F i g u r e 15: Cumulative production for different values of yAO.

4. C O N C L U S I O N S

An industrial process operating under cyclic conditions has been analysed. The 1-butene catalytic dehydrogenation process carried out in a three-reactor train was chosen to illustrate the effect of catalyst deactivation on the system performance. Catalyst deactivation due to both coke deposition (fouling) and thermal degradation (sintering) was included. The combined effects of sintering and coking reduce the catalyst activity significantly. The decrease in catalyst activity results in a decreased production

124 rate. Although the steepest production loss is observed during the early cycles, thermal degradation remains significant throughout the entire process. As a result, the process never reaches a steady-state condition and a continuous decrease in the butadiene production is observed. In contrast, when sintering effects are ignored, the process reaches a steady regime shortly after the reactor train start-up. Selecting different operating conditions, however, could compensate the sintering effect. Indeed, the analysis presented was performed for constant condition in all stages. An increase, for instance, in the oxygen levels as the system tends to quench itself could compensate deactivation effects and avoid the continuous decrease in the production rate. Alternatively, a gradual increasing of the inlet temperature would have a similar compensating effect. 5. R E F E R E N C E S

Blasco V., C. Royo, A. Monz6n and J. Santamar/a, AIChE J., 38, 237-243 (1992). Borio D. O., M. Men~ndez and J. Santamar/a, Ind. Eng. Chem. Res., 31, 26992707 (1992). Borio D. O. and N. S. Schbib, Comp. & Chem. Eng. 19S, $345-$350 (1995). Borio D. O., N. S. Schbib and A. A. Savoretti. XV Simp. Iberoam. de Catdlisis, C6rdoba, Argentina, ACTAS, Vol. 1, 133-138 (1996a). Borio D.O., N.S. Schbib, A. A. Savoretti. AIChE Annual Meeting, Chicago, USA. November (1996b). Craig R. G. and J. M. Dufallo, Chem. Eng. Prog., Feb., 62-65 (1979). Dumez F. J. and G. F. Froment, Ind. Eng. Chem. Proc. Des. Dev., 15, 291-301 (1976). Mickley H. S., J. W. Nestor and L. A. Gould, The Can. J. Chem. Eng., 61-68 (1965). Rielly T., In Encyclopaedia of Chem. Process. and Des. (McKetta Ed.), V.5, 110170, Marcel Dekker (1977). 6. N O M E N C L A T U R E acl, ac2: local activity coefficients by fouling (coking). as : local activity coefficient at coke-free conditions (sintering) Cc : coke concentration, kg. coke/kg, cat. k 9reaction rate constant p "partial pressure, atm. rc~, rc2 : reaction rate of coke formation, kg.c/(kg.cat h) rD : reaction rate of dehydrogenation, kmol]kg.cat h) t " time, hr. T : temperature, K z" axial coordinate, m

Subscripts

A = oxygen B = butene C = Coke D = butadicne H = hydrogen

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

125

Propyne Hydrogenation Over A Silica-Supported Platinum Catalyst Studied Under Transient Conditions D . R . K e n n e d ) a , B. C u l l e n ~ , D. L e n n o n ~ ~, G. W e b b ~ , P . R . D e n n i s o n b and S.D. Jackson ~ " Department of Chemistry, University of Glasgow, Glasgow, G 12 8QQ, Scotland. b Department of Chemistry, University of Strathclyde, Glasgow G 1 1XL, Scotland. ICI Katalco, PO Box l, Billingham, Cleveland TS23 1LB, England.

ABSTRACT Catalytic activity and selectivity is shown to be critically dependent on the hydrogen concentration. The surface processes responsible for propane formation are shown to be decoupled from those responsible for the selective formation of the alkene, with the observed chemistry adequately described by a 2-site model. 1. INTRODUCTION The hydrogenation of acetylenes is a critical reaction in the petrochemical industry [1]. Consequently there is a large body of work on the catalytic hydrogenation of alkynes, however the vast majority of this work to date has involved principally the hydrogenation of ethyne [2]. The literature on higher molecular weight alkynes, such as propyne, is surprisingly sparse [3-8]. A recent report has shown that rather than ethyne providing a good model for the hydrogenation characteristics of alkynes in general, it appears that the ethyne system is in fact a special case [9]. Clearly then, it is important to examine higher alkynes and to investigate how such systems react. Propyne hydrogenation over a silica-supported platinum catalyst was selected as it represents an unsymmetric alkyne and should provide information on the effect a substituent group (methyl) has on the reactivity, compared to ethyne. Also the selective hydrogenation of propyne is an industrial process, for which an increased database is desirable. An initial study of propyne hydrogenation in a flow regime over a platintun/silica catalyst was performed by Jackson and Kelly [ 10], which principally concentrated on the determination of orders of reaction. In addition, their study highlighted the effect carbonaceous residues have on the rate of reaction. This study investigated further how the nature of these overlayers affect the selective formation of the alkene by examining propyne hydrogenation over a range of reaction conditions. Under appropriate conditions the system undergoes an induction period before achieving steady state operation and, in order to probe this transition, preliminary ex-situ nuclear magnetic resonance experiments have been carried out. ~ICI Lecturer in Catalysis, correspondingauthor

126 2.

EXPERIMENTAL

Apparatus and Reaction Testing A pulse-flow microcatalytic reactor system was used throughout this study. The sample of catalyst was supported on a glass sinter in the centre of the reactor (8 mm i.d. down flow), which was placed inside a furnace. Catalyst temperatures were measured by means of a chromel-alumel thermocouple placed alongside the catalyst bed. Approximately 0.20 g of catalyst was reduced in a 6% H2 in N2 flow (30 ml min ~) while the temperature was raised from 293 K to 573 K at 5 K min -~, then held at 573 K for 2 hours. The sample was then flushed with pure hydrogen for 30 minutes, flushed with helium for a further 30 minutes and then allowed to cool to ambient temperature in a flow of helium. Pulses of reactant gas (24.3 lamoles of propyne/hydrogen at the ratios specified below) were injected into the helium cartier gas ~stream immediately above the catalyst. One pulse of gas corresponds approximately to one molecule for each surface Pt atom, as determined by CO chemisorption. On elution from the catalyst bed the full pulse was analysed by on-line gas liquid chromatography, using a thermal conductivity detector with a Poropak QS column. The amount of gas adsorbed/reacted, from any pulse, was determined from the difference between calibration peak areas and the peak areas obtained following injections of pulses of comparable size onto the catalyst. Fresh catalyst was used for each experimental run, including the nmr experiments. All reactions were performed at 293 K. The helium, hydrogen and 6% H2/N 2 (BOC, 99% purity) were purified using in-line deoxygenating and drying traps. The propyne (BDH, 96% purity) was purified through bulb to bulb distillations prior to use. Propane and propene (BOC 99% purity) were used for gas chromatography calibration.

Catalyst Preparation The 1% w/w catalyst was prepared from hexachloroplatinic acid (Johnson Matthey) by impregnation. The silica support material was M5 Cab-o-Sil (surface area 200 m 2 g-l). The mixed slurry was prepared in water then transferred to a rotary evaporator where the water was slowly removed by maintaining the sample at 353 K under a nitrogen atmosphere to produce a free flowing powder. The resulting catalyst was then dried for 12 hours at 373 K. Active surface area measurements were performed on a pulse-flow gas chemisorption apparatus. Carbon monoxide chemisorption revealed 16.29 x 10 lg surface Pt atoms g-l, assuming a Pt:CO stoiciometry of 1:1, which corresponds to a metal dispersion of 48%. The total surface area as determined by BET analysis was 187 m 2 g-Z. Temperature programmed reduction showed an onset of reduction at 468 K and complete conversion of the metal precursor to the metallic state at 518 K. Atomic absorption spectroscopy showed the actual metal loading to be 1.1%. NMR Proton nuclear magnetic resonance experiments were performed on a Bruker MS L 100 MHz spectrometer using a Doty high volume probe. For these experiments the glass reactor was modified so it filled the nmr probe and could be sealed after reaction. Approximately 0.40g of catalyst were used and the amount of gas passed over the catalyst scaled accordingly so that 1 gas pulse for the nmr sample gave the same number of reactant molecules per gramme of catalyst, as in the conventional reaction testing experiments described above. After pulsing the reactant mixture over the catalyst, the sample was isolated, evacuated to a pressure of ca. 1 x 10.5 mbar then the reactor sealed. The sample was then transferred to the spectrometer for ex-situ measurements. Spectra were recorded using a single 90 ~ pulse of length 7.625 ~ts.

127 Recycle delays were chosen such that further increases in this time produced no further increase in intensity. Delays of 200 s were used. FIDs were recorded using 2048 data points with a dwell time of 8 las. 256 scans were acquired for determination of a spectrum.

3.

RESULTS

Alkyne hydrogenation produces the corresponding alkene which itself can be hydrogenated further to the alkane. In order to evaluate the role of hydrogen concentration in the selective formation of the alkene, two reaction mixtures were used: C 3 H 4 : H 2 in a 1 : 1 ratio and C3H4 : H 2 in a 1 : 3 ratio. For the former, reaction was studied over a series of 10 pulses but this was extended to 30 pulses for the excess hydrogen mixture. Equimolar reaction mixture,

CsH4"H

2 =

1 : 1.

Figure l(a) presents the extent of conversion of propyne as a function of pulse number. Low conversions are observed, with activity substantially reduced after the first pulse. Catalyst deactivation is clearly a dominant process under these conditions.

83 %

.-,..

|

|

(b)

-e-~~ -~- R ' ~

43 23

%

.

Pulse

Fig. 1. Percentage conversion for (a) equimolar mixture and (b) excess hydrogen mixture.

.

.

.

.

PUlse Fig. 2. Variation of product distribution for the equimolar mixture.

The variation in product distribution with pulse number is shown in Figure 2. For the first pulse propane is the only product, with approximately 75% of the initial propyne content of the pulse unreacted. From pulse 2 onwards, the yield of propane decreases and propene is formed at a constant (low) amount over 10 pulses. Carbon mass balance data for this reaction sequence are presented in Figure 3. Retention of carbon by the catalyst is large for the first pulse then decreases sharply in pulse 2. It continues to decrease up to pulse 6 where no carbon retention is observed and thereafter small amounts of carbon are removed from the surface on continued pulsing.

128 The mass balance data appears to correlate well with the catalyst activity and shows that in the initial pulse there is dissociative adsorption of the propyne which goes on to form hydrocarbonaceous overlayers which, in this case, appear to block metal sites and hence reduce activity. Given that propene is only formed in the presence of this overlayer suggests it is responsible for the selective hydrogenation to form the alkene. The amount of carbon deposited on the catalyst in the first pulse is 2.00 x 10 ~9 carbon atoms g~ which corresponds to a surface Pt : C ratio of 1 : 1.23. The cumulative value at pulse 10, which is representative of steady state operation, is 1.97 x 10 ~9carbon atoms g " .

E x c e s s h y d r o g e n mixture, C3H4 : 1-12 = 1 : 3.

The variation of conversion with pulse number for the propyne hydrogenation when the excess hydrogen mixture is used is presented in Figure l(b). In comparison to Fig l(a) there is a significant improvement in catalytic activity, with initial pulses exhibiting conversions of ca. 75%, which gradually decrease to a steady state value of ca. 53%. The product distribution is more complex than that seen with the equimolar mixture (Fig 2) and is presented in Figure 5. Here, marked changes in selectivity are observed. Over the first 16 pulses the only observed elutants are propane and unreacted propyne. Then suddenly, at pulse 17, propene is detected and continues to be produced at a fixed amount. This constitutes a distinct step change in propene selectivity (defined as [alkene/(alkene + alkane)]), from 0% at pulse 16 to an approximately constant 54% from pulse 17 onwards. Interestingly, there is little change over this transition period in either the propane formation or the amount of unreacted propyne formation. The latter is reflected in Fig 5, which shows a relatively fiat profile arotmd this region.

Fig. 3: Carbon mass balance for the equimolar mixture.

Fig. 4: Carbon mass balance for the excess hydrogen mixture.

Carbon mass balance data for this reaction sequence are presented in Figure 4. Again, the profile is considerably different to that seen for the equimolar mixture (Fig 3). Initially high amounts of carbon are retained by the catalyst up to pulse 16, then from pulse 17 onward

129 effectively no carbon is retained at the catalyst surface. In a similar manner to that seen with the lower hydrogen concentration, the carbon mass balance data correlate well with the product distribution (Figure 5). Specifically, for the period when hydrocarbonaceous overlayers are being formed at the catalyst surface only propane is produced, but once formation of this overlayer is complete then propene production commences in quite a dramatic step-like fashion. The initial period corresponds to a non-steady state regime and post pulse 17 to a steady state regime, with approximately equal amounts of propane and propene detected in the product distribution. The cumulative amount of carbon deposited on the catalyst up to pulse 17 is 2.55 x 102o carbon atoms g~ which corresponds to a surface Pt : C ratio of 1 : 15.7. It seems reasonable to assume that complete formation of this overlayer is responsible for the selective formation of the alkene. Furthermore, the fact that propane formation, and also the extent of conversion, remain effectively constant over the period where propene production switches in is interpreted as indicating that the surface processes responsible for propane formation are decoupled from those responsible for the selective formation of the alkene. The extent of carbon laydown initially appears large and implies that, under the conditions examined here, alkene formation requires a hydrocarbonaceous overlayer rather than bare metal. Such a deposition process is well documented for the hydrogenation of unsaturated hydrocarbons [2,5,10]. The fact that the amount substantially exceeds the number of surface Pt atoms suggests a degree of spillover on to the support.

0

2

".

~i / \

~.mmt i n i . i i n i i l l

,,.,~ \,~ oo,',.o.,,,,,.ooo,,%,,,

fE .

~-

~ol~r~

~-

~Ol:~r~

- a - ~opyr~

Pulse

3o

Fig 5: Variation of product distribution with pulse number for the excess hydrogen mixture.

I

e

40,000

I

i

I

ul~

~

I

,

I

-~,000

Fig. 6. Background substracted ~H NMR specman of Pt/SiO2 catalyst after (a) 13 and (b) 25 pulses of excess hydrogen mixture.

NMR With reference to Figure 5, the onset of propene production is quite dramatic. How has the nature of the catalyst surface changed between pulses 16 and 17 to initiate this process? One probe that is sensitive to changes in the catalyst surface is solid state nmr spectroscopy. Provided experiments are performed with due regard to the complex relaxation properties of

130 the chemical system, quantitative measurements are possibi'e and resonance frequency positions can provide information on the electronic nature of the substrate [ 11,12]. The IH nmr spectrum of the catalyst recorded after the cleaning/reduction cycle shows an intense single feature at 0 ppm which is due to hydroxyl groups present on the support material [ 12]. Figure 6 shows background subtracted spectra of the catalyst samples recorded after 13 (Figure 6(a)) and 25 pulses (Figure 6(b)) of excess hydrogen mixture have been passed over the catalyst. For these spectra the spectrum of the clean catalyst has been subtracted from the dosed spectra in order to obtain the spectnun of hydrogen retained at the catalyst surface. The spectra are recorded under quantitative conditions. There is an increase in intensity on going from 13 to 25 pulses, with the latter having an integrated area 151% that of the former. This increase corresponds to the additional hydrocarbonaceous laydown between the two doses. The carbon mass balance data reveal a relative increase over this period of 26% (Figure 4). Proportionally, the earlier pulses represent a higher C : H ratio: by pulse 13 the catalyst has achieved respective C and H retentions of 79% and 66% that of their steady state capacities, as represented by pulse 25 data. The selectivity data (Figure 5) suggest that it is the formation of the later, relatively hydrogen-rich, deposits that are responsible for the formation of the alkene. The resonance maxima is the same for both spectra, at +1 ppm, though the 25 pulse spectrum exhibits a noticeable tail to low frequency, centered about -20 ppm. The position of the support hydroxyl peak is used to define a position of zero ppm [ 12]. Fig 6 represents the 1H nmr f'mgerprint of the species retained by the catalyst before and after the onset of propene production. The increase in intensity occurs only on the low frequency side of the main peak, which suggests that these deposits are interacting with the metal. Chemisorbed hydrogen produces a nmr signal that is attributed to a Knight shift arising from an interaction of the metal-chemisorbed hydrogen nuclei with polarised metal conduction electrons [12]. Consequently, the frequency of the chemisorbed species reflects the local electronic environment and no significant shift would be expected for residues associated only with the support. This nmr data appears to define at least two stages to the catalyst conditioning process. Initially, hydrocarbonaceous deposits resulting from the dissociative adsorption of the alkyne accumulate at the catalyst surface. Given the total quantities of carbon retained with respect to the number of metal surface sites, some of this material must reside on the support. This state saturates and then relatively hydrogen-rich deposits accrue, which are associated with the metal. Propene production clearly does not scale with the formation of this overlayer but rather switches in on the completion of the layer. Why such a step change, which is effectively reproduced in the equimolar data as well, is observed is uncertain at this stage. The fact that a distinct resonance is observed for this favourable deposit indicates that, on the nmr timescale, this state is not exchanging with the initial residues. Identification of this crucial species highlights the potential of nmr spectroscopy in catalyst characterisation studies.

4.

DISCUSSION

Catalytic activity is substantially greater for the excess hydrogen mixture, confirming the importance of the hydrogen concentration in controlling the reaction chemistry. The fact that this enhanced activity occurs despite a carbon retention 13 times that seen with the equimolar hydrogen mixture clearly demonstrates that the overlayer is not simply acting as a site blocker, but also has promotional benefits as well. These results show the hydrogen supply

131 affects the nature of the overlayer and this is reflected in changes in the selectivity of the substrate over the course of the reaction sequence. The overlayer must form from the originally adsorbed propyne but in the absence of sufficient hydrogen this overlayer will decompose to produce carbon rich residues which ultimately will block adsorption sites, and hence reduce activity. However, in the presence of an increased hydrogen supply at the catalyst surface the initial residue can be hydrogenated to form products, or also to form hydrogen rich overlayers which promote propene formation. This promotion occurs with the carbonaceous overlayer acting as a hydrogen transfer agent [13]. For both of the mixtures studied, propene formation only occurred in the presence of this overlayer, confirming the importance of this layer in controlling selectivity. Collectively these findings support the earlier findings of Jackson and Kelly [ 10] where it was found that the C3H4 : H2 ratio affected the extent to which the carbonaceous residue acted as a promoter or a poison. Ossipoff and Cant have observed a range of oligomerisation products produced from propyne hydrogenation over a Cu/SiO~ catalyst [7] but, given that only C3 fragments are observed in the product distribution in these reactions, suggests no such processes are occuring with the Pt/SiO2 catalyst studied here. Concentrating on the excess hydrogen selectivity profile, it is useful to consider a model of three different types of sites which has been used extensively to explain selectivity phenomena in alkyne and alkadiene hydrogenations over a range of metals [ 14]. Type I is responsible for the direct hydrogenation of alkyne to alkane, Type II is responsible for hydrogenation of alkyne to alkene, Type III is responsible for the hydrogenation of alkene to alkane. In our case only Types I and II appear to be active. Under all circumstances propane production is observed. Any contribution from Type III sites is unlikely as when propene formation steps in, no increase in propane production is seen (Figure 5). Type I sites are present throughout the reaction sequence. This site is deminished due to deactivation processes but achieves steady state operation by pulse 10. The Type II site appears to form as part of the catalyst conditioning process. Figures 4 and 5 suggest that this site is in fact the hydrocarbonaceous overlayer formed as a result of a favourable hydrogen supply. The 1H nmr fingerprint of this species is shown in Figure 6(b). With propene and propane formed in roughly equal amounts under steady state conditions, it would appear that Type I and II sites co-exist. The details of this multi-site approach are different to that observed by Jackson and Kelly. These authors suggest that deactivation processes convert Type I sites into Type II sites, which thereby increases the number of Type II sites [10]. Such a progressive scheme is not active in this case and, moreover, no reduction in propane formation is seen once propene production is established. One possible explanation for this discrepancy could be the difference in the metal particle size of the two Pt/SiO2 catalysts. The catalyst in Jackson and Kellys' study had a dispersion of only 3% whereas the catalyst used in this study has a dispersion of 48%. The differences in catalyst morphology suggest that the selectivity characteristics of this system are structure-sensitive. Differences in carbon retention are also apparent between the two substrates: Jackson and Kelly report a surface Pt : C ratio of 1 : 8.0 which was insensitive to propyne: hydrogen ratio. In contrast the better dispersed catalyst in this study produced ratios of 1 : 2.1 and 1:16 for steady state operation for the equimolar and excess hydrogen mixtures respectively. Collectively these results show the importance of hydrogen supply in controlling activity and selectivity. Such a concept has been discussed previously [ 15,16] but this work clearly demonstrates the principal. This study has only examined a small range of alkyne/hydrogen ratios yet observed marked changes in catalytic behaviour. It is intriguing to consider if gains in selective hydrogenation to form alkenes could be achieved by further

132 manipulation of the reaction conditions. currently underway in our laboratories.

Work on this industrially significant matter is

ACKNOWLEDGEMENTS The EPSRC (DRK) and ICI Katalco (BC) are thanked for the award of studentships. DL thanks ICI for the award of an ICI Lectureship. Professor C. Snape (University of Strathclyde) is thanked for providing access to nmr facilities.

REFERENCES

.

3. 4.

5. 6. .

8. 9. 10 11. 12. 13. 14. 15. 16.

M. Derrien in Catalytic Hydrogenation, Ed. L. Cerveny, Elsevier, Amsterdam, 1986, p. 613 G. Webb, Catalysis Today, 7 (1990) 139 G.C. Bond and J. Sheridan, J.C.S. Faraday Trans., 48 (1952) 651. R.S. Mann and K.C. Khulbe, J. Phys. Chem., 73 (1969) 2104. S.D. Jackson and N.J. Casey, J.C.S. Faraday Trans., 91 (1995) 3269. J.T. Wehrli, D. J. Thomas, M.S. Wainwright, D.L. Trimm and N.W. Cant, Applied Catalysis, 66 (1990) 199 N.J. Ossipoffand N.W. Cant, J.Catal., 148 (1994) 125. H.N. Choksi, J.A. Bertrand and M.G. White, J. Catal., 164 (1996) 484. S.D. Jackson and G. Kelly, Current Topics in Catalysis., 1 (1997) 47. S.D. Jackson and G. Kelly, J. Mol. Catal., 87 (1994) 275. A.T. Bell and A. Pines, NMR Techniques in Catalysis, Marcel Dekker, New York, 1994. M.A. Chesters, K.J. Packer, H.E. Viner, A.P. Wright and D. Lennon, J.C.S. Faraday Trans., 92 (1996) 4709 G. Webb and S.J. Thomson, J.C.S. Chem. Comm., (1976) 526. N.C. Kuhnen, S.J. Thomson and G. Webb, J.C.S. Faraday Trans. 1, 79 (1983) 2195. A.S. A1-Ammar and G. Webb, J.C.S. Faraday Trans., 1, 74 (1978) 195. S.D. Jackson and L.A. Shaw, Applied Catalysis A, 134 (1996) 91.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

133

A TAP reactor investigation of the oxidative dehydrogenation of propane over a V/MgO catalyst: experiment and modeling Y. Schuurman, T. Drcamp, J.C. Jalibert, C. Mirodatos Institut de Recherches sur la Catalyse, C.N.R.S. 2, avenue Albert Einstein, F-69626 Villeurbanne Crdex, France ABSTRACT The oxidative dehydrogenation of propane over a V/MgO catalyst was studied in a TAP-2 reactor. The experiments were carried out using oxygen rich and lean feeds at temperatures between 500 and 650~ The data can be described adequately by a rate equation based on a reaction sequence involving elementary steps. The rate-determining step is the irreversible adsorption of propane leading to an alkyl intermediate that reacts rapidly to adsorbed propene. Sorption of propene is reversible and further oxidation of adsorbed propene leads to COx. No information on the oxygen activation could be obtained. All kinetic parameters were found to be physically meaningful and sensitive to the oxidation state of the catalyst. INTRODUCTION Vanadium catalysts supported on magnesia a r e known to be among the most active and sdective materials for the oxidative dehydrogenation of propane (ODHP), leading to propene yields limited to around 20% [ 1,2]. Improving the yield requires an advanced knowledge of the mechanism, and especially of the total oxidation pathways. Most of the published kinetic studies considered exclusively a consecutive reaction scheme, where CO x are only formed from the further oxidation of propene [ 1,3,4]. In fact, both selective and non-selective products are always observed even at low conversion levels [2]. Therefore, both a parallel and a sequential scheme for the partial/deep oxidation of propane should be considered. In addition, no consensus on the type of oxygen species involved is reached in the literature. Several types of active oxygen are believed to activate the propane molecule : lattice and/or monatomic adsorbed oxygen [1-4]. Mostly power law models are used to describe the kinetic data. The propane conversion is found to depend only on the propane pressure with a kinetic order between 0.5 and 1. For the formation of propene and COx, however, an order for the oxygen dependence of 0.2-0.5 was found [1,3]. A few, more detailed, rate equations have been proposed, but it was not possible to discriminate between rival models, such as Langmuir-Hinshelwood and MarsVan Krevelen mechanisms [6]. Non-steady state methods have more potential in discriminating between different mechanisms and therefore the ODHP reaction over an optimized V/MgO catalyst was investigated by means of steady-state isotopic transient kinetics combined with msitu FT-IR [2]. A Mars-Van Krevelen type of reaction mechanism was proposed, where the propane activation and selective formation of propene takes place over surface vanadium 5+ ions and lattice oxygen anions. A preliminary study in a TAP reactor (Temporal Analysis of Products) confirmed qualitatively the key features of the ODHP mechanism [7]: both propene and carbon oxides are primary reaction products but CO x is also produced by a secondary oxidation of propene. Propane selective and deep oxidation occur at the same surface site but would involve nucleophilic lattice oxygen and adsorbed electrophilic oxygen, respectively. The oxidation of propene would involve both types of oxygen species.

134 In this work a simplified microkinetic model is developed in order to derive from the TAP responses the kinetic parameters involved in the main steps of the reaction. In order to test the sensitivity of the ODHP to the propane-to-oxygen ratio, two series of TAP experiments obtained under oxygen rich and oxygen lean conditions are compared. EXPERIMENTAL Catalyst. A V/MgO catalyst containing 14 wt.% V was found to give the highest propene yield as a function of the V content [5]. The catalyst preparation is described in detail in [2,5]. From HREM, XPS, UV-vis, XRD and m-situ electrical conductivity, the active surface was shown to be essentially a monolayer of amorphous VO43" units scattered over the magnesia as isolated and polymeric species [8]. The BET surface area amounts to 43 m 2 g-1. TAP experiments. Transient experiments under vacuum conditions were carried out in the TAP-2 (Temporal Analysis of Products) reactor. Narrow gas pulses of reactants are introduced in a microreactor (25.4 mm in length and 4 mm in diameter) which is evacuated continuously. The response of these pulses as a function of time is detected by a quadrupole mass spectrometer (QMS) located directly underneath the reactor exit. The shape of the response reflects the diffusion, adsorption, desorption and reaction, as described in [9]. To avoid temperature gradients over the catalyst bed, 75 mg of the catalyst (dp = 200300 lam) was placed in the microreactor between two layers of quartz (dp = 200-300 ~tm), with a thermocouple inside the catalyst bed. The catalyst was pretreated at 550~ by a flow of propane, oxygen and helium for 1 h at atmospheric pressure. Two different propane-to-oxygen ratios were used, C3H8/O2= 2 and 0.5. A smooth transition to the vacuum experiments was assured by first instantly switching by means of a four-way valve the reactive flow to pure nitrogen for 1 h, before exposing the reactor to the vacuum chamber. The same oxygen-topropane ratio during single pulse experiments was used as during the pretreatment. Additionally propene, CO and CO2 were pulsed over the catalyst. Pulse experiments were carried out at 500, 550, 600 and 650~ In order to record the fragmentation patterns and to calibrate the mass spectrometer, all reactants and products mixtures were pulsed over a quartz bed. In all cases argon was mixed in the pulse valve as an internal standard. A correction of the response signals at a specific m/e ratio was performed in the case of overlap of different substances. The following m/e were used : propane 29, propene 41, carbon monoxide 28, carbon dioxide 44, water 18 and argon 40. The water response was too broad to be accurately integrated. For each response 10 pulses were averaged to improve the signal to noise ratio. QMS calibration factors relative to argon were calculated by integrating the pulse responses over the quartz bed. The conversion and selectivities (e.g. for propene) were calculated according to the following equations:

(1)

=

c'H')IAc'H I rat AAr)

(2)

where Ai (Vs) is the integrated surface response area, Ni (mol) the molar concentration in the pulse valve and 7t (mol Vls q) the QMS calibration factor of the i component. To simplify the mathematical modeling of the transient responses, pulse intensities were kept sufficiently small to ensure Knudsen flow. This was verified by the independence of the response shape on the pulse intensity [9].

135 RESULTS

25t

C3He

50 *C

20

3= o 15

~n ":o~ 20

C3H

, I ~ R ~ 6 O0 *C -c

E

~ 10 C0

5

i

0.0

f

C02

i

i

i

i

i

0.1

0.2

0.3

0.1

0.2

tim e, s

0.3

time, s

Figure 1. Experimental (circles)and model (lines) responses at 600~ from a pulse of C3HdO2 = 2 over V/MgO.

Figure 2. Experimental (circles) and model (lines) responses for propene as function of temperature from a pulse of C3Hs/O2 = 2 over V/MgO.

Figures 1 and 2 show the transient responses of propane, propene, CO and CO2 from a pulse of propane and oxygen at a ratio of 2 at 600~ and the propene pulses as a function of temperature respectively. Similar responses were obtained with a propane-to-oxygen ratio of 0.5. The narrow pulse of propane indicates its consumption, while propene, CO and CO2 all have residence times greater than that of argon, typical of products. The residence time for argon amounts to approximately 70 ms. In all experiments with the two different propane-tooxygen ratios the oxygen conversion was almost complete. Table 1 gives the conversions, selectivities and propene yields as a function of temperature. A reasonable closed mass balance was found. This table shows that an oxygen rich feed leads to significantly higher yields in propene, mostly due to higher conversion, the selectivity towards propene remaining rather C3Hs/O2 independent. The propene selectivity decreases with increasing temperature while the CO selectivity remains constant and the CO2 selectivity increases.

C3H8 2O ~ 15 __.N o ,--

10

CO

i

i

0.1

0.2

0.3

time, s

Figure 3. Experimental responses at 600 ~ from a propene pulse over V/MgO.

Figure 3 shows that upon a pulse of propene over a catalyst pretreated with an oxygen lean mixture, CO and CO2 are formed as products, indicating a consecutive oxidation of propene. No formation of CO2 was found upon the introduction of a pulse of CO, except at temperatures higher than 650~ ruling out any consecutive formation of CO2 from CO. No interaction of propane, propene and carbon monoxide with the MgO support was observed, when pulsed over a calcined MgO support free from vanadium. However, carbon dioxide was completely adsorbed and no transient response could be observed.

136 Table 1. Conversions and selectivities. T (~ C3H8]O2 = 2 500 550 600 650

C3H8/O2 = 0.5

XCaH8 SC3H6 Sco

Sc02 YCaH6 XC3H8 Sc3H6

0.15 0.28 0.47 0.65

0.09 0.11 0.29 0.41

0.64 0.43 0.27 0.16

0.29 0.30 0.30 0.44

0.10 0.12 0.13 0.10

0.32 0.49 0.68 0.84

0.53 0.45 0.30 0.16

Sco

Sc02 YCaH6

n.d. 0.18 0.38 0.39

n.d. 0.16 0.40 0.48

0.17 0.22 0.20 0.13

Modeling. The model describes the Knudsen flow of reactants and products through the three reactor zones (quartz, catalyst and quartz) and accounts for all reaction steps according to the mechanism considered in the catalyst zone. Assuming no radial concentration gradients, the one-dimensional continuity equations for component A in the gas phase and adsorbed on the surface are given by: Quartz zone:

8CA _ DA 82CA

e 8t

8z 2

Catalyst zone: e 8CA Ot = DA 82CA cTz2 _(l_e)Ns(kaCA(l_OA)_kaOA )

(3) (4)

and the continuity equation for the adsorbed species (catalyst section only): r

cTt

-----k a C A (1 - 0 A ) - k d O A - E krt~A

(5)

In equations (3) and (4) the term 1-0a can be set equal to 1, since the number of active sites is much higher than the number of molecules A (pulse size < 1014 molecules, while the number of active sites > 101s), so that 0a < 104. The initial and boundary conditions applying for the TAP reactor are detailled in [9] The system of the above partial differential equations with the accompanying initial and boundary conditions was solved by transformation into the Laplace domain with respect to time [ 10]. Parameter estimation was performed by Marquardt's method [ 11 ]. The Knudsen diffusivities for all species were determined by regression of pulse experiments over a quartz bed. Then the Knudsen diffusivities in the catalyst zone were estimated from experiments in the three-zone reactor at 300~ to avoid any reaction. The diffusivities at other temperatures were calculated by the following equation:

DA (T:I=DA (T~)_~T~ /r~

(6)

These values for the diffusivities were then fixed during the estimation of the kinetic parameters.All other kinetic parameters were estimated by a non-linear regression of the transient responses at all temperatures simultaneously. For the regression analysis a reparametrized form of the Arrhenius and Van 't Hoff equations were used. A full statistical analysis was performed after regression, which allowed calculating the 95% confidence intervals on the estimated parameters. All parameters were found to be statistically significant with a 95% confidence interval of approximately 10%.

Reaction mechanism. No associative adsorption of propane (physisorption) on V/MgO was detected by SSITKA, i.e. under steady-state atmospheric conditions [2]. Such weak adsorption being still more unfavored under the low pressure TAP conditions, it can be ruled out. Therefore the direct dissociative adsorption is considered. This step involves a C-H bond breaking and leads to an adsorbed alkyl species and the formation of a hydroxyl group (step 1).

137 A second hydrogen can be abstracted from the adsorbed alkyl species leading to adsorbed propene and a hydroxyl group (step 2). The latter species were clearly revealed by in situ DRIFT experiments [2]. A strong interaction between the double bond of propene and the Lewis acid sites which were found to be maximized for this vanadium content [12] can be postulated. Therefore a reversible sorption of propene is considered (step 3). The formation of CO and CO2 proceeds through a series of steps of hydrogen abstraction, carbon-carbon bond breaking and carbon-oxygen bond formation. Close CO and CO2 responses were obtained in the TAP reactor either from propane or from propene (Figures 1 and 3, respectively) and no consecutive CO --, CO2 reaction was detected. The only stable species accumulated under steady-state conditions was found by in situ DRIFT to be carbonates (mostly bidentate) adsorbed on the magnesia [12]. This side accumulation which also corresponds to the irreversible trapping of CO2 by pure magnesia in the TAP reactor, cannot therefore be considered as belonging to the proper catalytic cycle on the vanadium sites. Accordingly, all the elementary steps leading either to CO or to CO2 will be lumped into one first order step, which is the rate-determining step for the formation of COx. It can be either formed from the alkyl species (step 4) or from the adsorbed propene (step 5). A reversible sorption step is also considered to account for the likely interaction i) between the basic lattice oxygen of the active phase and the acidic CO2 molecules, and/or ii) between the surface vanadium anions and CO (step 6). The reoxidation steps of the reduced sites, including oxygen dissociative adsorption and surface dehydroxylation into water should be included in a complete catalytic cycle. However, in this study no information on these steps was obtained due to the complete oxygen conversion and the weak broad water signal. Several different oxygen species were suggested to be present on the V/MgO surface, each having a specific role in the selective and non-selective oxidation of propane [1,10]. As mentioned above, under the TAP conditions very low surface coverages are achieved and as a result the relative changes in the oxygen surface concentrations are very small. Hence, no discrimination can be made between the reaction of adsorbed hydrocarbons with different oxygen surface species. In the mechanism below all surface sites are indicated with a * without differentiating them. All the above considerations give the following mechanism: C3H8 + 2* C3H7" + *

C3H7" C3H6" C3I%* COx*

~ ~

~

~

~ P' ~ ~

C3H7" + OH* C3I-I6"+ OH* COx* C3H6 + * COx* COx + *

{ 1} {2 } {3 } {4} {5 } {6}

DISCUSSION Table 2 presents the results of the parameter estimation based on the above reaction mechanism for the two reacting mixtures. The solid lines in Figures 1,2 are those calculated by the model. An excellent agreement between the model and the experimental response curves is obtained at all temperatures. The propane activation step is found to be truly irreversible, the reverse rate being estimated not significantly different from zero. For the sorption equilibria both the forward and backward rates were estimated but at each iteration their values would increase while their ratio remained constant. This indicates that the sorption is too fast compared to the diffusion and that only the adsorption equilibrium can be estimated. A good fit for the propene pulse response was obtained only if the consecutive reaction towards COx was

138 included. The formation of COx from the alkyl species (step 3) was found not to improve the overall fits. This can be understood taking into account the fast rate found for step 2. Hence, although this route to COx cannot be excluded, a good description of the data is possible by considering only the consecutive oxidation of adsorbed propene. In that sense, the COx can be considered both as primary (direct formation from propane) and secondary (from readsorbed propene) products, as suggested by contact time kinetic experiments [2]. Table 2. Reaction rate and sorption parameters for C3H8/O2 = 0.5 and 2 between brackets step TOF at 575~ / S"1 pre-exponential factor enerboy / kJ mol"1 5.3 103 (1.3 10 3) m 3 m o l "1 s -1 * 45 (27) 80 (73) 2.1 10 TM(1.2 10 TM) s-1 3500 (7000) 175(166) 4.5 1011 (4.6 1012) s"1 680 (590) 143(160) A S a d s / J mol"1 K "1 AHads / kJ mol"1 4 -132(-129) -54 (-47) 6 -120(-120) -66 (-66) * Calculated with a number of V5§ sites on the surface of 700 mol m-3eat, corresponding to the total amount of surface vanadium units coveting the MgO phase [2,8]. From Table 2 it can be concluded that step 1 is rate determining. This corresponds well with the general picture of the oxidation of hydrocarbons over oxides where the cleavage of the first C-H bond is the rate-determining step. The propane adsorption rate at 575~ corresponds to a sticking coefficient of 3 109. Values of the same order of magnitude are derived from TAP experiments of butane partial oxidation over VPO catalysts [9]. This result underlines the very low probability of reaction on oxide catalysts as compared to the much higher values (10 4-10 -2) obtained for catalysis on metals. In order to check if this rate constant determined under TAP conditions corresponds to the data obtained under steady-state atmospheric conditions, the related conversion levels have been compared at similar residence times. For the same catalyst under oxygen lean conditions (C3Hs/O2 = 0.66), 10% conversion and 65% propene selectivity were obtained at a space-time of 70 ms under atmospheric pressure at 500~ [2]. Under TAP conditions, at similar residence time (an average residence time of 70 ms is obtained in the TAP reactor, though the actual contact time with the catalyst is shorter due to the two inert zones), a conversion of 15% and a selectivity of 64% are obtained at 500 ~ (Table 1). In a first approximation, a reasonable agreement is observed under the two operating conditions, which tends to validate the kinetic values derived from the TAP analysis. However, the slightly higher activity obtained under TAP conditions with a similar selectivity could indicate that more active sites are concerned under low pressure conditions. This is in accordance with a lower activation energy (80 vs 100 kJ mo1-1) obtained under TAP conditions. As mentioned above the surface coverages are very low under TAP conditions. Furthermore, the distribution of site activity on the catalyst is non uniform since a non uniform distribution of surface acidity strength was deduced from the rapid change in NH3 heat of adsorption as a function of coverage (Fig. 3A in [ 12]). This non uniform distribution of surface reactivity could explain that at low coverage, the most active sites are preferentially concerned. A kinetic mapping of the active phase could thus be established according to the surface coverage. The TOF and pre-exponential factors of step 1 are found significantly larger under oxygen rich than under oxygen lean conditions (45 vs 27 s1 and 5.3 103 vs 1.3 103 m3 mol1 s-1 at 575~ for C3HdO2 = 0.5 and 2.0, respectively). A main effect of intrinsic activity as discussed above can be discarded since it should have lead to a lower activation energy (the reverse trend, 80 vs 73 kJ/mol, is observed in Table 2). This difference in TOF values therefore

139 simply derives from a difference in the number of active sites, since this parameter is included in the TOF parameter through the pre-exponential factor. As a matter of fact, the partial pressure of oxygen directly monitors the concentration of V § atoms which are involved in the C-H bond activation. This change in surface concentration of active sites was confirmed by m situ DRIFT showing much more OH and carbonates groups under oxygen rich mixture than under oxygen lean conditions [unpublished results]. In this respect, and still considering a non uniform distribution of surface reactivity, the larger concentration of sites under oxygen rich conditions leads most likely to a slightly lower averaged activity per site, as reflected by the decrease in activation energy. The model was found little sensitive to the rate parameters of step 2 making the difference in rates irrelevant. The sorption equilibria parameters (steps 4 and 6) are reported in Table 2 as Van 't Hoff parameters. All adsorption entropies meet the criteria formulated by Boudart et al. [ 13]: 42 < ASads < Sgas and ASads -< 51 -0.0014 AHads (7) which hold for Langmuir adsorption. The sorption parameters for propene are rather close for the two feeds applied, which confirms that the state of the surface differs essentially by the concentration of active site. However, if the ratio of the kl values (about 4) for both feeds is taken as the ratio of active sites concentration, the adsorption entropy for propene under oxygen lean conditions is recalculated to -124 J mol 1 K "l, to be compared to -132 J mo1-1 K -1 under oxygen rich conditions. This indicates that the sorption of propene is slightly stronger on a more oxidized surface, which is in line with the lower averaged activity per site. The larger concentration in V 5+ Lewis acid sites would favor logically a stronger interaction with the double bond of the adsorbed propene precursors. The sorption parameters for CO and CO2 were found rather close and lumped into one set of parameters corresponding to step 6 (Table 2). They do not vary with the feed composition, possibly due to a much more uniform distribution of surface basicity, measured by SO~ adsorption, as reported in Fig. 3B in [ 12]. At that point it can be observed that the ratio of CO to CO2 selectivity tends to be rather stable and close to one at temperature where the trapping effect of CO2 by magnesia is reduced (above 550~ in Table 1). CO and CO2 appear to be formed from a common intermediate depicted as adsorbed propene linked to the surface by the double C=C bond (formed from step 4). It can therefore be speculated that a first CCH3 cleavage with O insertion would lead to CO formation (via formyl species), while the remaining C2 alkene would be step-wise oxidized into acetate, and/or formate leading to CO2 preferentially. All these intermediates would be in fast equilibrium with the surface, leading to the estimated sorption parameters. CONCLUSIONS Modeling of the TAP transient responses allows a fast and detailed insight into the reaction mechanism. Data presented in this paper correspond well with other kinetic and characterization studies on V/MgO catalysts as well as for other oxidative dehydrogenation reactions. Moreover all kinetic parameters are physically meaning~l in the framework of the transition state theory and Langmuir sorption. Thus the RDS activation parameters of propane activation reflects the state of the surface, depending on the operative conditions. However no information is obtained on the kind of oxygen species involved and open questions remain about the non selective routes leading either to CO or to CO2. ACKNOWLEDGEMENTS Financial support from the EC JOULE programme contract number JOE 3CT950022 is gratefully acknowledged.

140 REFERENCES Chaar, M.A., Patel, D., Kung, H.H., J. Catal. 109 (1988) 463; Kung, H.H., Adv. Catal. [1] 40 (1994) 1. Pantazidis, A. and Mirodatos, C., Stud. Surf. Sci. Catal. 101 (1996) 1029. [2] Lars, S. and Anderson, S.L.T., Appl. Catal., A: General 112 (1994) 209. [31 [4] Creaser, D. and Andersson, B., Appl. Catal., A: General 141 (1996) 131. Pantazidis, A. and Mirodatos, C., ACS Symposium Series, Washington, 638 (1996) [5] 207-222. Smits, R.H.H., PhD thesis University ofTwente (NL), 1994, ISBN 90-9006885-6. [61 Pantazidis, A., Bucholz, S.A., Zanthoff, H.W., Schuurman, Y., Mirodatos, C., Catal. [71 Today, 40 (1998) 65. Pantazidis, A., Burrows, A., Kiely, C.J, Mirodatos, C., J. Catal. 177 (1998) 325. [81 Gleaves, J.T., Yablonski, G.S., Phanawadee, P., Schuurman, Y., Appl. Catal. 160 [91 (1997) 55. [10] Svoboda, G.D., Gleaves, J.T., Mills, P.L., Ind. Eng. Chem. Res. 31 (1992) 19. [11] Marquardt, D.W.J. Soc. Indust. Appl. Math., 11 (1963) 431. [12] Pantazidis, A., Auroux, A., H e r r m a ~ J.-M., and Mirodatos, C., Catal. Today 32 (1996) 81. [131 Boudart, M., Mears, D.E., Vannice, M.A., 1967, Ind. Chim. Beige, 32, 281.

NOTATION A cross sectional reactor area AAr pulse response surface area C concentration DA Knudsen diffusion coefficient E activation energy AH~ adsorption enthalpy ka adsorption rate constant kd desorption rate constant k~ reaction rate constant L reactor length NA inlet pulse size Ns number of active sites

m2 R gas constant Vs t clock time mol mg"3 T temperature rng3 mr-1 s-1 SC3H6 selectivity kJ mol-~ AS~ adsorption entropy kJ mol1 XC3H8 conversion m3 mol-~ s-1 z axial reactor coordinate s-~ 8z Dime forcing function s-1 void fraction m TC3HS QMS sensitivity factor mol 0A fractional surface coverage mol rn~-3

J mol1 K -1 s K J mol-~ K-1 m ms3 mr-3 mol V-1 s-1 -

Reaction Kinetics and the Developmentof Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

141

T R A N S I E N T B E H A V I O R O F AN I N D U S T R I A L A C E T Y L E N E C O N V E R T E R Noemi S. Schbib, Alberto F. Errazu and Jos6 A. Porras 12 de octubre 1842, 8000 Bahia Blanca, Argentina e- mail: [email protected] FAX. 54-91-861600

Abstract The dynamic behavior of an industrial acetylene converter is discussed in the present work. The reactor is used to remove unwanted unsaturated hydrocarbons by means of a hydrogenation. This exothermic reaction is carried out in an adiabatic fixed bed reactor train in series (using a Pd/A1208 catalyst). Undesirable reactions accompany the main one. Therefore, the selectivity of the catalyst is very important. It is necessary to maintain stable operation while meeting product specifications for extremely low acetylene concentrations (<1 ppm). A computer simulation program for the industrial acetylene converter was developed. The steady states estimated by simulation are in good agreement with those found in industry. Besides, the simulated dynamic behavior of the converter shows the same general trends as those exhibited by industrial equipment. 1. INTRODUCTION The selective hydrogenation of acetylene in the presence of large amounts of ethylene is an important step in the ethylene manufacturing process. Most commercial installations manage to reduce the acetylene impurity to the desired specification effectively. In practice, this unit may have control problems when the undesirable hydrogenation of ethylene becomes important, leading to a runaway effect. In industry the acetylene converter can be located at different points in the purification section of an ethylene plant [1]. In one disposition the converter is placed after the conversion section (front-end). Another alternative involves the hydrogenation of the stream taken from the top of the de-ethanizer (tail-end). A great deal of research on acetylene hydrogenation has been undertaken. Most of it refers to kinetic studies under conditions similar to those at the taft-end and only a few papers study front-end conditions [2, 3]. Only some works analyze the steady-sate or dynamic simulations of the industrial process in particular. Sughrue et al. [4] have studied the dynamic behavior of a reactor with tail-end arrangement using Speedup software (Aspen Technology) for the dynamic simulation. Brown et al. [5] carry out the control and economic optimization of the taft-end hydrogenation process. They assume pseudo steady states for the mathematical model of the reactor. Hobbs [6] has analyzed the dynamic behavior and control of the front-end disposition. He simulated the reactor with lumped parameter model. In this work, the dynamic behavior of an industrial converter with front-end disposition is presented. It is based on the mathematical modeling of each unit of the converter (i.e. the condenser, the fixed beds and the heat exchangers). The main features of the feed of the front-end are: the high H2/C2H2 ratio @ 100) and the presence of several acetylenic, olefinic and diolefinic by-products as well as carbon monoxide. The CO is produced in the cracking furnaces and constitutes the main inhibitor of the hydrogenation reactions.

142 Figure 1 shows the scheme of the industrial converter simulated in this work, which includes a condenser (I1) and three adiabatic beds (R1, R2, R3) with intermediate cooling (12, I3). T2e Feed Fg

Ti e wt I y-,,i

~

"

T~ U

T2s

(

t

I

Tls T~

water

Produet

F i g u r e 1: Scheme of the industrial acetylene converter 2. MATHEMATICAL MODEL The hydrogenation process can be described appropriately by the reaction scheme shown in table 1. The kinetic model for the C2H2 and C2H4 T a b l e 1: Reaction mechanism hydrogenation was previously obtained for a C2H2 + H2 --->C2H4 commercial Pd/a-A1203 catalyst, under conditions similar to those at the front-end industrial C2H4 + tt2-->C2H6 hydrogenation process [2]. The conversions of C4H6 C3H4 +H2-'-~C3H6 and C3H4 are calculated by using an empirical C4H6 + H 2 - + C 4 H 8 model, that allow their estimation as a function of the acetylene conversion. Each stage of the acetylene converter consists of a short adiabatic fixed-bed reactor. The transient behavior of the three reactors (R1, R2, R3) has been simulated by a onedimensional pseudo-homogeneous model. The dynamic models for the condenser (I1) and the heat exchangers (12, 13) are first-order interactive systems, which result from the thermal balances at both, the process side and heating/cooling medium: M o d e l f o r reactors R1, R2 and R3 Kinetic M o d e l

M a s s Balances:

aci c~t

Z ( - ri )PL

~ a(,,ci ) e

8z

_

1

c

- r~ =

a+~C~Cn

[I+(KHCH)'/2 +K+
(1)

_1

(3)

i=A, E

j = 1,2... NC

z.

Energy balance:

8T

a 8T

Z (-Atr--Ii)(-r+)PL i

8t

6T 8 z

6T

(2)

= 1 - ( 1 - z~) ~"

y

z,, = 1 - ( 1 - z~) ~

(4)

143 Model for the condenser 11 Energy balance for the process side: p g VgC pg t i T ; : Fg2Cpg (Tge- T; ) + Q dt

Mass balance for the condensed fluid:

(5)

dh pcS ---~= Wv - FcPM

(6)

Energy balance for the heating fluid (steam):

(7) Models of the heat exchangers 12 and 13 Energy balance for the process side:

Energy balance for the cooling stream (wateO : (7/tCpwPw + MsCps) a ~w~ = FwCpw(Twe - Two)+p

(9)

To solve the steady-state model, the time derivatives of equations 1 to 9 set in zero. Steady-state results obtained from the solution of this mathematical model showed good agreement with industrial steady-state data [7]. 3. RESULTS AND DISCUSSION

The dynamics of each fixed bed and of the whole converter was studied for different step disturbances at the inlet of each unit. Starting from the initial steady state, individual step changes in inlet flow-rate, temperature, CO concentration and C2H2 concentration were carried out. The progress of the disturbances throughout the process can be explained by means of three types of waves: a "pressure wave", which is practically instantaneous; a "convective wave", which depends on residence time and is, therefore, fast Oust few seconds) because industrial flow-rate are high; and "thermal wave", which depends on the all thermal capacities (fluid + catalyst + reactor) and is relatively slow. Although these three waves have different propagation velocities, the time taken to reach the final steady-states is similar for different disturbances due to the relationship between temperature and concentration represented by reaction rate expressions. This effect complicates the study of the dynamic responses of the individual bed and the whole converter. 3.1. D y n a m i c S i m u l a t i o n s o f the Bed For the dynamic simulations, the axial coordinate (z) (eqs. 1 and 2) was discretized using finite differences. A "Gear routine" [8] was used to solve the system of ordinary differential equations. The behavior of each bed under the disturbances listed above was analyzed. In the following sections, the dynamic responses of the first bed under two of those disturbances, are explained in detail.

3.1.1. Disturbance in the feed temperature for the first bed (R1) The first bed was simulated under a +8% step in the feed temperature. The axial profiles for the temperature and concentrations of C2H2 and C2H4, at six different times are shown in figs. 2 to 4. The axial temperature profile in the final steady-state (t---~) is higher than the initial one (t=0) (fig. 2). Due to the higher reaction rate, this evolution is accompanied by a drop in the acetylene concentration profile (fig. 3). At the same time, a maximum

144 appears at the final ethylene concentration profile (t=oo) (fig.4) because this is an intermediate compound in the following series of reactions: C2H 2 rA ) C2H4 rE ) C2H6. The t e m p e r a t u r e profile changes gradually since it is related to the slow propagation rate of the "thermal wave". A quick change of the concentration profiles simultaneously occurs (figs. 3 and 4). These different propagation rates cause the inverse response exhibited by the outlet temperature (fig. 5). This effect, called "wrong-way behavior", has been observed before by other authors [9, 10] 100

0.5 t=

9O

0.4 t=0 s s

60

U' ,o

1

t'q 0.3 O 0.2

80

o~

70

0.1

60

,

30

,

,

60

f

r

,

,

90 120 z [cm]

,

,

150

0.0

r

180

Figure 2" Axial temperature profiles in R1. Disturbance in the feed to R1 (To = 69.5 75~

0

30

60

90 120 150 180 z [em] Figure 3: Axial profiles for C2H2 in R1 Disturbance in the feed to R1 (To = 69.5 75~ 100

52.0 t=_Qo ~51.9 r

95

P

~51.8

/

/

/

90

51.7

85 0

30

60

90 120 z [cm]

150

180

Figure 4: Axial profiles for C2H4 in R1. Disturbance in the feed to R1 (To = 69.5 75~

,

0

,

50

,

r

i

f

,

,

f

,

100 150 200 250 time [s]

r

300

Figure 5: Outlet temp. for bed R1 (z=180em). Disturbance in the feed to R1 (Te = 69.5 --~ 75~

3.1.2. Disturbance in the acetylene concentration in the feed The dynamic responses of the first bed under concentration changes were analyzed. For example, when the feed concentration CAe changes from 0.47 to 0.40 %wt the internal amounts of C2H2 and C2H4 drop immediately. The corresponding axial profiles for five different times are shown in figures 6 and 7. The final steady-state t e m p e r a t u r e profile (t=oo) is under the initial one (t-0). This is due to the lower rate of heat generation associated to the smaller concentration of the reactant (C~Hg. In contrast with the concentration profiles, the temperature remains unchanged during the first few seconds.

145 0.5

52.0 r

~-

0.4 51.9

~0.3

r

t'q

~9

0.2

~51.8 t---- OO

0.1

--Os

0.0

'

0

I

30

'

I

~

I

'

I

'

I

'

60

90 120 150 180 z [cm] Figure 6: Axial profiles for CzI-Ig. in R1. Disturbance in CAe=0.47 -~ 0.40 %wt

51.7

I

0

30

'

I

'

I

'

I

'

I

'

60

90 120 150 180 z [era] Figure 7: Axial profiles for CzH4 in R1. Disturbance in CAe=0.47 -~ 0.40 %wt

3.2. D y n a m i c s o f t h e w h o l e c o n v e r t e r

The disturbances enter the converter through the condenser (I1). To perform the simulation, the differential equations (1) to (9) must be solved. With the help of this model, the dynamic responses for temperature and conversions of C9H2 and C9H4 in each bed (R1, R2 and R3) were obtained, as well as the temperature profiles for the streams leaving the condenser and the heat exchangers. The "push-forward" interactive effect, that generates a variety of disturbances throughout the units, was also analyzed. 3.2.1. Disturbance in the feed temperature The behavior of the whole converter under a +5~ step change in the temperature of the feed entering I 1 was studied. This disturbance causes an increase in the outlet temperature for the condenser. This change moves upstream, affecting the first bed. In Figure 8 the dynamic evolution of the temperature leaving each bed (at z=180 cm.) is shown. An inverse response, similar to those described for the first bed (item 3.1.1) was found for all the beds. This effect is magnified and delayed for beds R2 and R3. The dynamic profiles for C2H2 and C2H4 at the outlet of the third catalyst bed (R3) were plotted in Figure 9. It is easy to note that the temperature increase in each bed produces lower C2H2 and C2H4 concentrations at the final steady state in comparison with the initial one. 3.2.2. Disturbance in the inlet CO concentration The carbon monoxide is produced in the cracking furnace according to the water-gas shift reaction. Its concentration is related to the amount of coke deposited during the pyrolysis process. It is known that, in addition to the conversion and cracker operating conditions, the sulfur additives have influence on the coke and CO formation [11]. As regards stability, the variations in CO concentration are the most dangerous disturbances. A small step change in the amount of CO at the inlet (Cco= 0.04 -~ 0.03 %wt), rises the temperature in all the beds (see fig. 10). In consequence, as it can be observed in figure 11, the outlet C2H4 concentration decreases almost 1% with respect to the initial steady-state value, while the C2H2 practically disappears. Another feasible situation during plant operation is that the CO concentration decreases beneath its normal value. A high negative step change (e.g. -60%) causes runaway conditions in the second and/or the third beds (see fig. 13). In this respect, our simulation results agree with industrial practice. A lower level of CO leads to higher hydrogenation rates with a sudden drop in the acetylene and ethylene concentrations. This effect is shown in figure 13, where the dynamic outlet concentrations for the third bed (R3) have been represented.

146

100

R1

f

/

96

%wt C2H2 8E-6

%wt C2H4 51.3

/

~51.2

R3

o~ 92 415-6

88

51.1

/

84 80 I 0

,

I , 400 600 800 time [s] F i g u r e 8: Outlet temperature of the reactors. Disturbance in the condenser (To: 35 ~ 40~

200

0

, 51.0 400 600 800 time [s] F i g u r e 9: Outlet concentrations for R3. Disturbance in the condenser (To: 35 ~ 40~ 0

200

,

% w t C2H4 51.2

% w t C2H2 4E-6 96 R3

51.0 2E-6

1

50.8

88

84

0

200

400 time [s]

600

800

F i g u r e 10: Outlet temperature of the reactors Disturbance in Cco =0.04 ~ 0.03 %wt

120

50.6 400 600 800 time [s] F i g u r e 11: Outlet concentrations for R3 Disturbance in Cco: 0.04 -~ 0.03 %wt 0

0

200

%wt C2H2 4E-6

%wt C2H4 52

112 ~ ' 104 o

R2

R

f 2E-6

48

96 88 0

40

80 time [s]

120

160

F i g u r e 12: Outlet temperature for the reactors Disturbance in Cco: 0.04 -~ 0.015 %wt

0

44 80 120 160 time [s] F i g u r e 13: Outlet concentrations for R3 Disturbance in Cco: 0.04 ~ 0.015 %wt 0

40

3.2.3. D i s t u r b a n c e in t h e inlet C2H2 c o n c e n t r a t i o n T h e c h a n g e in c o n c e n t r a t i o n t r a v e l s t h r o u g h t h e r e a c t o r s w i t h t h e s a m e velocity as t h e g a s flow. T h e r e f o r e , w h e n t h e C2H2 c o n c e n t r a t i o n drops from CAe =0.47 to 0.4 %wt, t h e

147 temperature of the first reactor (R1) diminishes. The temperature of the other reactors (R2 and R3) suffers the same effect but it is delayed, as a consequence of the dynamic effects of the preceding heat exchangers and the catalyst beds (see fig. 14). %wt C2H2 4E-5 -

90 l ~ 87 ~

RI

~"

~

84 ~ 81

0

%~ C2H4 51.25

R3

. . . . 400 600 time [s]

51.20

-

51.15

2E-5 -

R2 200

-

r

800

Figure 14: Outlet temperature for the reactors. Disturbance in CAe:0.47 --~ 0.40

0

I

0

200

i

400 time [s]

,

600

,

800

51.10

Figure 15: Outlet concentrations for R3. Disturbance in CAe:0.47 ~ 0.40

When the acetylene concentration diminishes at the inlet to the converter, the simulator predicts (as commented in section 3.1.2) a drop in the outlet concentrations of C2H2 and C2H4 in the first bed (see figs. 6 y 7). However, the lower temperature in each bed (fig. 14) leads to an increase in the final values for the C2H2 and C2H4 concentrations at the outlet to the converter, with respect to the initial ones (fig. 15). 3.2.4. Disturbance in the flow-rate The flow-rate change affects all units of the converter, practically at the same time since it moves along the converter following the "pressure wave". The dynamic response of the outlet temperature condenser (i.e. inlet to the first catalyst bed) for this disturbance is shown in figure 16. Then, two disturbances enter RI: an increase in the flow-rate and a decrease in the inlet temperature. The addition of both effects lowers the C2H2 conversion. Three disturbances enter R2: an increase in flow-rate and in the C2H2 concentration and a decrease in the temperature at the outlet of R1. Similar effects are produced in reactor R3. These disturbances lead to an increase in C2H2 and C2H4 at the outlet of R3 (fig. 17)

70 1

%wt C2H2 5E-5

/#

4E-5 t r..) o 69

%wt C2H4 51.35

3E-5

51.30 51.25

2E-5

51.20

1E-5

r

68 l . . . . 0

200

400 time [s]

600

800

Figure 16: Outlet temperature for I1. Disturbance in Fg: 72000 ~ 75000 NmS/h.

0

200

400

time [s]

600

51.15

800

Figure 17: Outlet concentrations for R3 Disturbance in Fs: 72000 --} 75000 Nm3/h.

148 4. C O N C L U S I O N S The m a t h e m a t i c a l model presented here allows the simulation of an industrial converter. The steady states obtained by simulation are in good agreement with those found in industry. The axial t e m p e r a t u r e and conversion profiles can be obtained as a function of time, after applying changes in the feed concentration of C9.H2 and CO and/or in the inlet flow rate and temperature. The dynamic behavior of the converter estimated by simulation exhibits the same general trends as those found in industrial practice. The t r a n s i e n t t e m p e r a t u r e profiles obtained for each bed when, for example, a t e m p e r a t u r e disturbance enters the condenser, indicate t h a t the inverse responses are several times greater t h a n those found from the independent studies of the dynamic responses of each mdividual bed. This fact m a k e s the dynamic model of the whole converter a powerful tool to predict t r a n s i e n t system behavior. 5. N O M E N C L A T U R E Cj concentration of j-th component in the gas phase, kmol/m 8 A CD~ti specific heat, J/kmol/K heat of reaction, J/kmol F flowrate, kg/s ki Arrhenius type rate constant of component i Ki equilibrium constant of component i NC number of components PM molecular weight Q rate of heat transfer, kcaYs S cross section of the condenser, m e ri rate of reaction of component i, mol/kgca t s t time, s T temperature, K T" by-pass temperature, K u average velocity of the fluid through the bed, m/s V volume, m s Wvt steam flowrate, kg/s wt weight fraction z reactor length coordinate, m

Subscripts A B c CO E Et g H e o P s v w

acetylene butadiene condensed carbon monoxide ethylene ethane gas hydrogen input output propadiene solid steam water

s pL p a

void fraction of packed bed bed density, kg/m 3 fluid density deactivation factor

Greek symbols

6. R E F E R E N C E S 1. Derrien M.L., "Catalytic Hydrogenation", Cerveny (Ed.), Elsevier Science, Amsterdam, (1986). 2. Schbib N. S., M. Garcia, C. Gigola and A. Errazu. Ind. Eng. Chem. Res. (1996) Vol 35, No. 5, 1496. 3. Schbib, N.S. "Din~maica y Control de un Convertidor Industrial de Acetileno". Tesis Doctoral, UNS. March (1998). 4. Sughrue E.L., R.L. Hair and R.J. CaUejas. AIChE Meeting, March (1995), 19-23. 5. Brown M.W., A_ Penlidis and G.R. Sullivan. The Can. J. of Chem. Engn. (1991) 69, 152. 5. Hobbs J.W., "Computer Control of an Acetylene Hydrogenation Process". Ind. Proc. Control Proc. Workshop, (1979). 7. Schbib N.S., &F. Errazu, J./~ Romagnoli, J.& Porras. Computers Chem. Engn. (1994), Vol 18, $355. 3. Gear C.W., "Numerical Initial Value Problems in Ordinary Differential Equations". Prentice Hall, Englewood Cliffs, N.S., (1971). 9. Van Doesburg H. and W.A, Jong, Chemical Engineering Science, (1976) 31, 45. 10. Gatica J. E, J.A_ Porras, A. F. Errazu and J. A Romagnoli. Chem. Engng. Comm. (1989) 78, 73-96. 11. Santiago J . ~ , Francesconi J.D. and N.L.Moretti. O/l Gas J. (1982) 81(39), 78-82.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science Ltd. All rights reserved.

149

Modelling of alkane dehydrogenation under transient and steady state conditions over a e h r o m i a catalyst using i s o t o p i c l a b e l l i n g . S. David Jackson, John Grenfell, Isobel M. Matheson l, and Geoffrey Webb I. Synetix, Research Technology & Engineering Group, PO Box 1, Billingham, Cleveland TS23 1LB, United Kingdom. tDepartment of Chemistry, Joseph Black Building, The University, Glasgow G12 8QQ, Scotland, United Kingdom.

Abstract The dehydrogenation of propane has been studied using isotopic labelling techniques coupled to transient flow methods. Using deutero-propane a kinetic isotope effect of 1"1= 1.85 was calculated and the rate determining step was identified as C3D7(ads) ~ C3D6(ads) + D(ads). By comparing the macro-scale model of the reaction system against the atomic-scale isotope tracer results it is clear that although the macro-scale was adequately described by stoichiometric equations, the atomic-scale processes are far more complex and are often poorly described.

1.

INTRODUCTION

The dehydrogenation of light alkanes has been known as a catalytic process for a significant number of years [1 ]. The catalytic processes fall into two categories, those based on platinum as the active phase [2, 3] and those based on chromia as the active phase [4]. However in both cases the catalytic process is complex with a series of competing reactions occurring simultaneously including significant carbon laydown resulting catalyst deactivation. Both catalyst systems require frequent regeneration to maintain activity and the heat liberated from the regeneration is used to balance the endothermic nature of the principal reaction. None of the current industrial processes are ideal and ways of improving the process are always being sought [5, 6, 7]. This can be done through catalytic chemistry, reaction engineering, or a combination of both. In this paper we report on a study of alkane dehydrogenation over a chromia catalyst. Although this type of system has been examined previously much of the work has concentrated on the nature, and electronic state, of the chromia [6, 8, 9], i.e. the catalytic chemistry. In this study we undertook, by using carbon-13 and hydrogen-2 isotope labelling, transient flow regimes, as well as continuous flow, to model the reactions taking

150 place in various sections of the catalyst bed on the macro-scale, as well as modelling the surface processes occurring under reaction conditions. In this way both the catalytic chemistry and the reactor engineering could be brought together. In this paper the differences between the surface processes and the overall stoichiometric processes will be discussed with respect to catalytic chemistry. In a following paper [10] the same modelling approach will be addressed to the reactor engineering.

2.

EXPERIMENTAL

The catalysts used in this study were a Cr203-K20/alumina and a Cr203/alumina, both prepared by impregnation. A solution of ammonium dichromate or potassium hydroxide and ammonium dichromate was added to the support (Engelhard A1-3992E, S.A. 200 m2g-1) and the resulting suspension evaporated to dryness. The solid was then calcined in air at 823 K for 3 h. Analysis of the catalysts gave values of 7.6 % w/w Cr and 0.7 % K for the ChO3-K20/alumina and 8.6 % w/w Cr for the ChO3/alumina. Two reactor systems were used in this study. Pulsed reaction studies were performed in a dynamic mode using a pulse-flow microreactor system with on-line GCMS. Using this system the catalysts (typically 0.3 - 0.5 g) could be reduced in situ in flowing 5% dihydrogen in dinitrogen (50 c m 3 min-~) by heating to 873 K at 10 K min~ and then holding at this temperature for 0.25 h. After reduction had ceased the catalyst was maintained, at the desired temperature, in flowing helium (70 c m 3 min-~). The reaction gases were admitted by injecting pulses of known size (typically 6.06 ~tmol.) into the helium carrier-gas stream and hence to the catalyst. After passage through the catalyst bed the total contents of the pulse were analysed by GCMS. Continuous flow reaction studies were performed in a 0.101 MPa, continuous flow microreactor with the gas stream exit the reactor being sampled by on-line GC. Using this system the catalysts (typically 0.3 - 0.5 g) could be reduced in situ in flowing 5% dihydrogen in dinitrogen (50 c m 3 min ~) by heating to 873 K at 10 K min ~ and then holding at this temperature for 0.25 h. After reduction had ceased the flow was switched to propane (50 c m 3 min ~) and the first analysis taken after 3 min, subsequent analyses were taken every 9 min. Catalyst regeneration was performed by flushing the reactor with dinitrogen at temperature before switching the flow to 5 % dioxygen in dinitrogen. The amount of carbon dioxide produced was continuously analysed by a carbon dioxide monitor. The system was held under 5 % dioxygen in dinitrogen at temperature until no carbon dioxide was detected. The catalyst could be cycled through this process of reduction/reaction/regeneration.

3.

RESULTS

After a catalyst was reduced, it was subjected to a series of transient reactant pulses each equivalent to one second of continuous flow. In Table 1 the results of sequential pulses of deutero-propane and propane passed over the catalyst at various temperatures are reported.

151 The presence of hydrogen in the first pulse of deutero-propane was shown to be related to residual hydrogen retained after the reduction process. If a catalyst was reduced in deuterium and a pulse of propane was injected, then deuterium incorporation was observed in the propane. In Table 2 the species identified at 773 K and 823 K from each isotopmer are specified.

Table 1. Propene as a percentage of the C-3 components. Pulse

773 K

823 K

873 K

C3D8

0

36

43

C3H 8

29

45

61

C3D8

0

30

41

C3H 8

36

45

61

C3D8

0

31

40

C3H 8

28

45

60

Table 2. Isotopic product distributions Pulse

Temp (K)

Eluant Gas

Isotopic species identified

C3D s

773

C3X 6

no propene detected

C3Hs

C3X8

C3D8

823

C3X 6

C3D6 C3DsH

C3X8

C3D8

C3DTH C3D6H2

773

C3X 6

C3H6

C 3 H s D C3H4D2 C3H3D3

C3X8

C3H8

C3H7D

C3X 6

C3H6

C3HsD

C3X8

C3H8

C3H7D C3H6D2

823

C3D7H

C3H4D 2 C3H3D3

where X = H or D

A typical product distribution from reaction at 773 K would be 3.7 % methane, 1.3 % ethane, 0 % ethene, 28.9 % propene, and 60 % unreacted propane, expressed as % carbon; the difference from 100 % represents carbon deposition. At 823 K the values are 2.4 % methane,

152 1.4 % ethane, 0.7 % ethene, 40 % propene, and 42.2 % unreacted propane. Both of these product distributions can be described by the following set of stoichiometric equations: C3H 8 C3Hs C3H 8 C3H 8 C2H4

(t) (2) (3) (4) (5)

r C3H6 + Hz --~ 3C + 4Hz --~ CH 4 + C2H4 ~ 2CH4 + C + H2 r C2H6

To confirm that these equations were valid not just in the first few seconds of catalyst life, a catalyst was run for 0.5 h under continuous flow conditions (Figure 1). Clearly over this time frame there is a significant change in product distribution.

f

5]

120

I '

~" 4 ]1- ~ r

d 3

Ii /

..... +--------o

....

<> .

.

.

.

.

-------

9

- 100 - 80 60

~2

g, o !

o~

40 20 0

-

0

_

-

500

" ....

T"~ ......

,~1

1,000

c ........

".........~.................,, ~ 0 1,500 2,000 2,500

T i m e on s t r e a m (s)

Hydrogen --

Methane

Ethane

,.

~

Ethene

Propane

t

~

Propene ....

m

.....

J

Figure 1. Change in product distribution with time on stream.

153 The product distributions early and late in the run were modelled against equations (1) - (5). Note that although it is not claimed that these equations are a unique descriptor of the process other equations were tried and did not give as a good fit to the distributions. Table 3 reports the product distribution as analysed by the above equations in terms of percentage propane reacting by the specified equation. The figure reported against equation (5) relates to the amount of ethene hydrogenated.

Table 3. Analysis of product distributions after 1 s and 1800 s on-stream. Equations

773 K

823 K

1s

1800 s

1s

1800 s

C3H 8 r

C3H6 + H 2

28.0 %

4.3 %

40.2 %

18.0 %

C3H s ~

3C + 4H2

6.4 %

0.2 %

11.3 %

1.0 %

C3H 8 ---> C H 4 + C2H 4

1.9 %

0.1%

3.2 %

0.7 %

C3H 8 ----> 2CH4 + C

4.6 %

0.1%

2.0 %

0.2 %

C2H4 + H2 ---> C2H 6

100 %

31.8 %

65.0 %

69.4 %

When [2-~3C]propane was fed the ~2C:~3Cratio for the methane produced was 8:1 at 773 K and 2:1 at 823 K. When an unlabelled pulse was passed over the catalyst after a labelled pulse the methane had a ~2C:13C ratio of 36:1 at 773 K and 15:1 at 823 K. A typical ratio for unlabelled methane was 88:1. In no case was labelled propene/propane detected in an unlabelled pulse.

4.

DISCUSSION

In this study the dehydrogenation of propane has been examined by transient isotope tracer techniques. The system has been modelled on the macroscale according to a series of stoichiometric equations (1) - (5) and these have been used to describe the macroscopic properties of the system [ 10]. However comparison between the macroscopic modelling and the isotope tracer studies reveals in some cases a discontinuity between macroscopic and microscopic processes. In this discussion we will examine the differences and show how the microscopic has implications for catalyst design, whereas the macroscopic, in a following paper [ 10], will be shown to have implications for reactor design. As can be seen from Table 1 there is a significant kinetic isotope effect, such that at 773 K no deutero-propane was detected. Calculation of the size of the kinetic isotope effect at 823 K

154 gave a value of 1"1= 1.85. This represents the maximum effect possible at this temperature as defined by Eyring [ 11 ]. An equivalent isotope effect was also observed in ethane dehydrogenation over a chromia catalyst [12]. A kinetic study [ 13 ] has attempted to model the propane dehydrogenation system and proposed C3Hs (ads) --~ C3H 6 (ads) + H2 (ads) as the rate determining step by a best-fit methodology. Clearly such a rate determining step would have a kinetic isotope effect, however at a molecular level the loss of two hydrogen atoms to give an H2 adsorbed species seemed to us unlikely. Hence to determine whether or not this was the rate determining step in our system the deutero-propane reaction was analysed at a time when no deutero-propene was formed (Table 2). This showed that only C3DTH was formed. Consequently a proportion of the deutero-propane undergoes the following processes: C3D s (g) r C3D s (ads) C3D s (ads) r C3D 7 (ads) + D (ads) C3D v (ads) + H (ads) r

C3D7H (ads) r

C3DTH (ads)

C3D7H (g)

Therefore this defines the rate determining step as being : C3D7(ads) ~

C3D6(ads) + D(ads)

The desorption of propene cannot be the rate determining step as this would allow the introduction of two hydrogen atoms into deutero-propane, nor can the rate determining step proposed in the earlier study [ 13 ] be correct as it would result in no incorporation of deuterium into the propane. This interpretation is in keeping with the study on ethane dehydrogenation, where either the loss of the first or second hydrogen (as two separate processes) was believed to be the rate determining step [12]. By 823 K, and at higher temperatures, there is rapid adsorption/exchange/desorption equilibria operating for both propane and propene, as well as the interconversion between propane and propene, resulting in multiply exchanged species. Hence by using the deutero-propane we have been able to delineate the rate determining step for equation (1) and show that the macro- and micro- are in alignment. From the literature [12] and our own results [14] it is clear that a similar argument can be applied to equation (5). 5 If we now consider equations (3) and (4) and examine the isotope tracer evidence with respect to the mechanism. At 773 K the ~2C:13Cratio of methane was 8:1, given the relative percentage conversions of equations (3) and (4) it is likely that only [12C]methane is produced from equation (3) and that equation (4) gives rise to a ~2C:13Cratio of 6.7:1. Therefore at this temperature the methyl components of propane have a far higher propensity than the methylene component to end up as methane. At 823 K the ~2C:13Cratio of methane was 2:1, indicating that all carbon atoms are equivalent on the surface. Consequently the most likely common surface intermediates are CH(ads) or CH2(ads), with total molecular rupture of the propane. Also at 773 K and 823 K the 12C:~3Cratio from methane produced during the following pulses indicated that carbon deposited from one pulse of propane can be hydrogenated and desorbed as methane during the passage of a subsequent pulse of propane.

155 Therefore even though the overall stoichiometric equations are good descriptors at both temperatures we find that the mechanism has changed significantly. Indeed at 823 K it is likely that equations (2), (3), and (4) are intrinsically linked with respect to the surface processes occurring. Such changes in surface reaction mechanism can clearly have considerable implications in relation to catalyst design, yet they have little impact on the overall macroscopic processes.

5.

REFERENCES

1. K. Kearby, Catalysis Vol. 3, P. H. Emmet (ed.), Reinhold, New York, p 453, 1955. 2. T. Hutson, Jr., and W. C. McCarthy, Handbook of Petroleum Refining Processes, R. A. Meyers (ed.), McGraw-Hill, London, 1986. 3 P. R. Pujado and B. V. Vora, Hydrocarbon Process., March, (1990) 65. 4. G. F. Hornaday, F. M. Ferrell, and G. A. Mills, Advan. Petrol. Chem. Refining, 4 (1961) 451. 5. R. J. Rennard and J. Freel, J. Catal., 98 (1986) 235. 6. S. De Rossi, G. Ferraris, S. Fremiotti, E. Garrone, G. Ghiotti, M. C. Campa, and V. Indovina, J. Catal., 148 (1994) 36. 7. S. D. Jackson, P. Leeming and J. Grenfell, J. Catal., 150 (1994) 170. 8. L. L. van Reijen, W. M. H. Sachtler, P. Cossee, and D. M. Brouwer, Proc. 3rd Int. Cong. Catal., W. M. H. Sachtler, G. C. Schmidt, and P. Zwietering (eds.), Vol. 2 (1965) 829. 9. C. Marcilly and B. Delmon, J.Catal., 24 (1972) 336. 10. E. H. Stitt, S. D. Jackson, F. King, and D. Shipley, paper this symposium. 11. H. Eyring and W. Cagle, J. Phys. Chem., 56 (1952) 889. 12. P. Tetenyi and P. Konig, Acta Chimica Acad. Scient. Hung., 89 (1976) 123. 13. I. Suzuki and Y. Kaneko, J. Catal., 47 (1977) 239. 14. S. D. Jackson and J. Grenfell, unpublished results.

This Page Intentionally Left Blank

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

157

The Development of a Model Capable of Predicting Diesel Lean N O x Catalyst Performance Under Transient Conditions GP Anseii, PS Bennett, JP Cox, JM Evans*, JC Frost, PG Gray, A-M Jones, M Litorell, RR Rajaram, G Smedler and AP Walker Johnson Matthey Technology Centre, Blount's Court, Sonning Common, Reading, RG4 9NI~ UK Abstract Steady state kinetics data from a commercial Pt-based lean NOx catalyst have been used to formulate a kinetic model to describe the performance of the catalyst. It is clear from this analysis that steady state kinetics m isolation are not sufficient to provide a full picture of the operational performance of such a catalyst. However, when this kinetic analysis is combined with mechanistic information obtained over the catalyst, the resulting model is extremely powerful. Within this paper, the development of the kinetic model is described, and the requirement for both accurate mechanistic information and detailed kinetic measurements is clearly demonstrated. The use of the model to predict the performance of a light-duty diesel vehicle under light-off conditions is described, and the power and flexibility of the model within the lean NOx area are emphasised.

Introduction Minimising the emissions of hydrocarbons (HC's), CO and NOx from motor vehicles is a major goal in today's increasingly environmentally conscious society. Today, all new gasoline-powered vehicles sold in the USA and the EC are fitted with so-called three-way catalysts which provide very high conversions of HC, CO and NOx, and lead to significant improvements in air quality. The diesel engine is an inherently low emission technology which, until recently, has not required secondary emission control such as catalytic converters. However, the impending introduction of more stringent legislation in the EC and USA will require the implementation of such systems. The diesel vehicle operates under strongly oxidising conditions, which means that the oxidative catalytic removal of CO and HC's is relatively straightforward. Unfortunately, the reduction of NOx under oxidising conditions is difficult, but the legislation demands that significant NOx conversion be achieved. In recent years Pt-based catalysts capable of performing "lean NOx" reduction within the temperature range of relevance to the light-duty diesel vehicle (150-300~ have been developed (eg [ 1]). While the performance of these systems is encouraging, the peak NOx conversion obtained over such a catalyst is around 40-50%, and this conversion efficiency can only be sustained over a relatively narrow temperature window (Figure 1). In the presence of such intrinsic limitations on the performance of the catalyst, it is clearly desirable to optimise 100 not only the catalyst, but also the rest of the emission control system, of which the NOx / . . . . . 80 catalyst forms a part. Such optimisation NO is only possible if a comprehensive .~_ 6 0 understanding of the characteristics of the catalyst under operating conditions can be g 40 developed. This paper describes the 0 construction of such an understanding, 20 and shows that the combination of accurate mechanistic information and _ _ 0 -: = detailed kinetics is essential if one is to 450 200 250 300 350 400 150 construct a model capable of predicting Inlet Temperature [C] the kinetic response of a lean NOx Fig. 1. The NO, NOx, and propene conversionprofiles obtained catalyst under actual operating over a 1% Pt/AI203 catalyst during temperature programmed conditions. reaction under a full gas mix. N.B. Above270 ~ NO -+ NO2

158

Experimental The steady state kinetics measurements were made using a spinning basket reactor which behaves as a continuous stirred tank reactor (CSTR) [2]. The commercial, Pt-based catalyst in monolith form was loaded into the reactor (4 blocks with dimensions of 40xl 0x5 mm at 400cpsi). The sample was exposed to a simplified gas feed (C3H6, NO, 02, CO and N2) at a gas hourly space velocity (GHSV) of 22,000hr-~. The speed of rotation (2500 rpm) was selected after ensuring that any further increases did not alter the measured steady state rate of reaction i.e. no mass transfer limitation. The conversion of propene was measured using gas chromatography and the CO and NO levels were monitored using a non-dispersive IR system and a chemiluminescent analyser respectively. The steady state experiments took the form of a stepped temperature ramp where the reactor was heated by an external furnace. The system was allowed to reach equilibrium at each temperature (all other inlet conditions remain constant). At an inlet temperature producing a fixed reactant conversion (typically 10%) a concentration span was carried out running from low to high reactant concentration, ie C3H6:300-1000 ppm, NO: 100-600 ppm. Temperature programmed reaction experiments were also performed using a monolith sample. The dimensions of this monolith core were 25.4 mm diameter by 38.1 mm long. Within these experiments, the catalyst was heated at a rate of 5~ from 120~ to 400~ under a simulated leanbum exhaust gas, at a GHSV of 30,000 hr1. This mixture comprised 200ppm NO, 200ppm CO, 400ppm C3H6, 12% O2, 4.5% H20 in N2. The gas phase concentrations were varied from this base case as follows: C3H6 200-800ppm; NO 100-400ppm; CO 200-5000ppm; 02 1-12%. In addition, the space velocity was varied between 15,000 and 60,000 hr-1 and the heating rate between 2 and 10~ -1. The engine emission data were obtained using a diesel 2.5L TDI engine run on a test bench (500ppm fuel S). The engine was set up without Exhaust Gas Recirculation (EGR) but with additional fuel injection into the exhaust. It is now well accepted that the HC levels emitted from diesel engines are not high enough to effect substantial lean NOx conversion. Therefore, additional fuel needs to be presented to the catalyst to boost the HC/NOx ratio. This was done by directly injecting fuel into the exhaust pipe, upstream of the catalyst. The HC:NOx ratio was kept constant within each test and was controlled to a set point by a fuel injection system The test protocol comprised a series of stepped lightoff tests; the temperature was increased by approximately 15~ at a time by increasing the engine load and the system allowed to reach steady state at each temperature. Exhaust gas was sampled for analysis using the same method employed during FTP/ECE vehicle te~ting i.e. constant volume sampling, where the exhaust gas sample is diluted with air to a constant volume to prevent water condensation in the pipework since exhaust gas contains a lot of water (typically 6%).

Modelling To enable comparison of the predictions of the propene and NOx conversions with the experimental data, a 1-dimensional dynamic mathematical model was written to describe the catalyst monolith sample used in the temperature programmed reaction experiments. We are using the model to extract kinetic data and Heck et al [3] have verified that a 1-dimensional model is adequate for this purpose. In addition, a fast turn around of simulations can be achieved using this approach, since 1dimensional models run relatively rapidly. The following assumptions have been made in the model: (a) Uniform flow distribution at entrance to monolith. (b) Negligible radial temperature and concentration profiles. (c) Adiabatic operation. (d) Transport of mass and energy in the gas is by convection. (e) Transport of energy in the solid is by conduction. (f) Mass and energy transfer between phases is accounted for using expressions by Ullah et al [4]. (g) No radiative transport of energy to or from the inlet/outlet monolith faces occurs. (h) No diffusion resistance is present in the washcoat. The following expressions are the one-dimensional equations describing a monolith reactor:

159 Gas mass balance"

OvCg, c3z

Gas energy balance"

+ k=Sv(Cg, C~j) = 0 (1)

OTg pg CpgV-~z + h Sv (Tg TA = 0 OT,

Solid energy balance" He) Ps Cp, Ot Solid mass balance

(le)

0 Csi

Ot

=

s (~Z2

(2)

+ ZHi RATE, Sa + hSv (TgTQ

(3)

= kmS,, (Cg, C~ORATE, Sa (4)

Where Ratei can be substituted by one of the kinetic rate expressions mentioned in the following sections of this paper. The partial differential equations above were discretized in the spatial domain using a finite difference technique to convert them into ordinary differential equations. These were then solved using a standard numerical Ordinary Differential Equation (ODE) solver. Results and Discussion

a)

"=

Steady State Kinetic Model

NO EXP

T .

Power law rate expressions have been used extensively in catalytic converter flow modelling studies. Figure 2 shows Arrhenius plots obtained from spinning basket reactor experiments over the commercial, Pt-based catalyst. The concentration span data were used to quantify the dependence of the reaction rates on the gas phase concentrations of the different reactants by linear regression analysis in power law form. These were then incorporated into the Arrhenius model to redetermine the activation energies and preexponential factors. The power law expressions for C3H6 and NO were found to be of the following form:

.

C3H6 EXP .

.

.

(Linear Fit) I

.

NO Rate IZ ..J -12

HC Rate

-14

-16 0.0021

~

I 0.00215

~.

~

J

I

0.0022 0.00225 1/Temperature [1/K]

~

! 0.0023

t

" 0.00235

Fig. 2. Arrhenius plots for C3H6 and NO rates. Temperature dependence of the lean NOx reaction rates from the spinning

basket reactor data showingthe linear regression analysisfor power laws.

The orders of reaction, a, b, c, and d are positive numbers between 0 and 1. The ~v,jC~v, C-NbO (5) power law equations show that propene RATEc3n, = Ac~m ex ~ Eac Rg Ts concentration has a positive effect on its own conversion but NO concentration has a negative effect on propene conversion over ( EaNo1 (6) R A T E No = ANO exp R gTs) Cc3m C a the conditions of the experiment. Propene also shows a promoting effect on NO conversion as shown in equation (6). Initially the power law equations (5) and (6) were inserted into equations (3) and (4) of the monolith model and the predicted reactant conversions were compared with the experimental data from the temperature programmed reaction experiments (Figure 3). Although the power law kinetics are capable of qualitatively predicting the observed experimental trends at low temperatures, it is clear that the derived expressions do not satisfactorily fit the data. This is particularly true for the NOx conversion profile which decreases with increasing temperature far more slowly in the model than it does

1

160 experimentally. It is obvious from this comparison that a more rigorous description of the kinetics is required. This has been achieved by obtaining a greater understanding of the reaction mechanism.

lOO _

80

m

_ m

m m

mllm__

r........

c 60 .o

"/ EXP NO b) Mechanistic Studies of the Lean NOx > 9e ,-- 40 Reaction o The mechanistic studies were o carried out using a Temporal Analysis of 20 Products (TAP) reactor [5]. Experiments designed to generate a full understanding 0 ~ mmm of the mechanism of the lean NOx 1 O0 150 200 250 300 350 400 reaction over Pt-based catalysts were Inlet T~.mn~.mttJre f('.l carried out; these experiments are Fig. 3. Power law kineticfit to leanNOx data. Comparisonof the described in detail in reference [6]. experimentally observed HC and NOx conversions with those Figure 4 outlines the mechanistic steps predicted in a 1-dimensionalmonolithmodel. derived from this mechanistic investigation. The hydrocarbon species reduces a patch of adjacent Pt atoms from Pt-O to metallic Pt. NO adsorption and subsequent dissociation then occurs on the reduced Pt sites. Once the NO dissociation has begun, there are two ways in which the adsorbed N adatom species can be removed from the catalyst. At low temperatures, where the rate of NO dissociation is low, most of the NO species associated with the surface will be present in the molecular form (rather than as dissociated N and O adatoms). Therefore, under these conditions, the most likely means of removal of the N adatoms is via reaction with molecular NO to yield NzO. As the temperature increases, the rate of NO dissociation increases and at higher temperatures, most of the NO-derived species associated with the catalyst surface will be in the form ofN adatoms. Under such conditions, the major route ofN adatom removal is by reaction with another N adatom to form Nz.

The key feature of the mechanism in the context of kinetic modelling concerns the temperature dependence of the NOx conversion. When the HC removes oxygen atoms to generate a reduced patch

161 of Pt, there is an immediate competition between NO and 0 2 to re-oxidise the Pt sites. At low temperatures, NO tends to win this competition, and significant adsorption of NO molecules occurs leading to NOx conversion. However, at high temperatures the oxygen wins the competition, partly because the dissociative adsorption of oxygen becomes increasingly rapid as the temperature is increased, and partly because the rate at which molecular NO desorbs from the catalyst increases rapidly as the temperature rises. As a result of this competition, the surface of the Pt particles is predominantly oxidised at the higher temperature and any NO conversion is oxidation to NO2 rather than reduction (to N2 and N20); the overall NOx conversion is then diminished (see Figure 1). This temperature dependence of the NO/O2 competitive adsorption explains the profile of the lean NOx conversion traces obtained over Pt-based catalysts. Our initial kinetic model based on power law expressions provides an inadequate description of the catalyst performance at high temperatures because this temperature dependent NO/O2 competition was not explicitly incorporated into the model.

C) Mechanistically based model Historically, mechanistically based modelling of catalytic reactions has involved testing estimated mechanisms against experimental data. Voltz et al [7] pioneered this approach by deriving kinetic rate equations for various possible mechanisms of heterogeneous catalysis and then using the equations which best fit the experimental data in a catalytic converter model. However, in this case we already have a mechanism for the lean NOx reactions and so can go straight to a specific formalism. The mechanistic studies reveal that the rate of NOx reaction depends upon the rate of reduction of the platinum surface i.e. the rate of NOx reaction is strongly dependent on the rate of the propene oxidation. Therefore the propene and NOx reaction rate equations were coupled together. Additionally, an improved kinetic description can be obtained by describing more fully the self and cross-inhibition effects that arise as the surface concentrations of reactants and products change. This is done by fitting the data from the reactor to equations in a Langrnuir-Hinshelwood formalism where self- and crossinhibition effects are incorporated as site adsorption terms in the kinetic expressions. Combining these two improvements gives : When these expressions for Ratei were used within the ~Eac~H,] Cc3m 1 (7) one-dimensional model, the RATEc~, = Ac,.m ext R, T] (1 + X Cc~,) (1 + Y CNO) agreement between model and experiment data shown KNO CNO (8) RA TE No = RA TE c~m in figure 5 was obtained. (1 + Z CNo) 100

The Langrnuir-Hinshelwood expressions allow a much better agreement with the experimental data than did the power law equations. During the light-off phase, the equations model the data very closely indicating that the kinetic parameters (i.e. activation energies and pre-exponential factors) are more accurate than previously. However, there is still a major discrepancy in the NOx rate equations since the model curve does not taper down as steeply as the experimental data points, highlighting the fact that there is still a major process being omitted from the kinetic equations.

80

P HC EXP NOx

oe- 60

L-H fit

| =~ 40 0o

~176176176176 o

o

20 0 100

J" 150

%

I

200

:

1

250

',

I ;=1;

300

Inlet Temperature [C]

=I

=-----I

350

400

Fig. 5. Comparison of experimentally observed HC and NOx conversions with those predicted using a Langmuir-Hinshelwood kinetic formalism

162 The mechanism suggests that this process is the high temperature competition between the NO and 02 for the reduced platinum sites. The NO/O2 competition for surface sites at higher temperatures causes a decline in the NOx conversion, so it should be accounted for as an inhibitory term within the kinetic formalism. The Langmuir-Hinshelwood type equations described earlier, which fitted the data satisfactorily at low temperatures and conversions, were modified to account for this process. Using the one-dimensional model and the experimental data an integral type analysis was performed to find values for the parameters shown in equations (9) and (10).

C ) RATEc~m = Ac~m exP Eac~.m X' Cc,mCo~ " Y' ~, RgTs (1+ Cc~vr)(l+ CNo) RATENo = RATEc~n,

(

(9)

(EaN~176I(I+Z'CNo)

Kxo

CNO

(10)

1 + Kxoaa,e x p l ~ l C o ~ '

LR

Equation (9) shows that over the ranges of oxygen concentrations tested there is a positive oxygen effect on the hydrocarbon oxidation rate even at very high levels of oxygen. Apart from this, the propene reaction rate equation has not changed significantly from equation (7). However, an exponential, temperature-dependent NO adsorption term has been added into the denominator of the NOx equation (10). This term acts to decrease the predicted NOx conversion sharply at higher temperatures, which is consistent with the experimental observations. Figure 6 shows a comparison of the predicted and observed propene and NOx conversion profiles using these improved kinetic expressions. The match between the model prediction and the experimental data is now very good. In addition, the model is capable of accurately predicting the effect of variations in reactant concentration, space velocity and heating rate over the whole range

T,)

100 _

_..,

mmm

m ~ -mmm -

_

80 EXP HC

,- 60 .9

l-

~ 40 o

"

EXP NOx L-H-M fit

20 0 100

9

150

200 250 300 350 Inlet Temperature [C]

400

Fig. 6. Comparisonof experimentally observed HC and NOx conversions with those predicted in a 1-dimensionalmonolith model using the Langmuir-Hinshelwoodformalism based on meehani.~tic information

of experimental conditions evaluated.

d) Application of the Model to Vehicle Tests Ageing and evaluating catalysts using engines and vehicles is costly and time-consuming. Often different permutations of exhaust system design and refinement leading to potential improvements in catalyst performance are overlooked because of the expense involved in considering them experimentally. This study has been aimed at obtaining kinetic expressions for different catalysts in the laboratory and employing them within an exhaust system engineering model. The model uses engine outlet data as input and predicts catalyst performance and outlet emissions levels. This capability would greatly compliment and serve to reduce the amount of engine and vehicle testing currently being carried out to evaluate catalysts, and can therefore lead to substantial reductions in costs. A one-dimensional systems model was developed containing (optionally) one or two catalytic monoliths and a length of exhaust pipe. The equations describing the heat and mass balance in a monolith are described earlier by equations (1)-(4). The assumptions and heat and mass transfer parameters used are the same as for the laboratory monolith model described above. The onedimensional equations for the pipe are as follows:

163 c3T. P. Cv~ Ot

=

02T~ Ri hi 2~ o,Z2. + R~ AR (Tg T.)

OTg 2hi PgCwv Oz + R, (Tg T~) = 0

Ro ho R. AR (T.

T~..)

(11)

(12)

The kinetic expressions for the hydrocarbon reaction rate were obtained using propene as a representative hydrocarbon. The hydrocarbon species emitted from the engine do not behave in exactly the same way as propene, so the kinetic 100 i I'1' I"' I' 'il I 9 I parameters in equations (9) and (10) needed to be refined to account for the different reactivity of the engine-out HC species. The form of the 80 equations was kept exactly the same but the = EXP E 'P HC value of the HC pre-exponential factor was o~ 60 altered by fitting to data collected from a stepped light-off experiment performed using the diesel 2.5 L TDI engine. The stepped light40 j EXP NOx off was achieved by increasing the engine load L) in stages and holding it at each point until a steady state was obtained. This was 20 approximated to a very slow temperature ramp in the model. A 0.144m x 0.152m monolith 0 was used with a space velocity of 60,000/hr. It 150 200 250 300 350 400 450 must be stressed that all of the other kinetic Temperature [C] terms used in the model were those obtained during the microreactor tests. Figure 7 shows Fig. 7. Comparison of HC and NOx conversions observed during the fit of the simulated light-off to the engine bench tests for the 2.5L TDI diesel engine with those experimental data. predicted by the model following adjustment of the HC preexponential factor to take account of the different reactivities of The systems model was then used to the engine-out HC species and propene. predict the performance of the arrangement shown in Figure 8. The system comprises two catalyst bricks separated by 0.4m of exhaust HC HC [. . . . [ ~J I [LJ' pipe. HC was injected before each catalyst to 1%~'~/~0] CatalYst i Catalyst ! 1 give a constant HC/NOx of 2:1 and the G.H.S.V. was 30,000/hr over the whole system. Figure 9 shows good agreement between predicted and experimentally Fig. 8. Dual brick exhaust configuration with intermediate pipe observed HC emissions during the vehicle used for simulations. test. Figure 10 shows the predicted NOx emissions over the vehicle test compared with the experimental values. Again the excellent agreement between predicted and the experimental data demonstrates the level of accuracy to which the mechanistically based lean NOx kinetic expressions have been determined.

i • 1

t"

"

Conclusions

Steady state kinetic data have been used to construct a model to predict the lean NOx performance of catalysts under dynamic conditions. A simple power law model fitted to the steady state kinetic data from a gradientless reactor, was not capable of predicting the NOx conversion profile, suggesting that critical features were missing from the model.

164 1000

600 .

' ~'

.

.

.

.

.

.

.

.

.

~ 600

I!;,

.00 :, r

.

0 400

3O0

ooo.

~ 200 Z

200

T

0

1001

0

2

4 6 Time/lO00

8 [s]

10

12

Fig. 9. Inlet, experimental and predicted catalyst outlet hydrocarbon concentrations over vehicle test

0

,

f'

2

,

I '",

t"

~

I'

4 6 8 Time/lO00 [S]

I

i

10

,I

12

Fig. 10. Inlet, experimental and predicted catalyst outlet NOx concentrations over vehicle test

Mechanistic studies revealed that the missing features related to the competition between adsorbed reactants and decomposition products, the most important feature being the strong competition between NO and 02 for reduced Pt sites at high temperatures. Once the mechanistic features were incorporated into the model, excellent agreement was obtained between the model predictions and experimental data obtained using a laboratory microreactor. Following the success of the modelling approach under laboratory conditions, the model was extended to describe the light-off performance on a diesel 2.5L TDI engine. Once again there was excellent agreement between the model prediction and the experimental data, reinforcing our confidence in the power of this modelling approach. It is clear that such a model can be used to optimise the design of lean NOx catalyst systems, to investigate HC injection strategies which maximise NOx reduction whilst minimising any fuel economy penalties. References

[1] [2] [3] [4] [5] [6] [7]

G. Smedler, S. Fredholm, J.C. Frost, P. Loof, P. Marsh, A.P. Walker, D. Winterbom SAE 950750 C.N. Satterfield, "Heterogeneous Catalysis in Practice", McGraw-Hill, New York, 1980,p.359. R.H. Heck, J. Wei and J.R. Katzer, A.I.Ch.E. Journal 2._22(1976) 477. U. Ullah, S.P Waldram, C.J. Bennett and T.J. Truex, Chem. Eng. Sci. 47 (1992) 2413. J.T. Gleaves, J.R. Ebner and T.C. Keuchler, Catal. Rev.-Sci. Eng., 30 (1988) 49. R. Burch, P.J. Millington and A.P. Walker, Appl. Catal. B: Env., 4 (1994) 65. S.E. Voltz, C.R. Morgan, D. Leiderman and S.M. Jacob. Ind. Eng. Chem. Prod. Dev. 12(1973)

Nomenclature

Ai, Ai', Ai" Pre-exponential factors Cgi, Csi Gas & Solid species concentration Cp~ Gas specific heat capacity Ea~, Ea~', Ea~" Activation Energies Heat of reaction for species i h Heat transfer coefficient Mass transfer coefficient kmi Pipe inner and outer radii R~,Ro R~ Logarithmic mean pipe radius AR Pipe wall thickness

Rg Universal gas constant Sv Specific surface area Sa Specific surface area of catalyst Tg, Ts Gas and Solid temperatures X, Y, Z, X', Y', Z' Adsorption parameters z Spatial axial co-ordinate 3' Porosity pgi, Psi Gas & Solid densities ~Ls,Lw Solid & Wall thermal conductivities v Velocity

This Page Intentionally Left Blank

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

167

M o n t e carlo and lattice-gas studies of the kinetics of h y d r o c a r b o n hydrogenation reactions A.S. McLeod and L.F. Gladden Department of Chemical Engineering, Pembroke Street, Cambridge, CB2 3RA, United Kingdom Abstract

The hydrogenation of unsaturated hydrocarbons by transition metal catalysts demonstrates complex behaviour that cannot be represented by mean-field kinetic models. We illustrate these kinetic phenomena by presenting the results of an experimental study of the catalytic hydrogenation of ethene by Pt/SiO2. Monte carlo and lattice-gas methods are then applied to the simulation of the ethene hydrogenation reaction. Both numerical approaches are shown to reproduce the kinetic discontinuities observed experimentally for the hydrogenation of ethene and ethyne. The monte carlo model is then extended to consider the hydrogenation of a number of alkenes of varying chain length. The simulations of the hydrogenation of both ethene and of the longer chain hydrocarbons suggests that the transition between the two kinetic regimes occurs at the point where the hydrocarbon surface coverage exceeds the dimer jamming limit. It is proposed that the origin of the discontinuity can be accounted for by assuming that the adsorption of the hydrocarbon approximates a random sequential adsorption process. This interpretation of the origin of the kinetic discontinuity is shown to be consistent with previous studies of hydrocarbon hydrogenation kinetics. 1.

INTRODUCTION

The use of mean-field models of catalytic reactions, in which the independent variables are the average surface coverages of each adsorbed species, can fail to capture the fundamental characteristics of a catalytic reaction. In this paper, we illustrate the potential limitations of mean-field kinetic models by considering the kinetics of the hydrogenation of C1 - C 4 alkenes. In the first part of this paper, we outline an experimental study of the hydrogenation of ethene by Pt/SiO2. In agreement with previous experimental studies of both ethene [1] and ethyne [2] hydrogenation a discontinuous transition between two kinetic regimes is observed. In order to aid the interpretation of these data, a monte carlo algorithm is introduced and is applied to the simulation of the Horiuti-Polanyi mechanism for the hydrogenation of alkenes. We show that by using a numerical method that explicitly accounts for the non-random distribution of the reactants on the catalyst surface the discontinuities observed experimentally can be accounted for by a kinetic model derived from the basic Horiuti-Polanyi mechanism.

168 4.0

L

E 3.0 z

[] 283 K (increasing pressure) 9283 K (decreasing pressure) 0 303 K 0 O0 O,..,O,q(}SQq~ 0 0 0000

n~D

=_ > O t'I..

@

= 2.0 I-

D

OmO

D~~ 1.0 2.5

3.0 3.5 4.0 In Ethene Pressure

d~r~

DN _~2LJD

4.5

Figure 1. Ethene hydrogenation catalysed by dispersed Pt/silica at 283K and 303K. The total pressure was 760 Torr, the hydrogen partial pressure was 100 Torr with helium carrier as the balance. While the monte carlo method provides an exact method by which to integrate the differential equation describing the steps of a catalytic reaction the method can be very computationally demanding. In order to address this limitation we introduce a lattice-gas model for the hydrogenation of hydrocarbons. Both lattice-gas and monte carlo methods demonstrate that the origin of the discontinuous kinetic behavior can be explained by assuming that the adsorption of the hydrocarbon on the catalyst surface approximates a random sequential adsorption (RSA) process [3]. At low reaction temperatures, the adsorption of the hydrocarbon on the catalyst surface is almost irreversible leading to the formation of a hydrocarbon overlayer which inhibits the dissociative adsorption of molecular hydrogen. The inhibition of hydrogen adsorption is found to result in a change in the rate limiting step from that of the hydrogenation of the half-hydrogenated intermediate to that of the adsorption of hydrogen. 2.

EXPERIMENTAL

STUDY OF ETHENE

HYDROGENATION

A 2.5 wt% Pt/silica catalyst was prepared by impregnation of a porous sol-gel silica support by an aqueous solution of hexachloroplatinic acid. The resulting material was dried under vacuum and subsequently calcined at 200 ~ followed by reduction in a stream of flowing hydrogen at 300 ~ The mean metal particle diameter of the catalyst, as determined by powder XRD, was 6.3 nm [4]. Reaction rate measurements were performed in a tubular quartz differential reactor at atmospheric pressure with helium used as a carrier gas. As the reaction is both rapid and exothermic, the catalyst particle size and gas flow-rates were chosen as to ensure neither heat nor mass transport limitation influenced the kinetic measurements. Turnover frequencies for the rate of ethene hydrogenation are presented in Figure 1. Data are shown for a hydrogen pressure of 75 Torr and reaction temperatures of 283 K and 303 K. For the low temperature experiment, data points were obtained first by increasing the hydrocarbon pressure, followed by reducing the hydrocarbon pressure. No evidence of hysteresis was observed. For the low temperature experiment two kinetic regimes separated by a discontinuity are apparent. At low hydrocarbon pressure the reaction order is found to be positive with respect to the /

169 hydrocarbon. In the low activity regime, the kinetic order with respect to the hydrocarbon is found to become negative, in agreement with the simulation predictions. Very similar reaction isotherms have been obtained previously for the hydrogenation of ethyne [2]. At the point just above the discontinuity in the high activity regime the turnover number is found to be 12.3 and 5.1 immediately below the transition. On increasing the reaction temperature to 303 K the discontinuity is not observed within the range of hydrocarbon pressures studied. 3.

A MONTE CARLO MODEL FOR THE HYDROGENATION HYDROCARBONS

OF

The general form of the Horiuti-Polyani mechanism for the sequential hydrogenation of an alkene is assumed to take the following form: ka

H2 + 2, ~ 2H*

(1)

kb kc

C2H4 -+- 2, --~ C2H~

(2)

ka ke

CnH~n nt- H* ~ Cng~n+l -+- *

(3)

kl

CnH~n+l -k- H* -~

CnH2n+2-~- (n +

1),

(4)

where the * denotes a vacant adsorption site and the * superscript an adsorbed species. In the monte carlo method each of the steps of the hydrogenation mechanism is simulated by a series of transformations of a lattice of discrete adsorption sites representing the catalyst surface. Starting from an initially empty lattice the lattice sites are occupied and vacated according to a set of elementary rules that implement each of the steps of the reaction mechanism. The adsorption of the hydrocarbon is represented by conducting a random walk on the lattice, occupying a connected group of adsorption sites [5]. The number of adsorption sites occupied by the hydrocarbon on the catalyst surface is assumed to correspond to the number of carbon atoms in the molecule. Hydrogen adsorption results in the occupation of a pair of neighbouring lattice sites and is assumed to be dissociative. The kinetic parameters were estimated from the study of Rekoske et al. [6]. As the time-scales of each reaction step differ by several orders of magnitude, an efficient event driven monte carlo algorithm is required. Using this approach all the possible reactions that can occur on the simulation lattice are recorded in a set of lists, with one list for each of the reaction steps. The rate at which each step in the reaction mechanism occurs depends on the configuration of the lattice and on the intrinsic rate of each reaction. Ignoring the influence of substrate heterogeneity, the rate of reaction for each fundamental step of the reaction mechanism is a product of the rate constant and the number of possible reactions for a given lattice configuration. Let na denote the number of possible reactions of type a, given a is one of the steps { a - g } of the reaction mechanism given by Equations (1)-(4). The probability of a given reaction occurring will be proportional to the number of possible reactions on the simulation lattice at any time, ha, and the rate of each event, ks, thus the probability of selecting a particular reaction, p~, will be given by

170

naka P~ -

E ~k~

(5)

The simulation algorithm can be summarised as follows. Starting with an empty square lattice of N = 1282 lattice sites, all the possible hydrogen and hydrocarbon adsorption events are determined and added to the appropriate lists. The first event is then selected and time incremented by an interval 5t according to equation Equation (6). The event lists are then updated. This procedure is repeated until the rate of reaction and the surface coverage of each of the reactants reach steady state. The rate of hydrogenation is defined by the turnover frequency, rg, of the final reaction step

Ng

(6)

rg = N A t

where N 9 is the total number of reactions of step g, that is the hydrogenation of the ethyl intermediate, that have occurred in a time interval At. The relation between real time and each monte carlo step was calculated assuming the occurrence of reaction events corresponds to a Poisson process [7]. 4.

R E S U L T S OF M O N T E C A R L O S I M U L A T I O N S

The results of monte carlo simulations of the ethene hydrogenation reaction are presented in Figure 2. A discontinuity separating the two kinetic regimes is predicted in accordance with the experimental data. The transition from the high to low activity regimes is promoted by a reduction in temperature, hydrogen pressure or increase in hydrocarbon pressure. This would suggest that the low activity regime is associated with a high hydrocarbon surface coverage and v i c e v e r s a . We may interprate these data as follows. If the transition between the two kinetic regimes is related to the hydrocarbon surface coverage, as has been previously suggested [8], then it would be expected that above a critical surface coverage a strongly adsorbed overlayer of hydrocarbon will prevent the dissociative adsorption of hydrogen. It is proposed that this critical surface coverage corresponds to the dimer jamming limit for the lattice as the adsorption of ethene is effectively irreversible at low temperature. Thus, the adsorption of ethene approximates an RSA process in which the hydrocarbon molecules are deposited until no further adjacent pairs of vacant adsorption sites exist on the lattice. In the absence of any surface reaction, the hydrocarbon surface coverage at this point would be equal to the dimer jamming limit of the lattice [3]. Once the hydrocarbon surface coverage is equal to the jamming limit no molecule occupying more that a single adsorption site can adsorb onto the lattice, irrespective of the rate of adsorption of that molecule. In the absence of surface diffusion the lattice then becomes jammed in a non-equilibrium configuration. At this point the catalyst surface is saturated with hydrocarbon, and the rate limiting step of the reaction mechanism must become the dissociative adsorption of molecular hydrogen. If the hydrocarbon surface coverage attains the jamming limit in the surface reaction controlled regime, i.e. the high activity regime, then there will be a discontinuous

171

10 [] 9

X

v A O

v A O

[]

[]

D

9

v A O

v A O

v A

v A

v A

v

O []

O []

O []

0

0

0

v

[] o i v

A

O

9

[]

o~

9270 300 330 360 390

K K K K K

(-

0

-5 -10

0

.

.

0

2

.

[]

0

[] 0

.

.

4

A O [] 0

6

In PC2H4

Figure 2. Turnover frequencies for the rate of ethene hydrogenation predicted by monte carlo simulation. Reaction rate isotherms are shown for a constant hydrogen pressure of 75 Torr. transition to a lower activity regime for which the rate limiting step is the dissociative adsorption of molecular hydrogen. It is however possible that the rate limiting step will become the adsorption of hydrogen before the hydrocarbon coverage reaches the jamming limit. In this case the transition will be continuous. It would therefore be expected that at high temperatures the equilibrium hydrocarbon surface coverage will be below the jamming limit and the transition between the two regimes will always be continuous. In order to confirm that lattice jamming is responsible for the kinetic discontinuity, monte carlo simulations have been used to study the kinetics of the Horiuti-Polanyi mechanism for the hydrogenation of longer chain hydrocarbons. The kinetic parameters used in this case correspond to those obtained for the ethene hydrogenation reaction at 300K normalised with respect to the rate of step f. The reactant gas was assumed to be composed of a binary mixture, the mole fraction of the hydrocarbon being Yh. In Figure 3 the rate of hydrogen adsorption, the first hydrogenation step, and the overall reaction rate are compared with the hydrocarbon surface coverage for a discontinuous transition (n = 2) and continuous transition (n = 4). For the smaller hydrocarbon the adsorption of hydrogen becomes rate limiting at the point where the surface coverage of the hydrocarbon reaches the jamming limit. The transition in this case is shown to be discontinuous. For the larger hydrocarbon, the adsorption of hydrogen becomes rate limiting step before the surface coverage of hydrocarbon approaches the jamming limit. The transition between the two regimes is continuous in this case. This is a feasible explanation as to why a discontinuous transition is observed for ethene and ethyne, while for larger hydrocarbons, such as butadiene the transition is almost continuous. These predictions support deuterium exchange studies of alkene hydrogenation reactions which demonstrate that in the low activity regime the adsorption of hydrogen is the rate limiting step, while in the high activity regime the rate limiting step is the first hydrogenation reaction [9,10]. Experimental studies have shown that the transition between the high and low activity regimes is discontinuous for the C2 hydrocarbons, ethene and ethyne, but continuous for a C4 hydrocarbon, butadiene, in agreement with our predictions [1,2,10].

172 10 -1

10 -1 a a a a a a

a

0.8 -~

10 .2

in

10 .3

0

~

10 .4

0.6 ~"

~

0

0.4 o

lO-S

|

10 -8 10-~

o o

o o o o o o o o o o

0

o oo~| 0.2

0.2

~ o~176 0.4'

oo

oo1:]oo

oaoo

10.2

E~

Yh

018

1

0

o

o

0.8

0.6

A &

-r

a

10 .3

0.4 o0 10 .4

0.2 o

0.6'

o

i0 -s

e~,-,~oO

0

0.2

(a)

o

~ ~

0.4

Yh

0.6

0.8

1

'~

0

(b)

Figure 3. A comparison of the rates of hydrogen adsorption ([-1), the first hydrogenation step (A) (left axis), and hydrogen surface coverage (~) (right axis) for random walk lengths of (a) n=2 and (b) n=4. The overall rate of ethene hydrogenation is shown as the solid line. Note the change in the rate determining step from the first hydrogenation reaction in the high activity regime to the adsorption of hydrogen in the low activity regime. 5.

A LATTICE-GAS MODEL FOR ETHENE HYDROGENATION

In the preceding discussion it was demonstrated that the transition between the two kinetic regimes arose as a consequence of lattice jamming with the strong adsorption of the hydrocarbon inhibiting the adsorption of hydrogen. Thus, in order to describe the kinetics of the ethene hydrogenation reaction it is necessary to be able to represent the spatial distribution of the adsorbed surface species. Previous studies of RSA on a square lattice [3] have demonstrated that a lattice-gas model accounting for nearest neighbour correlations provides a reasonable description of the adsorbed overlayer that results from the random deposition of dimers on a lattice. We would, therefore, expect a lattice-gas model to represent the kinetics of the ethene hydrogenation reaction as in contrast to the mean-field model the formation of a jammed lattice can be predicted. By introducing independent variables for all possible combinations of site pairs we explicitly account for all possible single site fractional coverages, ai, and dual site coverages, oij. The triplet terms, O'ijk, w e r e estimated using the Kirkwood approximation. In the present example the i, j and k terms denote the H*, C2H~, C2H~, and vacant surface species. These lattice states are denoted by subscripts H, E, Y, and O respectively. As an illustrative example, the balance equations for the temporal evolution of the H and HH pair terms are given by Equations (7)and (8), d(TH

dt d(THH

dt

=4(k~aoo--k2auu)--3(k6auy--ksauE+kTayy)

(7)

: k l a o o - k2aHH

+ 3(klaoOH + k2aHHH -- k7aggY -- ksagHE + k6agyy~./ \

(8)

173

I..

c:

-2 o

o o Diz] D o []

-7 -12 0.( )02

0.003 1000/T

0.004

[K-~]

0.005

Figure 4. A comparison of the results of monte carlo simulations (O) and a lattice-gas model (K]) for the hydrogenation of ethene. Results are shown for an ethene pressure of 50.9 Torr and a hydrogen pressure of 111.7 Torr. The first two terms in Equation (7) and the first four terms in Equation (8) result from the dissociative adsorption and associative desorption of hydrogen. The remaining terms account for each of the possible surface reactions. Details of the algorithms used to derive the balance equations can be found elsewhere [11,12]. The rate constants for each step, the k~ terms, are equivalent to those defined by Equations (1)-(4). Similar expressions can be written for the other possible single and dual site coverages to give a total of 14 independent equations. The corresponding sum rules,

i

o-,

=

1

E aij = a,, i

(9) (10)

were used as an independent check on the equation set. The resulting set of coupled differential equations was integrated numerically using a semi-implicit Euler method for stiff systems [13]. The hierarchy of kinetic equations being truncated using the Kirkwood approximation [14]. An Arrhenius plot comparing the monte carlo and lattice-gas solutions is presented in Figure 4, reasonable agreement is obtained between the two methods. Significantly, both the lattice-gas and monte carlo solutions predict the presence of a discontinuous transition separating the two kinetic regimes. In agreement with the monte carlo simulations of this reaction the point of transition is found to occur at the point where the ethene surface coverage attains the dimer jamming limit. Agreement between the two models is excellent in the high activity regime, this is unsurprising as the distribution of the reactants in the high activity regime has been shown to be random [8]. 6.

CONCLUSIONS

We have introduced monte carlo and lattice-gas models for the simulation of hydrocarbon hydrogenation reactions on transition metal catalysts. Detailed simulations of the hydrogenation of ethene have been conducted and compared to experimental data obtained for the hydrogenation of ethene by a Pt/Si02 catalyst. In

174 agreement with experimental observations, a discontinuity is found to occur in the reaction rate isotherm separating two regions of differing kinetic behavior. We have attributed this discontinuous transition to a random sequential adsorption process that leads to the formation of a hydrocarbon overlayer inhibiting the adsorption of hydrogen. Both monte carlo simulations and a lattice-gas model accounting for nearest neighbour correlations are found to reproduce the discontinuous kinetic behaviour that is observed experimentally. 7.

ACKNOWLEDGMENTS

ASM thanks Peterhouse, Cambridge and Rolls Royce for the award of the Frank Whittle research fellowship. LFG thanks the Innovative Manufacturing Initiative and EPSRC for financial support.

REFERENCES [1] S.D. Jackson, G.D. McLellan, L. Conyers, M.T.B. Keegan, S. Mather, S. Simpson, P.B. Wells, D.A. Whan and R. Whyman, J.Catal. 127, 10 (1996). [2] R.B. Moyes, D.W. Walker, P.B. Wells, D.A. Whan and E.A. Irvine, Appl. Catal. 55, L5 (1989). [3] J.W. Evans, Rev. Mod. Phys. 65, 1281 (1993). [4] A.S. McLeod, K.Y. Cheah and L.F. Gladden. Proceedings of the Seventh International Symposium on the Scientific Bases for the Perparation of Heterogeneous Catalysts, in press.

[5] A.S. McLeod and L.F. Gladden, J. Catal. 172, 43 (1998). [6] J.E. Rekoske, R.D. Cortright, S.A. Goddard, S.B. Sharma and J.A. Dumesic, J. Phys. Chem. 96, 1880 (1992). [7] D.T. Gillespie, J. Comp. Phys. 22, 403 (1976). [8] A.S. McLeod and L.F. Gladden, Catal. Lett. 43, 189 (1997). [9] D. Vassilakis, N. Barbouth and J. Oudar, J. Chim. Phys. 88, 209 (1991). [10] J. Oudar, Z. Phys. Chem. 197, 125 (1996). [11] Y.K. Tovbin, Prog. Surf. Sci. 34, 1 (1990). [12~ .,~. ~.{a., V.N. Kuzovkov and W. von Niessen, J. Chem. Phys. 100, 6073 (1994). [13] P. Deuflhard. Num. Math. 41,398 (1983). [14] T.L. Hill, Statistical Mechanics, Dover, New York, 1956.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

175

Time distribution of adsorption energies, local monolayer capacities, local isotherms and energy distribution functions on catalytic surfaces Ch. Abatzoglou a and N. A. Katsanos a, A. Kalantzopoulos b and F. RoubaniKalantzopoulou b aphysical Chemistry Laboratory, University of Patras, 265 00 Patras, Greece bDepartment of Chemical Engineering, National Technical University in Athens, 157 80 Zografou, Greece

Abstract Many catalytic phenomena are governed mainly by the adsorption energy on the heterogeneous solid surfaces. A very simple experimental arrangement, based on a combination of Head Space Analysis and Reversed-Flow Gas Chromatography, permits the determination of: (1) the adsorption energies e~ on various active sites i of the surface; (2) the local monolayer capacities for the adsorbed analytes on each i; (3) the local adsorption isotherms O~(p, T, e) on the heterogeneous surface; and (4) the adsorption energy distribution function f(e) on the surface. I. I N T R O D U C T I O N Adsorption of gases on heterogeneous surfaces plays an important role in many scientific and engineering areas, like heterogeneous catalysis. The energetically heterogeneous solids have been the subject of many papers and books recently [1-3]. One is impressed by the various ways used to overcome the difficulty or rather impossibility to obtain an analytical solution of the classical integral equation at)

O (p, T)= ~0, (p, T, e)f(e)de

(1)

0

in order to find the adsorption energy distribution function f(e) from the global adsorption isotherm t0(p, T) determined experimentally. Various approximations for the local adsorption isotherm O,(p, T, e) of Equation (1) and/or f(e) have been adopted, and since 1991 numerical solutions and estimation methods were used as recently reviewed [4]. The latter methods open another way to the problem solution, but they need powerful computers, not easily available everywhere. It should be noted that all above methods offer approximate functions or values only for the probability density function f(e), through approximate solutions of Equation (1), without any determination of the actual values of the adsorption energy e, the respective local monolayer capacities Cmax , and the local isotherms Oi(p, T, e).

176 The contribution of gas chromatography (GC) to the problem described above was so far to determine the global adsorption isotherm by providing retention volume data [5]. However, this so-called inverse GC, i.e., giving properties of the stationary phase by classical GC procedures, needs an extrapolation to infinite dilution and zero carrier gas flow-rate to approximate true equilibrium parameters. Here, the method applied allows an "out of column" adsorption equilibrium between the analyte gas A and the heterogeneous solid. In order that adsorption rate processes are also taken into account, a diffusion column filled with stagnant carrier gas intervenes between the gas-solid interface and the chromatographic column. Finally, the sampling procedure into the column may be based on the RF-GC (Reversed-Flow Gas Chromatography) technique. In fact, this combination belongs in the domain of HSA and provides a simple method to measure the adsorption quantities mentioned before, namely: (1) The values of the adsorption energies ei on the various active sites i of the surface. (2) The local monolayer capacities Cmax, i.e., the maximum adsorbed concentrations of the gaseous analytes on each kind i of adsorption sites. (3) The local (homogeneous) adsorption isotherms 0i(p, T, e) of Equation (1) on the heterogeneous surface. (4) The adsorption energy distribution functionf(e) on the heterogeneous surface. These important physicochemical quantities are determined neither through Equation (1) nor by means of numerical estimations, but via another pathway.

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

SECTION

An outline of the experimental arrangement is shown in Fig. 1. It is based on a traditional gas chromatograph, which, apart from the separation column and the detector, is equipped with a four-port valve for gases and a T shaped cell constructed from 1/4 in. glass chromatographic tube filled at one end with particles of the solid under study. The diffusion column contains only stagnant carder gas (nitrogen or air), which also flows through the empty sampling column, either from D1 t o DE or vice versa. A small volume (0.5-1 c m 3 at atmospheric pressure) of the analyte A under study is introduced into the solid bed through the injector at y - L2, by means of a gas-tight syringe. Then, a sampling procedure starts by turning the four-port valve from one position (solid lines) to the other (broken lines) for 10100 s, and then back again many times. This procedure creates narrow and fairly symmetrical peaks like those shown in Fig. 2 (samplepeaks) owing to the flow-reversal in the direction of the carder gas through the sampling column.

177 l' . . . .

!

inlet of carrier gas

fourvan

l

I

sampling~e~

I

reference injector

4~

b

!I

DI

Z It i diffusion [" l t ~ " c~ zd-,~ ,[~. P:8"/~" sol~ ~~ruo~o~ ~,u~ . Y=La--IIIII~ . _ l injector of

separation column or simple restrictor

detector

Figure 1. Schematic representation of columns and gas connections showing the principle of the present method. i

i

!

I

,

~

i

i

i

I

'

i

I

9

.

.!

:

!i

'

2~5

:'i '

i

!

i ! i

i

I i,

9

I I

:

7

i

I

-~

-

:

I

1

210

....!!i

time (min) Figure 2. Sample peaks of propene in nitrogen carrier gas, with section L2 (9.6 cm) containing 0.67 g TiO2, at 50 ~

178

3. T H E O R Y To find the adsorption energy e responsible for the local isotherm 0i(p, T, e) of Equation (1), only a local isotherm model is required and a fairly general one for our purposes is that of Jovanovic:

O(p,T,o~)=l-exp(- Kp)

(2)

going over to Langmuir isotherm in middle pressures, and to a linear form at low pressures. The K is Langmuir's constant given [5] by the relation

K=K~

(3)

Where R is the gas constant, and K~ K0 =

h3

. u s (T__~)

(2zc m) 3/2 (kT) 5/2 bg (T)

is described by statistical mechanics [6] as (4)

k being the Boltzmann's constant, m the molecular mass, h the Planck's constant, bg(T) the partition function for the rotations and vibrations of the free gas molecule and os(T) the partition function of the adsorbed molecule over all possible quantum states. As a low temperature approximation, we adopt that os(T) ~ bg(T), as was done before [5]. A mathematical model [7] based on two mass balances of the analyte A (in the regions z and y of the diffusion column), together with the rate of change of the adsorbed concentration Cs (mol/g) and the total isotherm for the equilibrium concentration c~ (mol/g) at time t, leads to the following function H-f(t): 4

H1/a4 = ~ A, exp(B,t)

(5)

i=l

where H is the height of sample peaks of Fig. 2, and M is the response factor of the detector. From the coefficients of time Bt calculated by non-linear regression with a PC program in GW-BASIC [7] of the experimental pairs H, t, various physicochemical quantities have been calculated, among which is a total adsorption rate constant kl (s l ) of analyte A on the heterogeneous surface. The fraction of the surface covered is denoted as 0 = c s /Cmax , Cmax (mol/g) being the local monolayer capacity, i.e., the maximum adsorbed concentration of the gaseous substance A at time t. Instead of the partial pressure of A, p, we write cyRT, considering A as an ideal gas and Cy(mol cm 3) as its gaseous concentration in the region y of the solid bed. Therefore, according to Equation (2)

179

0 = - -c~ T - - = ] - - exp(-KRTcy)

(6)

Cmax

From this and Equation (5), one can calculate the derivatives ac s* /OCy and a 2 c s* /OC2y of the adsorbed concentration c~ with respect to the gaseous concentration Cy of the analyte A for any chosen time. It is through these two derivatives that e, Cmax , 0,(p, T, e) and fie) are calculated, by means of the following relations:

s/aC y = KRT a s/aC E=

(7)

leT[lnIKRT[- l n ( R T ) -

, ay oL 1 s c s = - - k 1- -

as

,

Ai

gDl ,=1-~i

(8)

In K 0 ]

(9)

[exp(B,t)- 11

OCs*/OCy

,

Cma x =C s +

(10)

~

KRT

O,(e,p,T)=cs/Cma x 1 [KRT(ac:/at)

J( I :

La

where ay

=

as /)

---

LI g D1 A;, Bj

= = = =

(11)

a c:/ac a, o:/aCy] + a(KRT)/at-

KR~

(12)

cross sectional area in void space of solid bed (cm 2) amount of solid per unit length of tube (g/cm) linear flow velocity of carrier gas (cm/s) length of section z of the diffusion column (cm) calibration factor for the chromatographic detector (cm per mol/cm 3) diffusion coefficient of analyte A into the carrier gas (cm2/s) pre-exponential factors and exponential coefficients of time in Equation (5).

From the sample peak heights H and the respective times t of flow reversals (cf. Fig. 2), all physicochemical parameters e, Cm~x , 0i(p, T, e) and f(e) are calculated using a simple PC program based on Equations (7)-(12). Naturally, the experimental values of ay, as, o, L1, T and m are required, together with the chosen initial tl and final t2 experimental times. The parameters are calculated within this range at a chosen time interval, revealing the various

t I min

t I min

t I min

Figure 3. Plots of adsorption energies (a), local monolayer capacities (b), local isotherms (c) and energy distribution function (d), as functions of experimental time, at 29 oC, for dimethyl sulfide on Penteli marble, in nitrogen atmosphere.

181 active sites i of the heterogeneous surface. An example is shown in Fig. 3 with 3000 plot points of time, drawn by means of the program MATHEMATICA 3.

4. CONCLUSIONS By using an out of column adsorption equilibrium between the analyte gas A and the heterogeneous solid, the RF-GC technique can lead to an easy measurement of adsorption energies on various active sites, local monolayer capacities of the latter for the gaseous analyte A, local adsorption isotherms of A, and adsorption energy distribution functions on the heterogeneous surface of solids. All above can be achieved by a simple experimental arrangement and a simple PC program.

5. REFERENCES S. Ross and J. P. Olivier, On Physical Adsorption, Interscience, New York, 1964. M. Jaroniec and R. Madey, Physical Adsorption on Heterogeneous Solids, Elsevier, Amsterdam, 1988. W. Rudzinski and D. H. Everett, Adsorption of Gases on Heterogeneous Surfaces, Academic Press, London, 1992. N. A. Katsanos, R. Thede and F. Roubani-Kalantzopoulou, J. Chromatogr. A, 795 (1998) 133-184. M. Heuchel, M. Jaroniec and R. K. Gilpin, J. Chromatogr., 628 (1993) 59-67. R. H. Fowler, Statistical Mechanics, 2nd ed., Cambridge University Press., 1936, p. 829. Ch. Abatzoglou, E. Iliopoulou, N. A. Katsanos, F. Roubani-Kalantzopoulou and A. Kalantzopoulos, J. Chromatogr. A., 775 (1997) 211-224.

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Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science Ltd. All rights reserved.

183

Molecular dynamical studies of the mobility of benzene and water on silica surfaces: correlation with the influence of surface chemistry and morphology. Sean P. Rigby a and Lynn F. Gladden b "Synetix, P.O. Box 1, Biilingham, Cleveland, TS23 1LB, U.K. bUniversity of Cambridge, Department of Chemical Engineering, Pembroke Street, Cambridge, CB2 3RA, U.K. Abstract The close inter-relationship between silica surface chemistry and surface morphology has been studied using nitrogen sorption and DRIFT spectroscopy. The influence of the nature of the silica surface on the molecular motion of non-polar benzene and highly polar water on that surface has been studied using NMR techniques. It was found that rougher, more convoluted and microporous silica surfaces are characterised by vicinal hydroxyl groups more resistant to condensation under thermal treatment. Higher degrees of surface roughness facilitate the motion of molecules such as benzene which interact via dispersion forces. More corrugated, microporous silica surfaces also provide pockets of strongly interacting hydroxyls which, it is suggested, hinder the motion of diffusing water molecules.

INTRODUCTION The surface diffusion of species may significantly influence the overall rate and selectivity of catalytic reactions. The reactivity and use of silicas in catalysis is inherently linked to the presence of hydroxyl groups on their surface. The surface activity of silica depends strongly on the extent of hydroxylation and the nature of the hydroxyl groups. They are able to interact specifically (via hydrogen bonding and/or acid/base interactions) with adsorbate molecules depending on the functional groups possessed by a particular molecule. The molecular dynamics of a non-polar species, benzene, and a highly polar species, water, on the surface of a variety of different types of silica has been studied. The mobility of benzene has been studied, using Deuterium NMR, on two different sol-gel silicas and a fumed (pyrogenic) silica, and the mobility of water has been studied, by variable saturation experiments, on the two sol-gel silicas. The two sol-gels had very (by a factor of over three) different specific surface areas despite having similar specific pore volumes (see table 1). The results of the mobility studies have been correlated with studies of the surface chemistry of the silicas using Diffuse Reflectance Infra-red Fourier Transform spectroscopy (DRIFTs) and the surface morphology by nitrogen sorption. These results represent preliminary findings in an ongoing study. EXPERIMENTAL METHODS

Mercury porosimetry

Mercury porosimetry measurements were made using a Micropore Autopore II 9220. The sample was first evacuated to a pressure of 50 mm Hg in order to remove physisorbed water

184 from the interior of the sample. The results were analysed according to the Washburn equation [ 1] and the parallel pore bundle model. The pore volumes of the two sol-gel samples G1 and G2 (see table 1) are 0.95+0.05 and 0.94+0.05 cm3g-1 respectively.

Nitrogen adsorption-desorption Nitrogen sorption experiments were carried out at 77 K by use of a Micromeritics ASAP 2000 apparatus. The sample tube and its contents were loaded into the degassing port of the apparatus and initially degassed at room temperature until a vacuum of 10.5 Tort was recorded. A heating mantle was then applied to the sample tube and the contents heated, under vacuum, to 100 ~ for one hour. This procedure was repeated in 100 ~ steps (or a 50 ~ step in the case of the last step to 350 ~ until a particular target temperature was reached. The range of final temperatures considered was 300, 350 and 400 ~ The sample was left under vacuum overnight at a pressure of 6x10 6 Torr. At this point the heating mantle was removed and the sample allowed to cool down to room temperature. The sample tube and its contents were then reweighed to obtain the dry weight of the sample before being transferred to the analysis port for the automated analysis procedure. The sample tube was then immersed in liquid nitrogen at 77 K before sorption measurements were taken in the relative pressure region ofP/P 0 = 0.05 to P/P0- 1.00.

Diffuse Reflectance Infra-red Fourier Transform (DRIFT) spectroscopy The sample was ground into a finely powdered form. The spectra were recorded on a Nicolet Magna-IR 750 Spectrometer Series III. The same background was accounted for in each case and, for the purposes of comparison, the spectrometer was not realigned between samples. Each spectrum was taken with 128 scans and a resolution of 4 cm 1. The sample was heated to 500 ~ under a stream of helium at a flow rate of 50 ml/min, in steps of 100 ~ leaving the system to equilibrate for 30 minutes at each step. The sample had not been evacuated under a vacuum prior to the experiment, and the stream of helium was passed through a water trap before passing over the sample. A spectrum was taken after the equilibration time at each of the steps up in temperature as the sample was heated. After heating to a temperature of 500 ~ the sample was cooled down again to room temperature and the spectra to be presented were recorded. In each case, the spectrum thus taken was the same in form as that at 500 ~ The spectra were processed by the Kubelka-Munk algorithm to account for the scattering from the sample particles.

Deuterium NMR experiments The silica samples were degassed at room temperature to at least 10.5 Torr before slowly being heated to 400 ~ over a period of approximately one hour. The samples were then kept at that temperature under vacuum for l0 hours. The sample was then cooled to room temperature before an amount of fully deuterated benzene (Aldrich I-[PLC grade with >99.5 atom % purity) equivalent to a monolayer coverage was adsorbed on the sample. The sealed sample was then placed in an oven at 60 ~ for 2 hours to ensure homogeneous distribution of the benzene-D 6 in the sample. Deuterium NMR measurements were recorded on a Bruker MSL 200 N M R spectrometer operating at 30.72 MHz using a quadrupolar echo sequence with a 90 ~ pulse of 6.9 ms and a time interval of 40 ms between pulses. Deuterium lineshapes of each sample were recorded at temperatures between 160 and 370 K. The temperature was stable to +_1 K. The number of echo signals accumulated depended on temperature, and varied from 100 scans at higher temperatures to about 50000 scans at lower temperatures, with a repetition time of 0.5 s.

185

Determination of surface relaxation properties The spin-lattice relaxation rate (1/T~) of water confined in a sample pore space is enhanced due to interactions at the solid-liquid interface. The simplest physical model that can be used to interpret Tz data is the two-fraction, fast exchange model [2] which, assuming that diffusion of the fluid to the surface is much faster than the relaxation process, leads to an observed relaxation rate that can be related to properties of the pore structure: 1

T---l-=

"T--~IB+ ~oo"r,----~'

(1)

where T1B is the relaxation time characteristic of the bulk fluid @3 s), Tls is the relaxation time characteristic of a surface layer thickness, 2 (=0.3 nm [3]), and S and V o are the surface area and pore volume, respectively. In cases of variable saturation, where the void space is only partly occupied with liquid, and it may be assumed T~B>>TIsand a liquid layer remains wetting of the whole of the solid surface throughout the drying process, then the relaxation rate is given, approximately, by:

E

VVo 1 T1 = Vo ' X-S Tls

1 TI~

1-1 '

(2)

where V is the volume of liquid in the pores. The overall average pore volume Vo, and surface area, S, were determined from mercury porosimetry because the pre-treatment prior to the mercury porosimetry experiment consists of the evacuation of the sample at low temperature (i.e room temperature); previous studies [4] suggest that this sort of pre-treatment simply results in the removal of the physisorbed water layer. This observation would suggest that this pre-treatment is suitable in order to obtain a value for the surface area (ignoring percolation effects) appropriate for the study of the water interaction with the "as-received" (i.e. non-reconstructed) surface. Equation (2) suggests that the observed average relaxation rate, 7"I, will vary linearly with the fractional saturation (V/Vo) and T~s may be obtained from the gradient. The average spin-lattice relaxation time (T~) was determined using the inversion recovery method [5] performed on a Bruker Spectrospin MSL 200 spectrometer operating with a proton resonance frequency of 200.13 MHz.

RESULTS

Nitrogen sorption experiments The nitrogen sorption data for the three silicas considered was analysed in three different ways. The data for low values of relative pressure were analysed using a standard t-plot employing the Harkins and Jura [6] equation with standard parameters: I t=

13.99 ] 0.034 - l o g ( P / P o )

1/2

(3)

In this case, the values obtained for the t-layer thickness from equation (3), in the range between 0.3 5 and 0.5 nm, plotted against the respective volume of gas adsorbed are fitted to a straight line using least-squares regression. The y-axis intercept is related to the micropore volume and the slope is related to the external (i.e. meso/macropore) surface area. The values

186 of these parameters obtained by this method are shown in table 1. The total surface area of the samples was also found from standard B.E.T. analysis. The difference between the "external" surface area and the B.E.T. surface area is attributed to micropores. From the results for the external and B.E.T. surface areas given in table 1 it can be seen that, over all of the different temperatures to which the samples were initially pre-treated, only sample G2 has a significantly large amount of its subsequent surface area associated with micropores. The pore size distributions were then obtained from the nitrogen adsorption and desorption isotherms using the standard method of Barrett et al. [7]. The desorption isotherm is generally recommended for thermodynamic reasons [8]. It has been proposed [9,10] that the total volume of pores of diameters >_ 2r, V, obeys: -

dV~ dr

oc

r2_&

(4)

where d is the fractal dimension of the pore surface. The surface fractal dimension may physically take values in the range 2 _
3O

e

I

~

a M

1

~

b

u n

k .

3800

3"I00

.

.

.

--

.

.

.

.

.

.

""

_

3600 3i00 Wavenurnbers {cm "1) Figure 1. DRIFT spectra of samples (a) C1 (b) G2 and (c) G1; all recorded at room temperature after activating at 500 ~

187 Table 1 Nitrogen sorption experimental results Sample Temperature 400 ~

350 ~

C1 fumed silica

Micropore volume (CC g-l)

External surface area (m2g1)

B.E.T.B.J.H. surface fractal area dim. (m2g1) (cutoffs, nm)

-~0"

202 +/-4

201.2 2.507+/+/-0.8 0.008 a (199-10) 2.04+/0.04 a (9-4) 2.211+/0.005 d (47-4)

Micropore External B.E.T.B.J.H. volume surface surface fractal (cc g-l) area area dim. (m2g1) (m2g1) (cutoffs, nm) 0.003

181

194.0

+/-0.001

+/-3

+/-0.4

2.52+/0.01 a

(144-15) 2.11+/0.03 a (13-4) 2.20+/0.02 d (52-4)

G1 N0* sol-gel

309 +/-4

311.2 2.816+/+/-0.7 0.002 d (55-11) 2.747+/0.003 d

-~0"

300 +/-4

302.3 2.996+/+/-0.5 0.001 d (157-18)

G2 0.0046 sol-gel +/-0.0006

86.6 +/-1.4

98.9 2.304+/+/-0.5 0.013 a (17-4)

0.0045 +/-0.0007

86.5 +/-1.6

98.9 2.22+/+/-0.4 0.02 a (23-4)

(10-3)

Notes: a adsorption; d desorption. *No significant micropore volume observed. For a pre-treatment temperature of 300 ~ the external and B.E.T. surface areas are 346+9 and 343.3+0.7 m2g1, respectively, for G1, and 95+2 and 111.5+0.9 m2g1, respectively, for G2.

Deuterium and Proton N M R

Deuterium NMR has been used by a number of workers to study the molecular dynamics of benzene on a variety of catalysts [ 12]. The temperature dependence of the :H NMR relaxation characteristics of an adsorbed :H-containing molecule has been modelled assuming two types of motion of the adsorbed species: (i) anisotropic rotational motion about the hexad axis; and (ii) translational motion where benzene molecules reorientate isotropically by site exchange jumps. If the hexad motion is very fast compared to the isotropic site-exchange motion and the correlation time for the translational motion, z>> 1/O3o, where 030 is the Larmor frequency, then the peak width at half height, 6v, for the Lorentzian line is given by:

1

EIe2 l

(5)

188 where x can be considered the jump time between adsorption sites and (e=qQ/h) is the quadrupolar coupling constant and equal to 187 kHz. The activation energy, EA, for the jump process can be obtained by introducing an assumed Arrhenius-type expression: = x0exp (Ex/RT)

(6)

and examining the temperature dependence of the Lorentzian line width. For the temperatures when the deuterium spectrum is unequivocally a Lorentzian line the correlation times are plotted against temperature for samples G1 and G2 in figures 2 and 3, respectively. The motion of benzene on the surface of G1 may be modelled as a single activated jump process with an activation energy of 21.8+1.0 kJ/mol (straight line fit ra=0.9924). Alternatively, a better fit may be obtained to the data with a two motional regimes model, where above 210 K the activation energy is 24.8+1.1 kJ/mol (ra=0.9963), whereas below this temperature the activation energy is 16.8+0.2 kJ/mol (ra=0.9999). For sample G2, in the temperature range 190-290 K, the best fit to the data is a single motional regime with an activation energy of 21.7+1.4 kJ/mol (ra=0.9825). It has previously [13] been reported that sample C1 has a single motional regime with an activation energy of 20.7+3.5 kJ/mol and a pre-exponential factor, x0, of 1.04+0.19x10 ~= s. This data fit for C1 is shown as solid lines in figures 2 and 3 and it can be seen that the correlation times for benzene mobility are generally much smaller (implying more rapid jumps) for the two sol-gels than for the fumed silica. From the results of the variable saturation experiments it was found that the values of T~s for water (1HaO) on samples G1 and G2 are 39+3 and 22+4 ms, respectively. These samples are known to be manufactured from the same starting material and thus contain the same concentrations of paramagnetic impurities, if any, as each other. Differences in T,s values between the two samples may thus, solely, be attributed to differences in the interaction between the silica surface and the diffusing water molecules.

-14

-14

II

9 I,,,~

,...11"""

9

o -16

o -16

9 v-,,,l

C1

,-"

~

~ -18

O

O

-18 O"

-.. -20 3.5 3.8 4.1 4.4 4.7 5.0 5.3 5.6

Thousandths 1/Temperature (1/K) Figure 2. Variation of benzene motional correlation time with temperature for sample GI: [] 180-210 K, x0=6.6+0.1xl0"2s; 9 210-270 K, "r0=6.3+0.6x10~4s.

-9 -20 3.0

i

!

I

i

3.5

4.0

4.5

5.0

5.5

Thousandths 1/Temperature (1/K) Figure 3. Variation of benzene motional correlation time with temperature for sample G2 (e), -c0=4.4+1.2x1013s.

189 DISCUSSION Using the data obtained from the nitrogen sorption and DRIFTs measurements the following description of the silica surfaces and the differences between them is postulated. From the nitrogen sorption experiments it has been found that the sol-gel silicas possessed a larger fraction of their total surface area associated with micropores and/or a higher value of surface fractal dimension over lengthscales f r o m - 3 nm up to -23 nm, than the fumed silica. This would suggest that over lengthscales from-0.35-23 nm the surfaces of the sol-gel silicas are generally more irregular and convoluted than that of the fumed silica. The high fractal dimension for G1 would indicate that the anomalously high surface area of this sample is associated with the mesopore lengthscale. In the literature [ 14-16] it is generally thought that for the thermal treatment of silicas below 150 ~ dehydration and the fleeing of strongly hydrogen bonded water from silanols occurs. At about 250 ~ condensation of closely associated clusters of hydrogen bonded silanols occurs leaving residual, isolated silanols on gel surfaces. At around 450 ~ condensation of more widely separated surface silanols takes place and also closely associated internal silanols again leaving a residue of isolated groups. Previous work [14,15] would suggest that the sharp peak occurring at-3750 cm 1 in the spectra for C1 and G1 in figure 1 is due to flee, single, isolated hydroxyl groups. However the shoulder occurring at -3740 cm1 in the spectrum for sample G1 and the prominent peaks at 3743 and 3736 cm-1 for the spectrum for sample G2 may be attributed to vicinal hydroxyl species. The spectrum for pellet G2 also shows a shoulder at-3750 cm1 indicating the presence of some single hydroxyls on the surface of this silica. It is suggested that the vicinal hydroxyls are retained even up to 400-500 ~ on the surface of the sol-gel silicas because the surface irregularities and convolutions mean that pockets of hydroxyls, particularly those in the micropores of sample G2, are very closely interacting and are able to resist condensation up to these relatively high temperatures. Below 400 ~ the fumed silica has a relatively flat surface where condensation occurs more readily leaving single, isolated hydroxyls remaining. Sample G1 has a surface chemistry somewhat intermediate between those of C1 and G2. Since the spin-lattice relaxation of the hydrogen nuclei in the diffusing water molecules is mediated by molecular motion, the change of relaxation rate of molecules in the surface layer, compared with those in the bulk, characterises the solid-liquid interaction responsible for impeding molecular motion at the surface. The lower value for Tls for sample G2, compared with that for sample G1, suggests that the motion of the water molecules is more impeded on the surface of this sol-gel because they interact more strongly with this surface. The more restricted mobility of water molecules on convoluted surfaces may contribute to (and latterly depend on) the greater thermal stability of vicinal hydroxyls because the water from condensation reactions is less able to escape from the surface. We now consider the results of the 2H NMR studies in the light of the proposed differences in surface chemistry and morphology. It has been proposed that benzene interacts with the silica surface predominantly through dispersion forces [17] and that this interaction is strongest when the benzene molecule is able to lie with its plane flat against the silica surface. While the single regime model activation energies for the jumping motion on all three silicas are nearly the same, suggesting similar natures for adsorption sites, the actual correlation time for motion on the fumed silica is higher than for both the more structurally convoluted sol-gels. The pre-exponential factor in equation (6) accounts for various factors, such as motional geometrical/steric effects and the surface packing of the molecules. It is proposed that the jump probability of benzene is enhanced, relative to a flat surface, on a more convoluted structure for two reasons. Firstly, the molecule may not need to rotate as far

190 between nearest neighbour sites at small lengthscales. At low temperatures, the pre-exponential factor for both C 1 and G2 is smaller (indicating an enhanced jump probability aside from the activation energy considerations) than that in the two-motional regime model for G1. After heating to 400 ~ both C 1 and G2 have some micropores and are hence rougher at smaller lengthscales. Secondly, recent work [ 18] suggests that, with increasing temperature, molecules may make multiple hops or longer flights to non-nearest neighbour sites. On a convoluted surface there are rather more adsorption sites within an accessable jump range over larger lengthscales. In the two-motional regime model for G1, at higher temperatures, the pre-exponential factor is smaller than both C1 and G2. Over larger lengthscales the surface fractal dimensions indicate that G1 is significantly rougher than C1 and G2. The results reported above suggest that the spatial distribution of different types of hydroxyl groups may be segregated and correlate with surface roughness. The regions of different hydroxyls on the heterogeneous surface may have different interaction strengths with benzene. For G1, at low temperatures, the fastest correlation time is seen for a shorter lengthscale lower activation energy intra-region jump, but with increasing temperature the molecule may gain enough energy to move across the whole surface, characterised by a different activation energy.

CONCLUSIONS The preliminary results described in this paper would suggest that there is an intimate coupling between surface morphology and chemistry for silicas. The surface morphology and its effects on the chemical heterogeneity of the surface has been suggested as an explanation of the different molecular mobilities of benzene and water molecules on different silica surfaces.

REFERENCES

,

5.

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

E.W. Washburn, Phys. Rev. 17 (1921) 273. K.R. Brownstein and C.E. Tarr, J. Magn. Reson. 26 (1977) 17. F. D'Orazio, S. Bhattacharja, W.P. Halperin, K. Eguchi and T. Mizusaki, Phys. Rev. B 42(1990) 9810. C.E. Bronnimann, R.C. Zeigler and G.E.Maciel, J.Am.Chem.Soc. 110 (1988) 2023. E. Fukushima and S.B.W. Roeder, Experimental Pulse NMR: A Nuts and Bolts Approach, Addison-Wesley, Reading MA, 1981. W.D Harkins and G. Jura, J.Am.Chem.Soc. 66 (1944) 1366. E.P. Barrett, L.G. Joyner and P.H. Halenda, J.Am.Chem.Soc. 73 (1951) 373. S.J. Gregg and K.S.W. Sing, Adsorption, Surface Area and Porosity, Academic Press, London, 1982. P.Pfeifer and D. Avnir, J. Chem. Phys. 79 (1983) 3558. P. Pfeifer, D. Avnir and D. Farin, J. Stat. Phys. 36 (1984) 699. P. Pfeifer, Appl. Surf. Sci. 18 (1984) 146. B. Boddenberg and R. Burmeister, Zeolites 8 (1988) 480. P. Chiaranussati, PhD thesis, Cambridge University, 1993. P. Hoffmann and E. Knozinger, Surf. Sci. 188 (1987) 181. B.A. Morrow and A.J. McFarlan J. Phys. Chem. 96 (1992) 1395. S. Kondo, M. Muroya and K. Fujii, Bull. Chem. Soc. Jap. 47 (1974) 553. B. Bilinski, J. Colloid Interface Sci. 201 (1998) 180. J.S. Raut and K.A. Fichthorn, J. Chem. Phys. 103 (1995) 8694.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

191

Experimental Study of Reaction Instability and Oscillatory Behavior during CO Oxidation over Pd supported on Glass Fiber Catalysts I. Yuranov, L. Kiwi-Minsker, V. Barelko 1 and A. Renken* Institute of Chemical Engineering, EPFL, CH- 1015 Lausanne, Switzerland, llnstitute of Chemical Physics RAN, 142432, Chemogolovka, Russia

Abstract The catalytic ignition-extinction limits and oscillatory behavior of the CO oxidation at atmospheric pressure over Pd supported on glass fibers were investigated as a function of catalyst specific surface area, surface concentration and dispersion of Pd. The ignition-extinction limits for each specific catalyst depend on the CO concentration and the gas flow rate. Three reactivity regions were found as a function of temperature and CO concentration: high reactivity region with conversion close to 100%, low reactivity region and the region of multiplicity of steady states. In the region of multiplicity the self-sustained oscillations with a period up to 6 hours were observed for the catalyst with Pd loading of 0.2%wt. High and low reactivity states of the catalyst were associated with reduced and oxidized state of Pd on the glass fiber surface. It is believed that the reaction rate oscillations are due to the cycles of Pd partial oxidation-reduction.

1. INTRODUCTION Supported noble metal catalysts based on glass fibers (GF) as a catalytic support have already shown promise for the catalytic converter technology (Nicholas et al., 1976, Neyestanaki and Lindfors, 1995, Barelko et al., 1996, 1997). Palladium supported on GF showed high activity in complete oxidation of hydrocarbons and CO. GF materials consist of long length filaments with diameter of 3-9 l,tm. This specific form of the elementary unit allows to design a variety of regular macrostructure of woven fabrics for the GF catalysts. Different types of glass fibers in the form of tissue, gauze, 3D-blocks can be produced in this way as catalyst supports. The GF supported catalysts have advantages regarding to conventional pellets or monolith catalysts. They combine texture elasticity with an open macro-structure allowing to avoid big pressure drop of gases when passing through the catalytic bed. For useful applications of Pd/GF catalysts, the understanding of the details of the combustion reaction mechanism is important. The detailed knowledge of the region of multiplicity of the steady states and oscillatory behavior is also necessary for reactor safety and economic operations. The present work focused on the catalytic activity, ignition-extinction limits and the region of multiplicity of steady states in the CO oxidation at atmospheric pressure over Pd/GF catalysts.

192 The study explored the catalyst activity as a function of specific surface area and porous distribution of glass fiber support as well as surface concentration and dispersion of Pd. 2. EXPERIMENTAL

2.1. Catalyst preparation and characterization Two types of glass fibers were used as a starting material for the preparation of the GF supports: silica glass fibers (SGF) and alumoborosilicate glass (E-type) fibers (EGF) (produced by "Steklovolokno", Polotsk, Belarus). The specific surface area (SSA) of the starting glass fiber materials was 2 m2/g. The samples of EGF in woven form were pre-treated in acidic media in order to vary the SSA of the supports from 2 up to 250 m2/g. The pre-treatment conditions employed in each case depended on the chemical composition of the glass used and the SSA desired. Palladium (II) chloride PdCI2 (purum, Fluka Chemic AG, Buchs, Switzerland) was used as a precursor for the catalyst preparation. The supported Pd/GF catalysts were made by the deposition of Pd from appropriate solutions. The content of Pd was varied from 0.002 to 1.0%wt. The main characteristics of the catalysts are presented in Table 1. The specific surface area of the SGF support (catalysts 1-3) did not vary and was kept constant at 2 m2/g. The Pd surface concentration was varied on this support from 10-4 to 102 mmol/m2 by changing the total Pd loading. For the EGF supports (catalysts 4-6) the total content of Pd was kept constant (0.2%), but the supports with a SSA from 2 to 70 m2/g were used to vary the Pd surface concentration. The SSA of materials was determined by BET method at low temperature N2 adsorption by means of Sorptomatic 1900 (Carlo Erba Instruments).

2.2. Experimental set-up The experimental set-up consisted of three parts: the gas supply system, the reactor and the analytical system. The gases CO, CO2, N2, 02, and Ar (Carba-Gas, Lausanne, Switzerland) were used without further purification. The feed was supplied through mass flow controllers and the inlet flow rate was kept in the range from 0.1 to 0.2 1 (NTP)/min. The CO concentration in the feed was varied from 0.1 to 2.5% vol., the 02 concentration was kept at 10% vol. and argon was used as diluent gas. Outlet CO and CO2 concentrations were continuously monitored by an infrared analyser Ultramat 22P (Siemens). The reactor used for this study was a continuous fixed bed reactor with an external recycle loop and a membrane compressor (16 1 (NTP)/min). The recycling ratio in the system was never less than 100. Therefore, the reactor could be considered as an ideal continuous stirred tank reactor (CSTR) (Frank et al., 1997). Rolled up woven Pd/GF catalyst was placed into the middle part of the reactor. The temperature in the catalytic bed was measured by a sheathed thermocouple. The reaction rate R was calculated as the amount of CO moles reacted per second per gram of Pd. 3. RESULTS AND DICUSSION The catalytic properties of Pd supported on glass fibers were explored as a function of the total loading of Pd, the Pd concentration per unit surface and the chemical composition of the fibers used. All catalysts studied demonstrated the temperature hysteresis during CO oxidation as shown on Fig. 1. The reaction rate increased slowly with temperature up to the ignition temperature

193 (T~g), where the jump into the ignited state (high reactivity regime) was observed. Above the Tig, conversion of CO was close to 100% for all catalysts studied. If the temperature in the catalytic bed was decreased, the CO conversion of 100% remains until the temperature of extinction (Tox) is reached. 100

,o ~

=

--

/

60

f

~

/i regime

/ /

I

/

~/

O

/

.2%c0

I

/ /

]

I//

/li

/

20

II - 1%C0

/~',

/

~O 40

.

--

/

i

~

/

/

,, T,,.: ~,0.o

1~.....~ 220oc

o

190

"-,

,

200

,

,

210

.

i

220

,

Tig n= I

230

240

T, ~

Fig. 1. Temperature hysteresis for CO oxidation over Pd (0.02%)/SGF; Co2 =10%; Flow = 200 Nml/min. At the Tox the reaction enters the low reactivity regime (extinguished state). The Tox is always considerably lower then the Tis, showing the temperature hysteresis in this reaction. In between Tis and Tox the bi-stability region is attained and the catalyst can be in low or high reactivity states under the same reaction conditions. The temperatures Tig and T~x depend on the reaction conditions like the inlet CO concentration (see Fig. 1), the gas flow rate and the ramp of the temperature change. Under the same operating conditions the values of Tig and Tcx are useful to compare activities of different catalysts. For example, from the Table 1 it is seen that the T~s depends on the Pd surface concentration and the type of the support used. Thus, for the same support type (SGF) the Tig was observed to decrease when surface Pd concentration was increased from 1.0 xl0 4 to 1.0 x 10.3 mmol/m2 and then the Tis remains unchanged up to the surface Pd concentration of 1.0 xl0 -2 mmol/m 2. The influence of the support composition on catalytic activity is seen by comparison of the catalyst 3 with the catalyst 4 (see Table 1). These two catalysts were identical in Pd surface concentration, SSA and the total Pd loading. Only the chemical composition of the fiber support used (silica glass and alumoborosilicate glass) was different. The Pd supported on the SGF demonstrated lower activity (reaction rate more then 5 times smaller) and higher Tis with respect to Pd supported on EGF. When the same EGF type of support was used, a more complicated dependence of the reaction rate and Tig on the Pd surface concentration was observed. The T~s first decreases with increase in surface Pd concentration and then again goes up. This behavior could be due to the influence of the catalyst porosity and surface microstmcture on the dispersion and accessibility of the active Pd, which is reflected in the catalytic activity.

194 Table 1. Characteristics of the Pd supported on glass fiber catalysts. Catalyst

Pd cont. wt%

Type of glass fibers

SSA m2/g

0.002 Pd/SGF-2

0.002

silica

2

1.0 10.4

2.

0.02 Pd/SGF-2

0.02

silica

2

1.0 10.3

1.1 10-4

230

3.

0.2 Pd/SGF-2

0.2

silica

2

1.0 10.2

2.4 l0 -5

225

4.

0.2 Pd/EGF-2

0.2

alumoborosilicate

2

1.0 10 -2

1.3 10 -4

185

5.

0.2 Pd/EGF- 15

0.2

alumoborosilicate

15

1.3 10 .3

2.0 10-4

165

6.

0.2 Pd/EOF-70

0.2

alumoborosilicate

70

2.7 10.4

6.6 10-5

220

7.

1.0 Pd/EGF-70

1.0

alumoborosilicate

70

1.3 10.3

2.5 10-4

155

.

Pd

R

sur.con. mmol/m 2

(180~ molco/~Pd S

Tig~ 1%CO

255

Temperatures of ignition and extinction were seen to depend on the inlet CO concentration. The temperature-concentration dependencies for the catalysts studied are presented in Fig. 2.

CO inlet concentration, % Fig.2. Dependence of the ignition and extinction temperatures on the inlet CO concentration (n~at = 0.4 g; Q = 0.1 1 (NTP)/min; Co2 = 10% vol.). The regions formed by the upper curves of ignition and the lower curves of extinction represent the bi-stability regions (shown dashed). The region of low activity (extinguished state) is under the curve of extinction. The region of the high catalytic activity (ignited state) is above

195 the ignition curve. The extinguished state is characterized by comparatively low conversion (<40%). In the ignited state CO conversion close to 100% is always observed even for a small catalyst loading into reactor and Pd concentration of only 0.002%. In the bi-stability region the catalyst can exist in high or low state of activity under the same conditions. Crossing the bistability region by vertical or horizontal lines allows specifying the temperature - concentration hysteresis. The self-sustained oscillations in CO oxidation are well known and their appearance may be connected with the observation of the hysteresis phenomena (SchOth et al., 1993; Slin'ko and Jaeger, 1994; Imbihl and Ertl, 1995). During the present study the oscillatory behavior was detected only for the catalysts 0.02 Pd/SGF-2 and 0.2 Pd/SGF-2. These catalysts have been synthesized on the base of silica fiber support with a small SSA of 2 m2/g. Oscillations were not observed for the catalysts with alumoborosilicate support independently of the SSA. Fig. 3 shows oscillatory behavior observed over 0.02 Pd/SGF-2. lOO

80

"~

60

0

40

20 0

5

10

15

time,

20

25

h

Fig.3. Oscillatory behavior of the CO oxidation over 0.02 Pd/SGF-2 (rn~at= 0.8 g; Q = 0.21 (NTP)/min; Cco = 1.0% vol.; Co2 = 10% vol.; T = 220~ The oscillations appeared after about 2 hours operation for the system in a high reactivity state. The oscillation parameters like amplitude, period and shape changed continuously during the experimental run. Finally, the oscillations disappear and the system turns irreversibly into low reactivity regime. Due to this deactivation behavior only the ignition curve for 0.02 Pd/SGF-2 is shown in Fig. 2. The observed lost of activity may be explained by strong interaction of Pd with silica glass support under the reaction conditions. The catalyst 0.02 Pd/SGF-2 seems to contain palladium in highly dispersed form. The oxidized form of supported palladium in this case is able to react with acidic silica support, leading to the formation of an inactive palladium-silicate phase. This assumption was confirmed

196 by thermo-treatment of the catalyst. After heating during 0.5h at 500~ in air the catalyst 0.02 Pd/SGF-2 was observed to deactivate irreversibly. When the amount of Pd supported on SGF was increased in ten times (the catalyst 0.2 Pd/SGF-2), the deactivation under reaction conditions or after heating in air was never observed. Deactivation was not observed for the Pd supported on EGF, irrespectively of the Pd surface concentration and Pd dispersion. Alumoborosilcate glass fibers (EGF) are known to be more basic than silica (SGF). Therefore, we suppose that a strong interaction of Pd with GF support does not take place. The self-sustained oscillations of high regularity were observed only over 0.2 Pd/SGF-2 catalyst. The examples of these oscillations are shown in Fig. 4. The main characteristic of the observed oscillations is their long period up to 6 hours. 100 -

80

= O

I

I

J r

60

r~

I

0

20

,

0

I

5

,

I

10

,

I

15

,

i

20

,

I

25

,

I

30

time, h

Fig. 4. Regular self-sustained oscillation of the CO oxidation over 0.2 Pd/SGF-2 (n%at= 0.8 g; Q = 0.21 (NTP)/min; Cco= 0.75% vol.; Co2 = 10% vol.; T = 210~ Under the reaction conditions toward the end of the bi-stability region (close to the extinction curve on Fig. 2) the oscillations tend to loose their regularity. One example of this irregular oscillation behavior is shown in Fig. 5. The catalyst remained in a low reactivity state during a relatively long period of 5-17 hours with irregular variations in the observed conversion. After long time on stream the system returns abruptly to the ignited state, remains in this state for 2-3 hours and alter some more irregular variations in the CO conversion comes back to the extinguished state. The ability of palladium to be partially oxidized and reduced under the reaction conditions used seems to be the reason for the oscillatory behavior observed. In order to confirm the proposed oxidation/reduction cyclic mechanism, a catalyst was tested undergoing reduction and oxidation cycles. When the catalyst was in low reactivity state, the reaction gas mixture (CO + 02) was switched to pure CO (reduction atmosphere) for 1 hour. After this treatment the catalyst was observed to return to the ignited state.

197 100

"i

'I

,?

80

o~ 60 s 0

/

L_

0

> tO

40

U

20

.

0

I

20

.

I

40

,

I

,

60

I

80

time, h

Fig.6. Irregular oscillations of the CO oxidation over 0.2 Pd/SGF-2. (rn~at= 0.8 g; Q = 0.21 (NTP)/min; Cco -- 1.6% vol.; Co2 = 10% vol.; T = 230~ When the catalyst in high reactivity (ignited) state was oxidized by 02 for 1 hour, it demonstrated a low reactivity state. Therefore, the high and low reactivity states of the catalyst were seen to be associated with reduced and oxidized states of Pd respectively. The oscillatory behavior of the system is due to cyclic partial oxidation-reduction of Pd on the silica glass surface (Sales et al., 1982). 4. CONCLUSION The Pd supported on glass fiber is an effective catalyst in CO oxidation. It showed a very high activity with conversions close to 100% at short residence times of about 10-2 g.s.ml-a and Pd loading of only 0.02-0.2 %wt. The Pd/GF catalysts seem to combine texture elasticity with open macro-structure avoiding pressure drop in the reactor catalytic bed. The catalysts used revealed the existence of three temperature-concentration regions with different activities: a) a high activity region with the CO conversion close to 100%, b) a low activity and finally c) a bi-stability region. Within the bi-stability region the regular self-sustained oscillations in the CO conversion were detected for the 0.2 Pd/SGF catalyst. The chemical composition of the glass fibers, the Pd surface concentration and the Pd dispersion were found to influence strongly the catalyst activity and stability. The Pd supported on silica glass fibers with the Pd loading of 0.02% wt. was observed to deactivate irreversibly during the reaction. Deactivation was not detected for the Pd supported on alumoborosilicate fibers for the same Pd loading under equal reaction operating conditions. Besides the chemical composition of the support, the properties of Pd/GF catalysts depend on the preparation conditions and on the Pd salt used as a precursor. The dependence of oscillation parameters on the reaction operating conditions as well as the variations of catalyst

198 surface temperature during the oscillations will be investigated and reported elsewhere at a later date. ACKNOWLEDGEMENT The authors gratefully acknowledge the financial support from the Swiss National Science Foundation under the contract N~ The authors also would like to thank Dr. M. M. Slin'ko for critically reviewing of the manuscript and for useful discussions. NOTATION R - reaction rate, molco.gvdl.s1, S S A - specific surface area, m2.gl; rn~at- catalyst mass, g; Q - total flow rate, l.min-1; C - concentration, %vol., T - temperature, ~ x - conversion, %. REFERENCES:

Barelko, V.V., Khalzov, P.I., Zviagin, V.N. and Onischenko V.Ya. (1996)"Catalyst for Chemical Processes", Russian Federation Patent N ~ 2069584, Bulleten Izobretenii (in Russian), N~ Barelko, V.V., Khrushch, A.P., Cherashev, A.V. and Yuranov, I.A. (1997) Catalysis on the Eve of the XXI Century, Abstracts, Part II, p.59, Novosibirsk, Russia. Frank, B., Doepper, R., Emig, G. and Renken A. (1997) Bistability and Oscillations of the NO/CO reaction on a Pt/Mo supported catalysts. Catalysis Today, 1, 1063. Imbihl, R. and Ertl, G. (1995) Oscillatory Kinetics in Heterogeneous Catalysis. Chemical Reviews, 95, 697-733. Marengo, S., Comotti, P., Scappetura, S. and Vasconi, M. (1997) Experimental Studies of Transient Thermal Effects During Catalytic Oxidation in a Packed-Bed Reactor. Dynamics of Surfaces and Reaction Kinetics m Heterogeneous Catalysis, p.429-437, Elsevier, Amsterdam. Neyestanaki, A.K. and Lindfors, L.-E. (1995) Catalytic Combustion of Propane and Natural Gas over Silica-Fiber Supported Catalysts. Combustion Science and Technology, 110-111, 303-320. Nicholas, D.M., Shah, Y.T. and Zlochower I.A. (1976) Oxidation of an Automobile Exhaust Gas Mixture by Fiber Catalysts. Industrial and Engineering Chemistry: Product Research and Development, 15, 29-40. Schtith, F., Henry, B.E and Schmidt, L.D. (1993) Oscillatory Reactions in Heterogeneous Catalysis. Advanced Catalysis 39, 51-127. Sales, B.C., Turner, J.E. and Maple, M.B. (1982) Oscillatory Oxidation of Carbon Monooxide over Platinum, Palladium and Iridium Catalysts: Theory. Surface Science, 114, 381394. Slin'ko, M.M. and Jaeger, N. (1994) Oscillatory Heterogeneous Catalytic system. Studies m Surface Science and Catalysis, 86, Elsevier, Amsterdam.

Reaction Kinetics and the Developmentof Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

199

E n h a n c e m e n t of Selective Conversion in Spatially P a t t e r n e d Reactors A. S. C6t~, W. N. Delgass, and D. Ramkrishna School of Chemical Engineering, Purdue University, West Lafayette, IN 47907

Abstract Dual-functionality is introduced into packed bed reactors via structured spatial patterns in order to circumvent the limitations inherent in by-product inhibited reaction systems. Simulation results demonstrate the ability of patterned reactors to vastly enhance the selective conversion of several systems, some involving undesirable side reactions which punish the system for operating too hot. Comparisons between well-mixed and layered patterns indicate that the temperature-tuning flexibility of the novel layered configuration offers significant performance improvements over the mixed bed configuration for systems involving temperature-sensitive reactions whose optimum temperatures vary widely. 1. I N T R O D U C T I O N Spatial patterns are structured reaction networks in which different "regions" of a reactor contain different catalysts so that as the reaction fluid navigates its way from one "region" to the next, it cycles between different reactions. The goal of these patterns is to introduce an auxiliary reaction system in order to enhance the selective conversion of the primary reaction system. More specifically, for systems involving a desired reaction that is restricted from proceeding in an ideal manner, patterns offer a strategy for circumventing the performance limitation by manipulating the reaction mixture. In general, patterns add design flexibility by allowing one to tailor a reactor that directs the reaction along a prescribed course.

Figure 1. Pattern Configurations: (a) Well-mixed pattern (b) Layered pattern The "regions" of different catalysts may be of various lengths, and it is these characteristic length scales which define the patterned configurations of interest. The two patterns considered in this paper are (1) a uniformly-mixed pattern in which the two functionalities are distributed homogeneously and (2) a discrete pattern in which the two functionalities are organized as distinct layers (Figure 1).

200 Since one of the premises of this work is that this switching between reactions is advantageous, it is clear that the number of switches will affect the performance improvement. Based on this alone, one would anticipate the mixed configuration with its infinite number of switches between reactions to be superior to the relatively few switches of the layered configuration. However, the layered configuration possesses a distinct advantage over the well-mixed case because its broad zones may be maintained at different temperatures. Mixed catalyst beds are forced to compromise the performance of one catalyst or the other as the different catalysts are maintained at a common operating temperature. The layered pattern allows each reaction's operating temperature to be tuned to its own optimum so that conversion and selectivity can be maximized. This flexibility is critical in the case of highly temperature-sensitive reactions whose preferred temperatures vary widely. It is this feature which makes the layered pattern particularly interesting and allows us to take the concept of dual-functionality in a new, novel direction. A previous paper [1] investigated a system in which the restriction imposed on the primary reaction was due to chemical equilibrium. The exothermic, equilibrium-limited primary reaction: A+B<=~C+D (where D is the desired product) was enhanced by introducing an auxiliary drain-off reaction intended to eliminate by-product C. Studies on this simple two-reaction system indicated that both the layered and mixed bed schemes could vastly bolster conversion to the desired product. Furthermore, operating the mixed bed hot (thus elevating both reaction rates) was generally superior to switching back and forth between "hot" and "cold" zones in the layered configuration. The exception occurred when the equilibrium constant for the primary reaction fell dramatically with rising temperature so that the mixed bed was "punished" for being hot. This paper focuses on systems in which a side-product of the primary reaction inhibits the reaction rate. This situation is similar to the equilibrium-limited case in that, as the product accumulates, the primary reaction shuts off. Therefore, the secondary reaction is implemented to remove this constraint by eliminating the inhibitor. Unlike the equilibrium-limited case, there is no punishment for running the simple primary/auxiliary reaction system too hot. Therefore, undesirable side reactions are introduced into the reactor simulations in order to incorporate a realistic punishment for operating the reactor at excess temperature. Specifically, these undesirable side reactions compromise selectivity by either competing for reactants or degrading the desired product. Simulation results will demonstrate how the utilization of patterns allows one to bypass the selectivity constraints by tailoring a reaction network to minimize unwanted reactions. These productinhibited systems are also of interest because the nonlinear kinetics they involve may lead to reactor multiplicity that will significantly complicate the analysis. The superposition of multiplicity-induced patterns (in which different catalyst pellets operate on different steady state branches) on top of the spatial patterns described herein will be the subject of later work.

201 2. P L U G F L O W R E A C T O R ANALYSIS Several systems are considered in this investigation, and the analysis is initiated by studying a steady state pseudohomogeneous plug flow model which has been derived in detail in an earlier work [1]. The nondimensionalized material and energy balances are given by, respectively: dYi = sir i dz

for i = 1, 2, ...,n

(1)

dO= ~6iRiBi _h(O_Oc) dz

(2)

i=1

where 6i is the volume fraction of the catalyst for the ith reaction, 0c is assumed constant, and n is the n u m b e r of reactions in the system. 2.1 System 1 The simplest system examined involves only two reactions. The primary reaction creates desired product D and by-product C which inhibits further reaction: reaction 1: A + B ~ C + D (catalyst 1) The auxiliary reaction is added to remove the inhibitor: reaction 2: C + F ~ G (catalyst 2) allowing the primary reaction to proceed forward freely. For this simple system, the plug flow analysis is straightforward and will be explained in detail. For the sake of brevity, this detail will be withheld from the discussion of later systems. The reaction rates were chosen as: r1 =

klCACB

(1 + K A C A + K B C B + K c C c ) 2

r2 = k 2 C c C F

(3)

and then nondimensionalized according to the formalism described previously [1]. The volume fraction of each catalyst is constant with reactor position for the u n p a t t e r n e d (61 = 1 and ~ = 0) and mixed bed (61 = ~ = 0.5) cases while, for the layered configuration, zl = 1 and 6~ = 0 in the odd numbered zones, and the reverse is true in even numbered zones. Generally, the two reactions have different optimal temperatures so t h a t the layered configuration switches between odd numbered zones fed material at 01 and even zones at 02. The mixed and u n p a t t e r n e d cases are evaluated for a feed at 01, 02, and the average of these t e m p e r a t u r e s (0ave) in order to determine which temperature is most appropriate for each. In this specific case (with two simple reactions and no further restraints imposed), the system always produces its highest D yield when r u n at high 0 because rates are increased and there is no penalty for being hot. Therefore, there is no reason to switch between

202 high and low 0 in the layered scheme, and the layered, mixed, and unpatterned configurations are all run at a common elevated temperature. In order to demonstrate the benefits of removing C without introducing the complications of temperature effects, this simple system is modeled under isothermal operation. Figure 2 plots the extent of reaction profiles for the different schemes. The well-mixed pattern converts 62.5% of the reactants to D, the layered 47.7%, and the unpatterned 23.4%. Clearly, both patterned reactors significantly enhance the yield over the unpatterned case. For this simple two-reaction network (absent of a penalty for operating too hot), the mixed bed, with its more frequent switching between reactions, is always superior to the layered pattern. E x t e n t of R e a c t i o n

~

~

o.~

~

O.6

- - ........ . .

O.4

~

9 9 9

:

Profiles

layeredextant 1 layeredextant 2

mixed extent 1 mixedextent 2 unpatternedextent 1 . . ~

.

7

;

'

_ ~ ~

"

~ " ~ l , , ~ =~ . ~=~ ~ ~ ~ 1 . 9m e

' ...... '"i

Primary Reaction Zone

1--"iAuxiliary Reaction Zone

0.2

0 0

O. 1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

z

Figure 2. Extent of Reaction Profiles for Isothermal System 1

2.2 S y s t e m 2 In real systems, it is frequently the case that as temperature goes up, performance suffers as other limitations such as catalyst stability, coking, or the introduction of unwanted side reactions restrict the maximum operating temperature. This effect is studied by modeling systems that involve one of two different types of undesirable reactions which turn on and dominate at high 0. The first of these reactions competes with the primary reaction by utilizing reactants nonproductively. The specific reactions for this system are: reaction 1: A + B ~ C + D (catalyst 1) reaction 2: A + B ~ E (catalyst 1) reaction 3: C + F ~ G (catalyst 2) Maintaining catalyst 1 at low temperature minimizes the rate of reaction 2 so that valuable reactants are not wasted. Therefore, the unpatterned reactor and primary reaction layers of the segmented pattern are fed cold fluid (at 01 = 0) and exposed to a coolant while the secondary reaction layers receive hot material (at 0~ = 0.2) and are hot-walled. Running the mixed bed cold compromises the rate of the auxiliary reaction, but running it hot activates reaction 2. The relative weight of these factors determines the appropriate mixed bed operating temperature. When reaction 2 is relatively inactive, the well-mixed pattern does best when run hot, and its performance surpasses that of the layered pattern. However, when reaction 2 is active, running the mixed bed hot may not be advantageous due to the

203 domination of reaction 2. Figure 3 shows that in such cases both patterns can offer performance improvements over the unpatterned reactor, but the layered configuration is vastly superior to the hot uniformly-distributed reactor. The arrows in Figure 3a m a r k the D yield at reactor's end for each configuration (layered = 64.9%, mixed = 34.5%, unpatterned = 20.0%). In this case, running the well-mixed reactor at a compromise temperature, 0ave, improves its product yield to 54.3%. However, the layered pattern still maintains a much higher selectivity as it wastes only 6.9% of the reactants while the mixed bed wastes 45.6%. (a)

(b)

Extent of Reaction Profiles

0

o_1

02

03

0.4

0.5

06

o7

o.t~

0s

1

Temperature Profiles

0

o. I

0.2

03

OA

Z

_I 9I 9I I . . . ..,.,,

.

.......... - - -

0.5

0.6

0.7

o~

0.9

!

Z

l a y e r = extent 2 , l .a y.e r _ ~ l toxL~mt 3 mixed e x t e n t I i x = extent 2 mlxeo extent 3 pattemeO,,m,.t unpattemeo extent 2

"~

~lB

I--I

layered

mixecl unpatterneo

~Z] m i x e d

"I unpattemed

Figure 3. (a) Extent (b) Temperature Profiles for System 2 (0m= --02) Frequently, the 0-tuning flexibility of the layered pattern produces tremendously superior selective conversion regardless of the 0 at which the mixed bed is run. This is particularly the case when the unwanted reaction is very 0sensitive (as denoted by a high activation energy) so that the penalty for operating hot is stiff. Figure 4 depicts such a scenario. Here, the mixed pattern produces its best yield at 0ave but is still significantly inferior to the layered scheme in that it generates less D (layered = 50.5%, mixed = 23.1%) while wasting over five times as much reactant (layered = 14.3%, mixed = 76.8%).

(a)

(b)

Extent of Reaction Profiles

0

o. 1

02

0.3

0.4

0.5

0.6

07

0.8

0.9

1

Z I

9 9 9 I 9I 9 I

- - . . . . . .

........ - - -

layered e x t e n t I layered e x t e n t 2 layered e x t e n t 3 mixecl e x l e n t 1 mixeO e x t e n t 2 mlxea extent 3 unpatternecJ e x t e n t 1 unpattern~l extent 2

Temperature Profiles

0

o. I

02.

0.:3

0.4

0.5

o6

0.7

0.6

o.g

Z

(I

layered mixed

111 ,~

l-- ,,~- I mrr~l unpattemea

unpatternexl

Figure 4. (a) Extent (b) Temperature Profiles for 0-Sensitive System 2 (0m=- 0ave)

1

204 2.3 S y s t e m 3 The second undesirable reaction implemented in the analysis is one which degrades the desired product D through further reaction with B. The complete system of equations is defined by reactions 1 and 3 of System 2 with reaction 2 replaced by: Reaction 2: D + B ~ E + C (catalyst 1) Once again, due to its high activation energy, reaction 2 generates a significant adverse effect only at high temperature. In this case, the punishment for operating hot is twofold as product is destroyed and additional inhibitor created. Clearly, then, it is critical to minimize this reaction so that running the mixed bed hot is frequently unacceptable. Figure 5 plots the results of a simulation involving a highly 0-sensitive degradation reaction that dominates at high 0. Here, the hot mixed bed actually performs worse than the unpatterned case because more than 90% of the D generated in the mixed bed is degraded at the elevated temperature. With its temperature-tuning flexibility, the layered pattern is able to offer a significant yield enhancement. While the mixed pattern's selective conversion can be improved for such 0-sensitive systems by operating it cooler, the layered pattern will generally offer superior performance regardless of the mixed bed's temperature. (a) (b) Yield of D Profiles

0

0.1

0.2

0.3

0.4

0~5

0.6

0.7

0.0

0_~1

1

0

0.1

0.2

03

0.4

m m i 9 9 IIIIIIII , ," ", ,,

........ - - -

m y e r e d exlm111 iaywr ~ t 2 layerer ~(tB1t 3 mixed extent 1 mixecl e x t e n t 2 nixed extent 3 u n p s t t e r n e d recent 1 unpattemed extent 2

o.5

o.6

0.7

o~

o.s

Z

Z ---

ymla: m~ea D y]elcl: unpa~ernecl

Figure 5. (a) Extent (b) Product Yield Profiles for 0-Sensitive System 3 (0~=

= 0ave)

If R2 is only moderately 0-sensitive, both patterns can provide benefits over the unpatterned case. This is demonstrated by Figure 6, generated for a mixed bed operated at intermediate 0. For this system, both patterns nearly triple the product yield produced by the patternless reactor (layered = 36.6%, mixed = 33.9%, unpatterned = 12.8%). However, the layered pattern provides a superior selectivity by minimizing reaction 2, converting only a few percent of the reactants to unwanted E rather than the nearly 20% with the mixed configuration. It is not always the case that the layered pattern is superior. In fact, when reaction 2 is even less 0-sensitive or relatively inactive, the advantage of the mixed pattern's infinite number of reaction switches outweighs the 0-tuning advantage of the layered bed. For such situations, it is not uncommon for the mixed bed to double or even triple the yield enhancement provided by the layered pattern. It

205 follows, then, that the choice between patterns will largely depend on the 0sensitivity of the reactions involved.

(a)

(b)

Extent of Reaction Profiles

Yield of D Profiles

. . . . . . . .

0

(11

02

0.3

0.4

0.5

06

0.7

0.8

o_s

0

o. 1

o.z

0.3

0.4

0.5

Z 9 9 9 I II I I I I I

0.6

0.7

0.8

o.~s

Z

- - -

la,/erecl extant 2 layered e x l ~ t 3 mixed extent 1 mixed exWnt 32

........ - - -

u n p a t t e m e d e~terrt 1 u n p a l t e m e d s~le~t 2

yield: mixed

/

D yleld: unpatterned

mixed extent

Figure 6. (a) Extent (b) Product Yield Profiles for System 3 (0m~ -

0ave)

3. AXIAL D I S P E R S I O N The plug flow analysis demonstrates the potential advantage of running product-inhibited systems with patterns. More importantly, it validates the value of the novel layered patterns proposed in this work. One concern with this analysis is that it assumes a discontinuous 0 profile in going from one layer to the next. Since this is an unrealistic idealization, it is necessary to assess the extent to which the layered pattern's advantages are compromised by the smoothing of the 0 profile by axial heat dispersion. This is done by conducting simulations based on a mixing cell model containing a specific mechanism for the backflow of heat between cells. The model equations and information about the solution strategy are presented in an earlier paper [1].

(a)

........

ji"

(b)

Extent of Reaction Profiles

Temperature Profiles

I=

9

Z

m

Rxnl

Z

layere0

9= 9 Rxn 2 , a y ~

IElllll

Rxn :3 layered

-.r

~

~ll l a v c r e ~ l

----

~ . ~ rob,=

~ mixed

........

Rxn 1 u . ~ , . ~

4 unpattemcd

- - ,-,,,,,, - - -

Rxn 2 m i x e d Rxn 3 m i x e d

Rxn 2 unpattemer

I m

layered

~

I

u np~Ittenled

Figure 7. (a) Extent (b) Temperature Profiles for Mixing Cell Model with Figure 4 Parameters (0== - 0ave)

206 The mixing cell model is used to determine whether or not the layered pattern retains its supremacy (in cases where the plug flow model predicts it) despite axial dispersion. Applying Figure 4's parameters to the mixing cell model generates Figure 7. A comparison of Figures 4 and 7 reveals that the performance of the layered pattern is slightly compromised by the backflow of heat into the main reaction zones (D yield falls from 50.5% to 46.8% when axial dispersion is included in the analysis), but this effect cannot undo the large advantage promised by the plug flow analysis. When the plug flow analysis predicts only a slight advantage for the layered pattern over the mixed, the relatively small drop in the layered reactor's yield can be enough to tip the scales in favor of the mixed bed. The bottom line is that axial dispersion will have a slight adverse effect on the selective conversion of the layered pattern; but if the advantage predicted by the plug flow analysis is significant, it will be largely maintained. 4. CONCLUSION Spatially patterned reactors can provide conversion and selectivity enhancements over traditional unpatterned reactors for systems involving byproduct inhibition. The uniformly-distributed and layered patterns each have their own operational advantage. The well-mixed configuration benefits from more frequent switching between primary and auxiliary reactions. The layered pattern provides an added degree of flexibility by allowing the layers of different catalyst to be maintained at their own optimal temperatures. This temperature-tuning ability is particularly important when conducting highly temperature-sensitive reactions, and it is the degree of this temperature-sensitivity that is critical in determining which pattern strategy will be most advantageous for a particular system. 5. NOTATION Bi

h Ri

yi z 8i

0 Oc

Dimensionless heat of reaction i Dimensionless heat transfer coefficient between fluid and reactor wall Dimensionless rate of reaction i Dimensionless extent of reaction i Dimensionless reactor coordinate Volume fraction of catalyst for reaction i Dimensionless temperature of reacting fluid Dimensionless temperature of heat transfer media

6. R E F E R E N C E S

A. S. C6t~, W. N. Delgass, and D. Ramkrishna, Chem. Eng. Sci., submitted.

This Page Intentionally Left Blank

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science Ltd. All rights reserved.

209

A novel laboratory reactor for gas-phase transient kinetics based on time-of-flight and quadrupole mass spectroscopy H. T. Randall a, P. L. Millsb and J. S. McCracke b aFirmenich SA, CH-1283 La Plaine, Geneva, Switzerland bDuPont Company, Central Research and Development, Experimental Station, E304/A204, Wilmington, Delaware 19880-0304 USA Abstract The development and application of a novel laboratory reactor system for investigating the kinetics of gas-phase heterogeneous catalyzed reactions under transient conditions is described. The system contains a fixed-bed microreactor that allows a bulk form of the catalyst particles (dp= 100 to 350 tma) to be exposed to either pulse or step-up inputs in concentration. The transient kinetic products are monitored using either a quadrupole or time-of-flight mass spectrometer. Application of the system to investigate the oxidation and reduction of 7-VOPO4 and &VOPO 4 pure oxide phases using oxygen and n-butane is also described. The redox kinetics are interpreted using a surface kinetics model combined with a model that accounts for sub-surface diffusion of lattice oxygen to the surface active sites. Novel features and performance of the reactor system are also described.

1. INTRODUCTION The use of transient reactor operation to create a performance advantage over conventional steady-state operating modes has received increasing attention in recent years [1-4]. It is now generally well-recognized that improved reactor performance on a time-averaged basis is possible in certain instances by the use of various transient techniques, such as periodic forcing of reactor input flow parameters, flow reversal of the reactor feed and product gas, and catalyst circulation through zones having different reaction environments. Particular examples include the oxidation of SO 2 to SO3 [5], the destruction of NOx's by reduction with NH a [6], and the partial oxidation of n-butane to maleic anhydride [7]. Additional details on catalytic and reaction engineering aspects of forced unsteady-state operation are provided in a recent monograph by Silveston [8]. Development of kinetic models t h a t can be used in detailed reaction engineering models for forced unsteady-state reactor systems is an essential element of the systems analysis [9]. However, collection and interpretation of kinetic data under transient conditions requires more sophisticated experimental

210 hardware and mathematical modeling techniques than those used for steady-state kinetic studies. Particular challenges include generation of well-characterized input concentration transients, achieving an ideal flow pattern for the gas and solid catalyst, ensuring t h a t interparticle and intraparticle gradients are negligible, obtaining reproducible catalyst oxidation states, and collecting the time-resolved gas-phase reaction products. For these and other reasons, specialpurpose laboratory reactors for transient kinetics have been developed over the years, which was the subject of a recent review [10]. One of these is the TAP T M (Temporal Analysis of Products) reactor system, which has been used in DuPont and other catalyst research laboratories over the past 5 to 8 years since it was first commercialized [11]. Experience in using the TAP TM reactor within DuPont has lead to development of a new system that has improved sensitivity, a simpler vacuum chamber design, simpler yet robust valves for generation of gas pulse and step-up inputs, dual quadrupole and time-of-flight mass spectrometers for collection of single ions or the entire mass spectrum, and standardized computer software for modeling of the transient responses and extraction of the kinetic model parameters. The primary objective of this work is to summarize the design features of the new reactor system, and to briefly demonstrate its' utility by determining the oxidation and reduction kinetics of well-characterized T-VOPO4 and 5-VOPO 4 catalysts using oxygen and n-butane as the oxidizing and reducing gases, respectively. A comparison between the sensitivities of the original TAP TM reactor system and the new system is also provided.

2. NEW TRANSIENT KINETICS SYSTEM A description of the new system and a brief overview of its performance is provided in this section. Other details are available elsewhere [12] to which the interested reader is referred.

2.1 System description A schematic of the new system is shown in Figure 1. It contains a single eightinch vacuum chamber in which the fixed-bed microreactor and pulse valves are mounted in a vertical orientation in the center of the top plate. The microreactor and pulse valve manifold are located on the top of the vacuum chamber for ease of operation and maintenance. The reactor is isolated from the primary chamber using a custom-designed flange valve. The quadrupole mass spectrometer (UTI, Model 100C) is mounted on a 4 3/4-inch flange so that the ionization cage is perpendicular to the flight of the reactor product gas. To improve the sensitivity, the distance between the reactor exit and the ionization cage has been reduced in the new design to about 4 cm versus 22 cm in the original TAP TM system. A linear time-of-flight (TOF) mass spectrometer (Jordan Associates) is connected to the main vacuum chamber in a horizontal orientation using a six-inch flange. The flight tube has an overall length of ca. 36 inches that terminates with a micro channel plate (MCP) detector. The electron gun for the TOF system is mounted on a 2 3/4-inch flange directly across from the flight tube grid assembly. To monitor the pressure in main vacuum chamber and the flight tube, full-range ion

211 gauges are mounted on 2 3/4-inch flanges. Vacuum in the primary chamber is provided by a 520 liter/sec turbomolecular pump, while additional vacuum for the TOF flight tube is provided by a 210 liter/sec turbomolecular pump. With this combination of pumps, a background pressure of ca. 10 -8 torr is achieved in the flight tube, while 10 -7 torr is obtained in the primary chamber. Continuous Valve Pulse Pulse Valve A Jw Valve B Leak Valve " ~ ~ : > ~9 ' Chamber Isolation "~ ~ ~ Valve Quadrupole ~ ,, ~ n Mass Spec ~ |

Full Range Ion Gauge

Flight Tube ~ ~ /

Grid Assembly

Full Range Ion Gauge

.._

etector

210 I/s Turbo Pump

520 I/s Turbo Pump

Figure 1. Schematic of the new system for gas-phase catalyzed transient kinetics. The fixed-bed microreactor is constructed of type 316 stainless steel and has an inner diameter of 5.5 mm with an overall length of 41.7 mm. It seals against the heated valve manifold using a Kalrez@ O-ring. A custom-designed splitter valve for directing a continuous controlled leak of the reactor product gas to the vacuum chamber can also be attached to the reactor exit. The reactor is packed by sandwiching the catalyst particles between two sections of 250-425 pm quartz beads. The front inert section serves to preheat the inlet gas, while the post inert section minimizes the dead volume and broadening of the response once the product gas exits the catalytic section.

2.2 Data collection and TOF operation When the quadrupole mass spectrometer is used as the detector, the system is operated using the s t a n d a r d TAP TM reactor automation program t h a t was supplied by Autoclave Engineers. This is based upon a Hewlett-Packard Model 360 UNIX workstation. Because the TOF system was being evaluated for the first time as a new technique for collection of time-resolved transient kinetic data from a heterogeneous catalyzed reaction, a special-purpose control and data acquisition system was developed using commercially available components. Key aspects of the time-of-flight operation and data acquisition are summarized below. Referring to Figure 1 and material given by Cotter [13], the molecular beam from the fixed-bed microreactor is first ionized by an electron beam that intersects

212 the beam at right angles. The ions are then repelled by a repeller plate that is maintained at ca. 3150 volts. They are then drawn through the extraction grid, which is set at ca. 2780 volts. The ions are then accelerated through the ground grid into the TOF flight tube. During this transit, the ions pass through focusing and beam steering plates that direct the beam onto the detector grid and compensate for initial beam velocity vector. The ions are then detected with a microchannel plate (MCP) detector. A schematic of the relationship between the various timed events that must occur is illustrated in Figure 2. In this example, the output from the pulse valve driver is set to occur once every 500 ms, which also triggers the TOF pulse generator. The pulse generator initiates a pulse every 2 ms, which corresponds to the allotted time for collection of a single TOF spectra. This pulse also triggers the extraction grid pulser, which sets the extraction grid potential equal to the deflection plate potential to 3150 volts for 4 ps. During this time, the ions are created. After 4 ps, the extraction grid potential is set back to 2780 V, which causes the ions to be extracted and accelerated down the flight tube towards the detector. The falling edge of the extraction pulse is used to trigger the oscilloscope. A 2 ps delay is applied in order to avoid acquiring unnecessary data before ions have reached the detector. Acquisition then occurs for 10 ps, which can be adjusted so that the heaviest ion has sufficient time to reach the MCP detector. The signal generated by the MCP detector is acquired with a LeCroy 9374L digital oscilloscope and transferred to a PC for subsequent processing. 500 ms (valve repetition rate)

=l v -

I I I I

Valve driver out put 12 ms (TOF repetition rate)

Pulse generator out put 10 ps Ext ract ion

OV

pulse

-470 V 4 psI

Lecroy

9374 L

I I I

I I 2 ps (delay)

Figure 2. Diagram of the timing of an experimental cycle for the TOF system.

213 2.3 TOF system calibration The TOF spectrometer was typically calibrated by injecting a gas mixture containing He, Ne, N2, Ar, and CO2, Xe, and other components using the TAP TM reactor pulse valve manifold and measuring the time-of-flight of the detected ions. The resulting time-of-flight versus m/z data was fitted to a second order polynomial as proposed by Cotter [13]. The spectrum was then converted from a time-of-flight basis to one in terms of m/z. Figure 3 shows a typical TOF spectrum of a gas mixture containing He, Ne, N2, 02 and CO2 that is based upon signal averaging thirty spectra after converting the flight times into m/z values using the calibration equation. In this example, the time-of-flight for CO 2 is about 10 ~s. By comparison, the time required for the UTI quadrupole to scan over the same range, assuming 4 points per amu, would be a few seconds. The ability of the TOF to provide rapid detection of multiple ions when compared to the QMS is obvious.

N2 +

He + i,_ .Q I,.

c02 +

t

u~ c'-

I

0

5

c§ 10

ikl+ 0 +

i

15

Ne+ 20

[ 25

30

35

40

45

50

m/z (amu)

Figure 3. Time-of-flight spectrum of a mixture of He, Ne, N2, 02 and CO2 obtained after converting the flight time to a mass basis using the calibration equation. The flight time for CO 2 is about 10 ~s in this example. 2.4 O t h e r e x p e r i m e n t a l details The VOPO 4 precursor material, which is VOHPO4o(1/2)H20 , was prepared by first slurrying 15 grams of V205 into 900 m] of iso-propano] in a round bottom flask to which 38 grams of 85 % phosphoric acid was added. The resulting mixture was stirred and refluxed under nitrogen for 24 hours during which the yellow slurry had became pale blue. The cooled blue slurry was then filtered and the solid collected and washed with acetone until the washings were colorless. The pale blue was dried under suction and stored in dry nitrogen for use in preparation of the VOPO4 phases, which are described below. Preparation of the 5-VOPO 4 and T--VOPO4 phases was performed by using the above precursor and following the methods described by Bordes [14-16]. Powder X-ray diffraction of these materials gave patterns that were identical to those reported for these phases. The samples were stored in dry nitrogen until used.

214

3. RESULTS and DISCUSSION 3.1 Comparison of sensitivity A comparison between the sensitivity of the new system and the TAP TM system was performed by measuring the response of the QMS using different pulse intensities of Ne at room temperature. To provide a common basis, the same pulse valve manifold, fixed-bed microreactor, and QMS probe were used. The pulse intensifies were determined by measuring the pressure drop from a gas reservoir having a measured volume of 13.6 ml for a given number of gas pulses. This technique allowed pulse intensities as low as 10" molecules/pulse to be measured. The effect of distance between the reactor outlet and the QMS ionization cage on sensitivity was also determined for the new system by performing experiments with and without a special-purpose spacer that could be inserted into the primary eight-inch vacuum chamber. This allowed the distance between the reactor outlet and the QMS ionization cage to be reduced from approximately 3.25 inches (8.26 cm) without the spacer to ca. 1.5 inches (3.81 cm) with the spacer in place. Figure 4 compares the integrated QMS signals at different Ne pulse intensities from the TAP TM system to the new system where the latter was operated both with and without the spacer. By taking the ratios of the areas, it can be shown that there is a 100 to 200-fold increase in sensitivity for the new system without the spacer when compared to the TAP TM system. The increase in sensitivity for the new system when the gap is reduced is more than a factor of 300. Moving the ionization cage even closer to the reactor exit had a negligible effect on the measured pulse area and also resulted in some discoloration of the flange plate due to heat transfer from the hot ionization cage to the metal. Collectively, these observations suggested that the above distance of 1.5 inches was near optimal for this particular vacuum chamber design and range of pulse intensities. 100000

-

10000

9TAP-TOF with spacer ,, TAP-TOF without spacer

lOOO-

=

9TAP

100 -

10I

-

1E+14

-I

I

1E+15 Pulse

1E+16 intensity

"I

1E+17

1E+18

(molecules/pulse)

Figure 4. Comparison between pulse areas of the TAP TM and the new system with and without spacers versus pulse intensities of inert neon gas.

215

3.2 Oxidation of reduced VOPO 4 catalysts A comparison between the model-predicted and experimental t r a n s i e n t responses for reoxidation of reduced T-VOPO4 at 380 ~ and 420 ~ is shown in Figure 5. The step response of Ne gas is included to illustrate the difference between the responses for the inert and reacting gases. The model predictions are based upon a surface reoxidation rate equation that is first order in gas phase oxygen and second order in the concentration of surface lattice oxygen vacancies. rox = kC02 (1 - Oo,s)2

(1)

Here, k = kox Ctot,s is the product of the true reoxidation rate constant and the total concentration of surface sites, while 0o,s is the fractional coverage of surface lattice oxygen sites. To obtain satisfactory agreement with the oxygen response data, it was necessary to account for solid-state diffusion of surface lattice oxygen into the lattice oxygen vacancies in the subsurface that were created during the reduction cycle by n-butane. The two model parameters that emerged were the surface reoxidation rate constant k and the solid-state diffusion constant for lattice oxygen Do,ss. The latter parameter was expressed in the form of a characteristic time for solid-state diffusion t d = 82/Do,ss where 8 is the characteristic depth for bulk diffusion. Both k and t d were determined by nonlinear parameter estimation in which the experimental output responses for oxygen at each temperature were compared to the model predicted output responses. The model predictions were obtained by solving the transient form of the axial dispersion model where gas transport through the three packed zones was described and also coupled to the reaction kinetics in the central catalytic zone. The detailed modeling equations and numerical solution methodology are omitted here for brevity, since they are described elsewhere [17]. It suffices to say that the method for model discrimination is robust and no approximations are introduced. The results show that an increase in the reoxidation temperature from 380 ~ to 420 ~ reduces the characteristic time for diffusion from 148 s to 93 s. This occurs due to either a smaller effective path length, or as a result of an increase in the solid-state diffusion coefficient for lattice oxygen. The same quality of agreement between the model-predictions and experimental responses was obtained for the 8-VOPO4 catalyst, so these are omitted here for brevity. The activation energies for the surface reoxidation of the 8-VOPO4 and T-VOPO4 phases were found to agree with each another within the estimated standard error. However, those for the characteristic time for solid-state diffusion differ by more than a factor of three since Ed = 41 + 8 kJ/mol for 7-VOPO4, while Ed = 12 + 9 for 8-VOPO4. This shows that the reoxidation of the T-VOPO4 phase exhibits a greater sensitivity to temperature than the 8-VOPO4 phase.

3.3 Reduction of oxidized VOPO4 catalysts Interpretation of the transient responses for the reduction cycle of T-VOPO4 and 8-VOPO 4 using n-butane was performed using both parallel and series reaction

216

networks. According to the first scheme, n-butane is assumed to react with surface lattice oxygen (O)s to form maleic anhydride, CO2, CO and H20 as products via three parallel reactions. Subsurface lattice oxygen (O)ss can diffuse to the surface to replenish surface lattice vacancies ( )s created by reduction of the surface lattice oxygen (O)s at some characteristic time constant t d. The second scheme is similar, except that maleic anhydride can undergo further reaction with surface lattice oxygen to form CO2. The rate of surface reduction for these species was assumed to be first-order with respect to gas-phase n-butane or maleic anhydride and fractional-order with respect to the concentration of surface lattice oxygen. In the rate equations given below, r 1 denotes the main reaction where maleic anhydride is formed from n-butane, whereas r 2 and r a denote the secondary reactions where maleic and n-butane undergo combustion, respectively: rl = kl CB eo~,~s ;

r2 = k2 CMAN eo~,2s ;

r3 = k3 CB eo~,3s

(2)

It was assumed that the reaction order 7i for each species can be unique, i.e., 71 v 72 7a. Attempts to fit the experimental transient responses using the above rate form and others gave poor agreement with the model predictions. Inclusion of solid-state diffusion of the subsurface oxygen as described above was necessary to obtain satisfactory agreement. For the initial condition, it was assumed that the catalyst was completely oxidized prior to reduction.

1.2 0 c

o Q.

"0

.~

Ne

1.0 0.8

- - Data, T=380oC

0.6

- - Data, T=420~ --" Model 5

m

E 0.4 o

z

0.2

T!~ c)

t, (s)

380

148 _+ 1.8 93 + 0 . 4

420

0.0 0

100

k <-102 (m3/kg's) .....

Time, s

1.28 +_0.002 3.92 +_ 0.07 200

300

Figure 5. Comparison between the model-predicted responses and the experimental responses for a step-up input of oxygen over a reduced ~-VOPO4 catalyst at different temperatures. Conditions: P = 1.01-105 Pa, Y%o = 0.02, Q = 25 ml/min. The catalyst was previously reduced 15 min with 1% n-butane in N2 at 420 ~

217 The model-predicted and experimental responses for n-butane at 380, 400, and 420 ~ for the parallel pathway model using the surface reduction model coupled with subsurface lattice oxygen diffusion are compared in Figure 6. The agreement between the d a t a and model predictions is excellent. From a qualitative perspective, the shapes of the responses suggest that the reduction is occurring on two distinct time scales. The first time scale is on the order of seconds and corresponds to a fairly rapid reduction of the catalyst surface. This leads to a rapid increase in the production of CO2 and maleic anhydride (not shown). The second time scale is on the order of tens or even hundreds of seconds, and corresponds to the consumption of subsurface oxygen as it diffuses to fill the surface vacancies created by the various catalyst-reducing reactions. This behavior agrees with the expected one, since the solid-state diffusion has a time constant that is on the order of 100 seconds.

Time, s Figure 6. Comparison between the calculated and experimental responses of butane to a concentration step of butane over a reduced 5-VOPO4 catalyst at different temperatures. Conditions: P = 1.01.105 Pa, Yc4,o = 0.01, Q = 11.3 ml/min. The catalyst was previously reoxidized for 15 min with 2% 02 in N2 at 420 ~

4. SUMMARY and CONCLUSIONS The reaction kinetics of the oxidation and reduction of pure phase 5-VOP04 and T-VOPO4 metal oxides have been investigated using a new experimental reactor system. This system is primarily designed for investigation of the transient kinetics of gas-phase reactions, and can use either a quadrupole or time-of-flight mass spectrometer for detection of the reaction products. Improved sensitivity,

218 simpler vacuum system design, and simultaneous detection of multiple masses are key advances that have been demonstrated. Kinetic models that were solely based on redox processes on the surface layer were not able to provide an adequate representation of the experimental transient responses. The coupling of subsurface lattice oxygen transport via solid-state diffusion with a surface kinetic model that was first-order in oxygen, butane, and maleic anhydride, fractional-order in oxidized sites, and second order in reduced sites was shown to provide satisfactory agreement between the model-predictions and experimental data. The model for solid-state diffusion is a one-dimensional representation, but more complex models that consider the geometry of the solidphase crystallites and other details associated with the inorganic chemistry may be required for more complex mixed-phase oxide catalysts. A correlation between the reaction rate parameters for the two different catalyst phases was not attempted, since data for a well-characterized mixed-phase material would be needed as a basis for comparison. This could be a topic for future work. 5. R E F E R E N C E S

1. Y. S. Matros (ed.), Unsteady-State Processes in Catalysis, VSP Press, Utrecht, The Netherlands, 1990. 2. Y. S. Matros, Can. J. Chem. Eng., 74 (1996) 566. 3. C. W. Robinson (ed.), Proc. of the 2nd International Conference on UnsteadyState Processes in Catalysis, Can. J. Chem. Eng., 74 (1996) 563. 4. Y. S. Matros (ed.), Proc. of the 3rd International Conference on Unsteady-State Processes in Catalysis, (to appear), VSP, Utrecht, The Netherlands, 1998. 5. Y.S. Matros, in Preprints of the Int. Conf. Sulfur 87, British Sulfur Co. Ltd., Gosport, Hants, England, 1987. 6. Y. S. Matros and G. A. Bunimovich, Ind. Eng. Chem. Res., 34 (1995) 1630. 7. R. M. Contractor, D. I. Garnett, H. S. Horowitz, H. E. Bergna, G. S. Patience, J. T. Schwartz, and G. M. Sisler, in Studies Surface Science and Catalysis (V. C. Corberan and S. V. Bellon, eds.), Vol. 82, Elsevier, Amsterdam, The Netherlands, 1994, C.3-1. 8. P. L. Silveston, Composition Modulation in Chemical Reactors, Cambridge University Press, Cambridge, England, 1997. 9. M. P. Dudukovic, F. A. Larachi and P. L. Mills, Chem. Eng. Sci. (1999), to appear. 10.P.L. Mills and J. J. Lerou, Revs. Chem. Eng. 9(1/2) (1993) 1. l l . J . T . Gleaves, J. R. Ebner and T. C. Kuechler, Catal. Rev.-Sci. Eng., 30 (1988) 49. 12.H.T. Randall, P. L. Mills and J. S. McCracken, presentation at the 3rd World Congress on Oxidation Catalysis, San Diego, CA, September 21-26, 1997. 1 3 . R . J . Cotter, ACS Symp. Series, Vol. 549, American Chemical Society, Washington, D. C., 1994. 14.E. Bordes and P. J. Courtine, Chem. Soc., Chem. Commun., (1985) 294. 15.E. Bordes, Catalysis Today, 1 (1987) 499. 16.E. Bordes, Catalysis Today, 16 (1993) 27. 17.P.L. Mills and H. T. Randall, Multiregion distributed parameter dynamic model of a fixed-bed reactor, SIAM Annual Meeting, Toronto, July 13, 1998.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

DEVELOPMENT

OF

NEW

PHOTO-CATALYTIC

219

METHODS

AND

REACTORS FOR WASTE WATER TREATMENT A. Starosud a, A. Bhargava b, C.H. Langford a, and A. Kantzas b*. University of Calgary, Calgary, Alberta, Canada T2N 1N4 a: Department of Chemistry b: Department of Chemical and Petroleum Engineering

Abstract The efficiency of photocatalytic reactors using immobilized photocatalysts have been found to be lower than those using dispersed titanium compound particles (slurry). However, for practical applications, immobilized catalyst is preferred as it does not need an additional separation process to recover the photocatalyst. Zeolites, as catalyst supports, offer the advantage of having high absorptivity for organics in wastewater aiding at preconcentrating pollutants on their surface. More than 40 different types of zeolites and their modifications as a support were synthesized and tested. Different methods of loading TiO2 on zeolites and special techniques to characterize distribution of the photocatalyst were developed. Adsorbency of various organic materials commonly present in wastewater was also evaluated. Lab-scale experiments on the newly developed immobilized photocatalyst particles (100-900 microns) have already been conducted. The design and development of the appropriate photocatalytic reactor requires a study of the hydrodynamics of the reactor coupled with the intrinsic rate kinetics to achieve higher quantum yields and optimum photocatalyst requirements. An annular fluidized bed photocatalytic reactor operating in absorption and regeneration modes has been constructed for the purpose of kinetic studies. The reactor has been modelled and simulation results are presented in this study. Kinetic studies are being carried out to obtain reliable intrinsic rate data using the newly developed photocatalyst. The data so obtained will be used to verify the reactor model to facilitate its scaleup. 1. INTRODUCTION Heterogeneous photocatalysis offers an attractive alternative for wastewater treatment especially when treating low concentrated, high volume fluids. For the photocatalytic process, the most suitable semiconductor is TiO2. It is characterized by chemical inertness, nonphotocorrosivity and nontoxic influence on microorganisms. These properties, coupled with its ability to create highly reactive oxidant (hydroxyl radical) on excitation with UV radiation make titanium compounds highly suitable for application in wastewater treatment. The hydroxyl radical (OHe) which is known to be one of the most powerful oxidizing species, mineralizes organic pollutants to carbon dioxide and inorganic ions. This technology has been developed and it is currently being evaluated for application at pilot plant and semi-industrial scales of testing. The efficiency of photocatalytic reactors using immobilized photocatalysts have been found to be lower than those using dispersed titania particles (slurry). Application of fine powder of TiO2 is technologically impracticable because powders could be easily washed out. To prevent this phenomenon, it is necessary to construct additional equipment [1 ] which causes a sharp rise in the cost of the process.

220 There is a rigid practical necessity to immobilize titania [1,2]. Many attempts have been made using glass beads, fiber glass, silicon, quartz, activated carbon and zeolites as support [ 1,2,3]. 2. PHOTOCATALYST SYNTHESIS Zeolites, as a catalyst support, offer the advantage of having high adsorption for organics in wastewater aiding preconcentrating of pollutants on their surface. More than 40 different types of zeolites and their modifications were synthesized and tested as supports. All of the synthesized materials can be divided in four classes: 1. Mesoporous molecular sieves (MCM-41) and relatives. 2. Titanium silicates (TS class with Ti incorporated into zeolite lattice) of both micro and meso pore types. 3. Aluminophosphates sieves (VPI-5 and relatives) containing both Si and Si-Ti. 4. Siliceous zeolites of the MFI and MEL structure families (TS-16, ZSM 48, etc.). 3. PHOTOCATALYST EVALUATION Adsorption characteristics of all synthesized materials were evaluated using static conditions. The static experiments involved slurring the material into solutions of organic contaminants for periods up to 48 hours. The aliquot and filtrate were analyzed by HPLC or GC. Most of these synthesized zeolites exhibited relatively low adsorption toward chosen contaminants as compared to activated carbon as a standard. Some results for Isopronanol (IPA) acetone and 2,4 - dichlorphenol are given in Figure 1 for the most promising materials. The major rival for zeolite materials is activated carbon. From Figure 1, it can be seen clearly that two of our materials (163ARC and TS-16) have nearly the same capacity as carbon and are even superior toward polar contaminants (acetone, IPA).

2,4-DCP

.

. .

u

~

Mtteritla

Figure 1. Static Adsorption of Organics on Zeolite Materials The most interesting zeolites were loaded with titania to obtain an integrated photocatalystadsorbent system (IPCA). TiO2 loading was accomplished by stirring the zeolite with a previously prepared colloid of TiO2, evaporating the solvent, and calcining the resultant solid at a temperature between 200 ~ and 500 ~ This was the preferred method of TiO2 loading which has been identified in earlier

221 work [4]. Two modifications were also used. First, the standard colloid with particle size in 10 nm and higher range were replaced with a preparation of"quantum size" or "Q-sized" particles of 2-4 nm size [5]. A further modification with loading Q-sized particles under these conditions, in this case filtration, replaced evaporation of solvent. Physical characterization of IPCA has included Raman spectroscopy to examine the state of TiO2, SEM to identify TiO2 clustering, and XRD to verify the identity of the zeolite phases and the apparent particle size in TiO2 loading. An interesting result is that all TiO2 reflections in XRD have line width suggesting particle size smaller than the size of colloidal TiO2 particles used in loading onto the support. This led us to question the TiO2 distribution. A novel energy filtered transmission electron microscopy technique was developed to obtain a Ti distribution map [6]. It was found that whenever zeolite pores are smaller than the particle size of T i Q , the colloidal particles used in loading titania remains on the surface. For general screening of photocatalytic behavior of IPCA, a larger diameter Teflon cell with a quartz window at the top was used. Several of these could be irradiated simultaneously under a pair of 40W fluorescent UV lamps with peak output at 350 nm. Substrate oxidation was measured using a mass balance approach where reduction of solution concentration was monitored and IPCAs were extracted with organic solvent at the end of runs to recover any material adsorbed. The overall results of screening for photocatalytic activity can be summarized as follows" 1. a "classical" MCM-41 exhibited higher activities than zeolite Y (used as a reference); 2. titanium silicates showed very low activity; 3. silicon and titanium containing aluminophosphates had low activity; 4. several members of MFI and MEL families were quite active. Unfortunately, the adsorption characteristics of MCM-41 class are not satisfactory, so emphasis turns to MFI and MEL family of materials. Several locally synthesized systems including ZSM48, ZSM-11 and TS-16 are very promising. As well, commercial ZSM-5 and silicalite I are competitive. In order to reduce the cost of the IPCA, a new approach has recently been developed. The general idea was to use inexpensive and readily available particles with a desired diameter as a carrier for titania and zeolite. The carrier has to be mechanically strong, should not reduce photoactivity and have a "golf ball" like surface to reduce attrition of zeolite/titania by depressions. 4. P H O T O C A T A L Y T I C R E A C T O R

4.1 Reactor Design and Construction For heterogeneous photocatalytic reactions the contact among reactants, photons and catalysts must be maximized. Mixing and flow characteristics of the photoreactor may greatly enhance these contacts. If a fixed bed reactor is used, the irradiated aliquot of catalyst is limited to a thin layer and a large reactor volume is required. For the liquid-solid and gas-solid systems, continuously stirred tank photo reactors and fluidized bed photo reactors, respectively, are the most suitable ones for enhancing contact efficiency even if their operation is quite expensive and troublesome. The rate of photocatalytic reaction is greatly affected by flow rates. The rate enhancement is not due to elimination of mass transport resistances, as expected in classical catalytic systems, since such considerations do not apply for most heterogeneous photo-processes that are characterized by low reaction rate with respect to mass transport rate. The enhancement is determined by the fact that on increasing flow rates, the frequency of exposure of the catalyst

222 particles to irradiation increases. The catalyst particles continuously receive diffuse radiation of reduced intensity due to absorption by other catalyst particles. They are directly irradiated intermittently due to shielding effect of particles which randomly intercept direct irradiation. By increasing flow rates, the frequency with which the catalyst particles may be directly irradiated increases and eventually, the reaction rate is enhanced. With the above considerations, an annular liquid fluidized bed photocatalytic batch reactor with full recirculation was constructed. A schematic of the reactor showing the adsorption and regeneration cycles is depicted in Figure 2. The above reactor was then modelled based on the study done by Rideh et. al. [7]. In this kind of reactor, the extended light source is placed at the axis of a reactor composed of two coaxial cylindrical tubes. The emitted radiant power is absorbed by the reaction system contained in the annular reactor volume. Irradiance diminishes in a filled reactor with increasing radius. This geometry called the negative geometry of irradiation makes the most efficient use of the light emitted by an extended light source. In fact, this geometry is used in all immersion type photochemical reactors, and most industrial photochemical production units [ 10].

Adsorption cycle Regeneration cycle

\7

A

U

N

N

Figure 2. Schematic diagram of the photocatalytic reactor

4.2 Photocatalytic Reactor Modelling The purpose of this study was to model and simulate a liquid fluidized bed photocatalytic reactor for 2-Chlorophenol degradation operating in once through and dual mode (adsorption followed by regeneration). Also, to analyze the effect of various operational parameters such as initial pollutant concentration, flow rate, oxygen partial pressure, absorbed UV light intensity, initial bed height using the developed simulator with a view to find out the range of operational parameters for optimum reactor performance.

4.2.1 Hydrodynamic Modelling A hydrodynamic model was set up using the Wen and Yu Correlation [8] to predict the bed voidage as a function of superficial velocity. The number of particles in the elemental volume under consideration were calculated using geometry. We could arrive at the average light

223 intensity reaching a photocatalyst particle surface using the line source with parallel plane (LSPP) emission model [9]. In this model, the lamp is considered to be a linear source in which each point emits radiation in parallel planes perpendicular to the lamp axis. Combining this hypothesis with the Beer-Lambert equation, the integration of the differential equation describing the radiant power or irradiance as a function of the optical path yields the profile of radiant power or irradiance.

4.2.2 Kinetic Modelling Kinetic modeling of the primary degradation steps is essential for any practical application of the process. Rideh et al. [7] showed that 2 - Chlorophenol was degraded in an illuminated suspension of TiO2 according to the following stoichiometry: 2C6H4OHCl + 1302

TiO2/ho

> 2HCI+ 12CO 2 + 4H20

4.2.3 Model Assumptions A pseudo-plug flow model has been used to simulate the process. A number of discrete plug flow reactors of an annular shape have been considered to represent the reactor. The assumptions used in the development of model are: 1. Complete radial mixing and no axial mixing in the elemental volume. 2. No mixing between the adjacent annular PFR's. 3. No reaction in the part of the reactor above the expanded bed (i.e., no direct photolysis). 4. Even distribution of photocatalyst particles in the expanded bed region. 5. Same inlet concentration to all the annular PFR's. 6. Complete mixing of the wastewater while effluent is recirculated. 4.2.4 Model Development 4.2.4.1 Prediction of Absorbed Light Intensity Total number of particles in the reactor can be calculated from the voidage and height of bed at minimum fluidization : Nj, - 4

r3

Particle number density in the expanded bed region can be calculated from the total number of particles in the reactor, assuming equal distribution of particles in the expanded bed region. N1, Nj'D

-" 7~ ( R ~

_ R21)I__IEx

Where, the expanded bed height is calculated from the voidage predicted by Wen & Yu Correlation:

HEx = ~( -l -~e":f) ) H.,f Thus, number of particles in the elemental volume can be calculated as : FI I, --

NpDTt(g[, - g~ ) ~

The light intensity incident to elemental volume under consideration is given by the LSPP Emission Model :

224

1"c :

$I~,~ exp[-l.t(r - R, )] ~(R,~ - R 2 )

Hence, the net light intensity absorbed by the photocatalyst particles present in the elemental volume is given by : ertr 2 2rtr~~[

I h, c n lAh,, ' =

5

where, c~ - fraction of particles which are irradiated by the incident photons. [3 - fraction of incident light intensity absorbed. 4.2.4.2 Material Balance for Once Through Process Material balance based on the amount of pollutant (2-Chlorophenol) inside the elemental volume can be written at steady state as follows : [Mass of 2-CP in] - [Mass of 2-CP out] - [Mass of 2-CP degraded by photocatalysis] =

[Accumulation]

At steady-state, [Accumulation] = 0. Also, [Mass transfer of 2-CP from bulk to photocatalyst surface] = [Mass of 2-CP degraded by photocatalysis]

Q.(Ci )r,z - Q'(Ci)~,~+~ = A . A z . ( r a t e ) dC i

(rate)

dz

UL dC i

I (rate) -

1

U J"dz

where, the rate expression is evaluated in the same manner as Rideh et al.[7]. 4.2.4.3 Material Balance for Dual Mode Process A similar material balance conducted for 2-CP in the bulk liquid and solid catalyst phase results in a set of ordinary differential equations.

For bulk liquid phase :

For solid catalyst phase :

dCi,

1

dz

u

dC s A F, ' ,, dz - - I W [KLa(Cs - C,~) + k'rk Iah.,.C;,2Cs,

IkLa(Cs, - C L)

The liquid-solid mass transfer coefficient was evaluated using the correlation proposed by Hassanien et al. [ 11 ]: Sh = 0.33(Ga. Mv.Sc)I/3[l + 0.22 Mv -~ (U~ / UL) 0"77] 4.2.4.4 Solution of Model Equations The material balance conducted for the pollutant in the liquid and solid phase in a PFR over the elemental volume at a fixed radial distance resulted in a set of simultaneous first-order ordinary differential equations. The equations were solved numerically using fourth order Runge-Kutta method with adaptive step-size control to obtain 2-CP concentrations at various axial distances. The above steps were carried out for all the annular PFRs. The values of 2-CP concentrations

225 obtained at the outlet of the reactor were then averaged and formed a new initial concentration to the inlet of the reactor for the next time step. The same procedure was repeated for a number of circulations to obtain 2-CP concentrations at the outlet of the reactor corresponding to time. For the adsorption cycle, break-through of the photocatalyst bed was considered when more than 95 % of it was saturated. For the regeneration cycle, bed was considered as regenerated when 99 % of the initial adsorbed concentration was removed from the solid phase due to reaction or transfer to the bulk liquid phase. 5. R E S U L T S AND DISCUSSION The dimensions and process variables required as an input for the simulations are presented in Table 1. Table 1

Dimensions and process variables Parameter Mean particle diameter Particle density Sphericity of particles Viscosity of water @ 25~ Density of water @ 25~ Molar absorption coefficient Reactor length Inner radius of the reactor Outer radius of the reactor Initial bed height Voidage at min. fluidization Oxygen partial pressure Recirculation flow rate (regeneration cycle) Superficial velocity Residence time Power emitted by source ~x (Fraction of particles incident to light) [~ (Fraction of incident light absorbed)

Value 500 microns 3000 kg/m 3 0.67 0.001002 kg.m/s 998.2 kg/m 3 0.01 0.80 m 0.06 m 0.10m 0.45 m 0.49 0.575 atm 27 lpm 0.0224 m/s 0.595 min. 1.98 x 10 .5 Einstein/m.sec. 0.50 0.75

The model for once through process was run using the above process variables and dimensions. The effect of initial concentration on 2-CP degradation was studied by maintaining the water recirculation rate at 27 lpm (when bed is fully expanded) and the initial 2-CP concentration was varied from 0.0003 to 0.004 mol/L. The effect of oxygen partial pressure on 2-CP degradation was studied by maintaining the water recirculation rate at 27 lpm and an initial 2-CP concentration of 0.0005 mol/L. The results of simulations for the once through process are presented in graphical form in Figures 3 and 4. The trends obtained with respect to the effect of initial concentration and oxygen partial pressure on 2-CP degradation satisfactorily match those obtained by Rideh et al. [7]. However, the values obtained do not match perfectly due to the following reasons : (a) our reactor was much larger than theirs (b) they used the power emitted

226 by the radiation source measured by uranyloxalate actinometry whereas a LSPP model was used in our study to arrive at the light intensity in a particular volume element under consideration (c) difference in the photocatalyst properties such as particle size, particle density, etc. 1

i

o.5 0.8

ox

O7

07

~o5 ~ o~

~.~ 115

s

o4

""

] o'

"1"

O3

03

02

O2 01

Of

0

0 20

0

30

40

50

60

70

80

: 0

10

20

3o

40

0 (XX)5

0 (XX)7

_ _

so

6o

ro

Tlme (mln)

T i m e (rain)

O.(X)l

....

O.(X)2

1).003

0 (XM n~d/L

Figure 3. Effect of initial concentration on 2-CP degradation

,i,,,,,-o

......

0 o575

0 2875

0.46

0 575

....

Figure 4. Effect of oxygen partial pressure on 2-CP degradation

For studying the effect of operational parameters in dual mode process, the range of values of various operational parameters used in simulations are presented in Table 2. Table 2

Range of values of operational parameters

Parameters Water recirculation rate Oxygen partial pressure Power emitted by source Initial bed height

Values 9 - 42 L/min. 0.0575 - 1.0 atm (3.96 x 10 -6 ) - (9.9 x 10 .5 ) Einstein/m.sec. 0.1 -0.65 m

4(X)

350

....... \1 ..... ,\ \

~2

21X)

1i0 IIX}

\

\ ",, ""-..

-o,.,

5O

1o

15

20

25 R~r,*,

30

3.~

40

0

[)2

()~

I

04

5

06

07

().x

09

(1~,)

F i g u r e 5. E f f e c t o f f l o w rate on regeneration time

Figure 6. Effect o f oxygen partial pressure on regeneration time

The effect of varying water recirculation rate on regeneration time is depicted graphically in Figure 5. There is a pronounced drop in regeneration time required when the flow rate is increased to a value where the bed is fully fluidized, this is because of the fact that on increasing flow rates, the frequency of exposure of particles to irradiation increases. Also, the height of bed exposed to irradiation increases and the decrease in particle density facilitates more effective

227 capture of the incident radiation. The decrease in regeneration time is not so pronounced once the bed is fully expanded at a flow rate of 27 lpm, as the only beneficiary factor is the increase in frequency of exposure. According to Figure 6, there is a sharp nonlinear decrease in the regeneration time on increasing the oxygen partial pressure. It confirms the fact that the partial pressure of oxygen is a crucial factor in the photocatalytic reaction and the limitation of the rate of photocatalytic degradation is attributed by most researchers to the recombination of photogenerated electronhole pairs. Since, oxygen adsorbed on titanium dioxide surface prevents the recombination process by trapping electrons, it can be inferred that the reaction rate is a function of the fraction of adsorption sites occupied by oxygen. Hence, oxygen adsorption becomes a governing factor at very low dissolved oxygen concentrations. i

~r

i*i i

t

i

,

I

p

L

,

i

i

I---i

i

t2o

! /

L .... "

'

i

.

i

9 d~

I

/

~ .....i

15o

(i

o (~1(io i

o

(RI(X)2

00~RX)'~ 000IX)4

0 (H)(X)5

L i g h l hllenz,,iCy

0 (XI(X)6

0 00~1(17

0 (XI(R)Ir

0 00009

0

(Ein.~ldn/in.,)

Figure 7. Effect of light intensity emitted by source on regeneration time

Figure 8. Effect of initial bed height on regeneration time and bed exhaust time

Results reported in Figure 7 show an exponential decrease in the regeneration time required with an increase in the light intensity. The slope decreases after a certain value when it approaches saturation of the catalyst by the incident photons. Since, the annular irradiated thickness is quite large in the reactor, it requires a very high intensity lamp to reach saturation. A trade-off depending on the lamp availability and its power requirement has to be investigated to optimize reactor performance. Simulations were carried out to study the effect of initial bed height on regeneration time and bed exhaust time, keeping the flow rates in both the modes constant at 27 lpm and a pollutant concentration of 0.0005 mol/L in the adsorption phase. The results shown in Figure 8 depict that there is a linear dependence of bed exhaust time with initial bed height and the slope will depend upon the flow rate and pollutant concentration in the adsorption cycle. The regeneration time is around 80 min. and remains about the same for initial bed heights of up to 0.4 m and then increases linearly for higher initial bed heights. The water recirculation rate of 27 lpm suffices to raise the initial bed height of 0.45 m to full expansion. The bed exhaust times are more than ten times higher than regeneration times: so, there is a possibility of increasing the flow rates in adsorption by a factor of ten. This can serve to equalize the bed exhaust and bed regeneration times which will facilitate continuous operation of two units in parallel, one in adsorption and the other one in regeneration phase, switching over from one to another.

228 6. C O N C L U S I O N Computer simulations on a novel photocatalytic reactor configuration reveal the possibility of using such a system for purification of water containing organic contaminants. 7. N O M E N C L A T U R E 2-CP E ml

2-Chlorophenol voidage at minimum fluidization molar absorption coefficient

I" O~

A

Ci

CL Cs Ga HEX Hmf labs Iinc

radial distance, m fraction of particles irradiated by incident photons annular cross-sectional area, m 2 fraction of incident light intensity absorbed pollutant concentration, mol/L pollutant concentration in bulk liquid phase, mol/L pollutant concentration on solid surface Galileo number expanded bed height, m bed height at minimum fluidization, m absorbed light intensity, Einstein/L.s incident light intensity, Einstein/L.s

kL K 2-cP LSPP Mv Np np NpD Q Ri Ro rp Sc Sh SL; x

U6 UL

mass transfer coefficient adsorption constant for 2-CP Line Source with Parallel Plane density number total number of particles in the reactor number of particles in elemental volume particle number density, particles/L flow rate, L/s inner radius of an annular photochemical reactor, m outer radius of an annular photochemical reactor, m mean particle radius, m Schmidt number Sherwood number photon rate of the light source per unit length, Einstein/rr superficial velocity of gas, m/s superficial velocity of liquid, m/s

8. R E F E R E N C E S 1. A. Haarstrick, O.M. Kut, E. Heinzel, Environ. Sci. Technol., 30 (1996), 817. 2. I.R. Bellobono, A. Carrara, B. Barni, A. Gazzotti, J. Photobiol. A: Chem., 84 (1994), 83. 3. A.P. Davis, in "Process Engineering for Pollution control and waste minimization", N.Y., 1994. 4. Y. Xu, Cooper H. Langford, J. Phys. Chem., 99 (1995), 11501. 5. W. Choi, A. Termin, M.R. Hoffman, J. Phys. Chem., 98 (1994), 13669. 6. A. Starosud, D.P. Bazett-Jones, Cooper H. Langford, Chem. Com. (1997), 443. 7. L. Rideh, A. Wehrer, D. Ronze, and A. Zoulalian, Ind. Eng. Chem. Res. 36 (1997), 4712. 8. C.Y. Wen, and Y.H.Yu, Fluid Particle Technology, Che. Eng. Prog. Symp. Ser., 52 (1966), 100. 9. P.R. Harris, and J.S. Dranoff, AIChE J., 11 (1965), 497. 10. A.M. Braun, L. Jakob, E. Oliveros, and C.A.O. Nascimento, Advances in Photochemistry, 18 (1993),235. 11. S.M. Hassanien. H. Delmas, and J.P. Riba, Entropie, 119 (1984),17.

Reaction Kinetics and the Developmentof Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

229

An i n t e g r a t e d d e h y d r o g e n a t i o n - h y d r o g e n a t i o n m e m b r a n e reactor for the s i m u l t a n e o u s production of styrene and c y c l o h e x a n e T.M. Moustafa-, I. Ashour ~ and S.S. Elnashaie c a Chemical Engineering Department, Faculty of Engineering, Cairo University, Cairo, Egypt b Chemical Engineering Department, Faculty of Engineering, Elmenya University, Egypt c Chemical Engineering and Pilot Plant, National Research Center, Dokki, Cairo, Egypt Abstract A rigorous heterogeneous model is used to study the performance of the membrane catalytic reactor for the dehydrogenation of ethylbenzene to styrene. The mathematical model is extended to simulate a novel hybrid membrane reactor. This reactor is composed of two catalytic sections separated by selective composite membrane for hydrogen separation. One side of the reactor is a dehydrogenation section in which ethylbenzene is dehydrogenated to styrene, while the other catalytic side is a hydrogenation section in which benzene is catalytically converted to cyclohexane. Hydrogen gas is transferred from the dehydrogenation side to the hydrogenation side through the selective composite membrane. The continuous removal of hydrogen from the dehydrogenation section leads to the shift of equilibrium conversion in this section, thus higher styrene yield is obtained. Detailed paprametric investigation has been carried out for the membrane reactor and the hybrid reactor configurations. The effect of co-current and counter-current flow pattern is investigated. Superior performance in terms of ethylbenzene conversion enhancement far above the equilibrium value was observed in the hybrid reactor configuration. The selectivity and styrene yield highly exceed the industrial figures. 1. I N T R O D U C T I O N The use of palladium membranes in laboratory scale catalytic reactors to break the thermodynamic barrier of reversible reactions has been tried successfully by many investigators for a number of important catalytic dehydrogenation reactions [1-4]. However the industrial implementation of this technology requires cheaper and more durable membranes. Recent advances in the manufacture of composite membranes represent an important step forward in the direction of the industrial exploitation of the membrane reactor technology [5-7]. The industrial production of styrene (ST) from ethylbenzene (EB), over iron oxide catalysts and with steam as diluent is a typical industrial example of such reactions where a membrane reactor configuration can improve conversion considerably well above the equilibrium conversion. Most research work suggested the hydrogen permeated to the other side of the membrane to be swept by any inert gas or vacuum [8]. Reaction couphng between the catalyst bed and the sweep gas sides was proposed by a number of investigators [9,10]. If hydrogen permeated from the catalytic dehydrogenation side is utilized in a hydrogenation reaction on

230 the other side of the membrane, a new more efficient configuration can thus be introduced which we will call the Hybrid Reactor (HR). The proposed hybrid configuration in the present work allows the utilization of the permeated hydrogen from the dehydrogenation side to produce cyclohexane through the catalytic hydrogenation of benzene (BZ) on the other membrane side. Any other hydrogenation reaction can be used depending on the industrial situation.Figure (1) gives a schematic diagram of the proposed hybrid reactor. 2. D E V E L O P M E N T O F T H E M O D E L E Q U A T I O N S The following assumptions are used in the derivation of the model: 1. The reaction mixture is an ideal gas in both catalytic reactor sections. 2. Both sections of the reactor are operated at steady state conditions. 3. Radial and axial diffusions are neglected in the two catalytic beds. 4. Concentration profiles are symmetrical around the catalyst pellet center and pellet is isothermal due to high catalyst thermal conductivity. 5. For the catalyst pellets in the dehydrogenation section (DS): The diffusion model used is based on the rigorous Stefan-Maxwell equations (dusty gas model) [ 11-12] with negligible external resistances. 6. For the catalyst pellets in the hydrogenation section: The diffusion inside the pellet is described by Fickian diffusion with negligible mass transfer resistance. 2:1 R a t e E x p r e s s i o n s

i. Catalytic d e h y d r o g e n a t i o n o f e t h y l b e n z e n e Most literature suggested six independent reactions [ 13]: m a i n reaction: C6H6CH2CH 3 <=>C6H5CHCH 2 + H 2 s i d e reactions: (:(;H5CH2CH 3 ~ C6H 6 +C2H 4, C6H5CH2CH 3 ~ C6H5CH 3 + CH 4 2 H20 + C2H 4 ~ 2 CO + 4 H 2 , H20 + CH 4 ~ CO + 3 H 2, H20 + CO ~ CO 2 + H 2 The rate equations are given as follows [13]: R1 = kl(PEB - Ps'r P u z / KEB) , R e = ke pEB, R.~ = k:~ pEB pile, R4 = k4 pHeo (pETH) 0"5, R5 = k5 pH20 pMET a n d R6 = k6 ( P r / T'9 pH20 p c o (1-6) The equilibrium constant of the main reaction is given by: KEB -- exp(-AFo/RT) an(! AF o = a + b T + c T 2 (7,8) where a = 122725.16 ; b = -126.27 and c = -2.19x10 s The rates of reaction have the units of kmol/h.kg Cat., and the kinetic constants of the reactions are expressed by: k i = e x p ( A i - E i / R T ) (9) The intrinsic rate constants of these reactions are given in Table 1 by [13]. The catalyst for this EB dehydrogenation reaction is iron oxide (Fe20 3) promoted with potassium carbonate (K2CO 3) and chromium oxide (Cr203). ii. C a t a l y t i c h y d r o g e n a t i o n of b e n z e n e On the other side, the catalytic hydrogenation reaction taking place is: C6H 6 + 3H 2 ~ C6H 12 The catalyst is supported nickel and the kinetic

Table 1 F r e q u a n c y factors and Activation energy Reaction Ai Ei 1 0.851 90891 2 14.00 207989 3 0.5(; 91515 4 ().12 103996 5 -3.21 65723 6 21.24 73628

231

rate equation is given by[14]:

kKBPBZ PH2 -- RRBZ = (1 + KBPBz ) (PH, + PBZ )

Where the constants in the rate equation are expressed by the following equations [ 14]: k=k o exp(-E/RT) and KB=KBoexp(EB/RT ) (11,12) Constants for the above equations are given in Table 2 [14].

(10)

Table 2 Constants of benzene kinetic equation

Constant ko E KBo ER

Value 4.36x105 12000 7.88x10 ~ 6000

2.2 M o d e l e q u a t i o n s

For the six reactions in the DS the six material balance equations are given by: dZi/dL=qip~AnRi/FFEB and dXj/dL=qjptr4eRj/FFH2o (13,14) where i represents the reactions 1,2 and 3 and Xi is the fractional conversion of EB in each of these reactions; while j is for reactions 4,5 and 6 and Xj is the fractional conversion of the steam in each of these reactions. The differential mass balance of the hydrogenation reaction is: dXBz2/ dL=rip B:A Be(R') / FFBz 2 (15) The energy balance equations for the two sections are as follows: 1o

dT

6

~" F~Cp - - = ~ (-AH , )rljRjpBA , +U(rcND)(T- T') i=! dL J=i

(16)

i

3 dT' ~'F~Cp ~ + ,-1 ' dL

dOp , , - Cp,, (T - T ) = ( - A H ' ) R qp~ As, - U ( z N D ) ( T - T') dL

(17)

The last term in equations (16) and (17) represent the heat exchange between the reactor sections. The Ergun equation is used for the pressure drop in both reactor sections: r

dP dL

-

10- 5 (1- ~)Go /150pc~(1- ~) [ dp t;3p(;g~ dp + 1.75Go.

(18)

L

The permeation rate of hydrogen through the composite palladium-ceramic membrane is given by [5,6] "

------~--:P dL

". - p ~ . )

(19)

2.3 C a t a l y s t p e l l e t m o d e l The dusty gas model equations for the catalyst particles in the DS are given by [11]: dC i N, ~ Yi N I - Y j N, =--+ z_., " (20) dz D x, ~ g*, D:I

with the boundary conditions: z=O at N~=O a n d z=z o at Ci(zo)-Cis where z is the characteristic length of the pellet. The mass balance equation for the catalyst particles in the HS are given by: d2C, 2 dC, p.,.R' ~ + . . . . (21) dr 2 I" dr D ea

232

de,

with the boundary conditions: Ci=Ci~ at r=rp and--~r =0

at

r=O

In the presence of external heat transfer resistance for the catalyst pellet, but negligble intraparticle heat transfer resistances (which is the case of nickel catalyst), the energy balance equation can be reduced to: Ts = T + ( - A H-' -! ps ! r 2R' dr '

9

(22)

-

h r;

.

The performance of the catalyst pellet in both reactor sections is expressed in terms of the effectiveness factors of each reaction. The effectiveness factors are computed using the following well-known relation:

l

r,

qi = , [ R i dr rpR,, "o

(23)

3. S O L U T I O N O F THE M O D E L E Q U A T I O N S The catalyst pellet equations in both beds are solved using the orthogonal collocation method [15]. The bulk phase differential equationd are integrated using subroutine DGEAR (IMSL Math/PC-Library) [16]. 4. R E S U L T S AND D I S C U S S I O N The data used to simulate the DS is that of the industrial reactor at Polymer Corporation, Sarnio, Ontario, Canada and is given in Table 3. Before exploring the hybrid reactor performance, we will first study the behavior of a catalytic dehydrogenation reactor producing styrene and having membrane with sweep gas in the permeation side.

Table 3 Specifications Item

Bed length Cross sectional area Catalyst density Catalyst diameter Inlet pressure Inlet t~ml)erature Feed: Ethylbenzene Styrene Benzene Toluene Steam Total molar feed

a n d f e e d c o n d i t i o n for D S Value Dimension

1.70 2.98 2146.3 4.7 2.4 922.59 36.87 0.67 0.11 0.88 453.1 491.87

m m2 kcat/m 3 mm Bar K kmol/h kmol/h kmol/h kmol/h kmol/h kmol/h

4.1 M e m b r a n e r e a c t o r f o r t h e p r o d u c t i o n o f s t y r e n e w i t h i n e r t s w e e p g a s The heterogeneous model of the DS is used to simulate this case in which the industrial reactor is fitted with hydrogen selective membrane and the sweep gas is inert and flowing counter currently. Figure (2) shows a considerable EB conversion and ST yield increase over the case without membrane, espicially at higher sweep gas flow rates. For high enough flow rates of sweep gas, EB conversion approaches 52.7% compared with 45.75% for the reactor without membrane. This represents an improvement of 15.2%. The corresponding values of ST yield are 48.2% and 40.17% indicating an improvement of 20%. The same figure depicts the counter-current case. At low flow rates of sweep gas, and using this configuration, ST yield is less than the case without membrane. A typical value is 41.8% at a sweep gas flow rate of 500 kmol/h. At higher gas

233 rates, styrene yield in case of counter-current mode exceeds t h a t of co-current. Values of 48.8% and 48.2% at 3000 kmol/h are observed for the two configurations respectively. The same pattern is observed for EB conversion with values of 53.1% and 52.7% respectively for the same gas rate. Figure (3) shows the hydrogen partial pressure profiles in the reaction and permeation side for a case of counter current mode having a sweep gas flow rate of 100 kmol/h (which is relatively a low flow rate). As shown from the figure, there is a zone of hydrogen back diffusion (i.e. from the permeation side to the reaction side). This explains the reason for the low conversion values. To overcome this region of hydrogen back diffusion, partial blinding of sweep gas can be used. Figure (4) demonstrates the profiles of hydrogen partial pressures in both reaction and permeation sides in case of sweep gas flow rate of 100 kmol/h. The optimum location of partial blinding is at a length of 0.765 m from reactor inlet. EB conversion and ST yield are increased by 2.0% and 6.5% compared with the counter current without partial blinding mode.

4.2 Hybrid R e a c t o r s i m u l a t i o n Data used for simulation of the HR is that present in Tables 3 and 4. Table 5 gives the results of reactor simulation. HR variables are demonstrated versus the reactor length through figures 5 to 8. As shown in figure 5, the trends of EB conversion and ST yield are similar to that of the conventional reactor, but with higher values and higher selectivity towards ST. The same figure shows the profile of the DS temperature. The profile declines with reactor length as the main reaction is endothermic.

Table 4 Specification for HS Item Value and unit Benzene Feed 20 kmol/h Cross sectional area 3.0 m 2 Catalyst diameter 1.86x10 "3 m Catalyst porosity 0.35 Catalyst density 1200 kcat/m 3 Inlet temperature 400 K Inlet pressure 1.1 bar Bed length 1.7 m

Table 5 Results of the HR On the other side of the membrane, and by Item Value and unit the continious transfer of hydrogen to the HS, EB conversion 0.518 benzene conversion increases. The rate of ST yield 0.466 benzene hydrogenation is increased sharply ST production rate 17.855kmol/h after 0.5 m from reactor inlet due to the high D.S exit temperature 841.4 increase in reaction temperature. The other D.S exit pressure 2.36 BZ conversion (H.S) 0.40 part of the figure depicts the temperature of H.S exit temperaure 957.4 the HS of both the bulk gases and the catalyst H.S. exit pressure 1.097 pellets. Small differences between the two temperatures are due to high external heat transfer coefficient (h) between the bulk gas and the catalyst pellet. Figure (8) demonstrates the hydrogen partial pressures in the two reactor sections. Both profiles exhibit a maximum and that of the HS approaches zero after 1.2 m from reactor inlet. This behavior offers a maximum driving force for hydrogen permeation between the two reactor sections.

234

4.3 Effect of use of i n d u s t r i a l m e m b r a n e on reactor p e r f o r m a n c e In order to predict the performance of the reactor on a more practical basis, a commercial membrane was used in the Simulation. Figure (9) shows the profit gain (over the conventional reactor) when using a commercial composite membrane of Bend Inc. in the proposed hybrid reactor. Membrane area was varried to predict the range of ecconomical optimum design. Calculations were based on membrane cost of $300 /ft 2 [17]. The increase in styrene yield was calaulated with a MT price of $660. Rate of return on investment was calculated based on the profit gained with respect to the additional investment (membrane cost). The optimum membrane area can be taken at the point of maximum rate of return. 4.4 Possible process design i m p r o v e m e n t s o f styrene p r o d u c t i o n p l a n t s The above simulation results indicate some of the benefits associated with the use of the HR configuration. The benefits are not limited to the increase in ST yield and selectivity but also includes the useful utilization of the permeated hydrogen to produce cyclohexane. The improvements in the design of such reactors can allow for the complete elimination of the recycle step for the unreacted EB. Exchange of heat between the two beds is assumed to take place espicially that new approaches of membrane preparation on highly conductive substrates are now feasible. Figure 9.a demonstrates the profiles of EB conversion and ST yield, reaching values of 93 and 87% respectively. This value of ST yield represents an enhancement of 116% over the industrial conventional reactor. Figure 9.b depicts the temperature profile in the DS, which shows a minimum. This behavior is due to the competing effect between heat lost in the endothermic dehydrogenation reaction and heat gain by the transfer of heat from the HS. 5. CONCLUSIONS Hybrid configuration improves the performance considerably, namely increasing EB conversion, ST yield and selectivity. The magnitudes of the increases are quite appreciable together with the production of additional product, which is cyclohexane. A design for the HR was proposed with longer length compared with the industrial case, as thermodynamic barrier has been broken, and was found to give values of styrene yield which is 116% more than the industrial reactor. The quantitative predictions obtained using this rigorous model should be reasonably reliable, however, they still need to be checked experimentally. 6. R E F E R E N C E S 1 2 3 4 5 6 7 8 9

N. Itoh, AICHE, 33 (1987) 1576. N. Itoh, J. Chem. Eng. Jpn., 24 (1991) 664. Y.V. Gokhale, R.D. Noble and J.L. Falconer, J. Memb. Sci., 77 (1993) 197. E. Gobina, K. Hou and R. Hughes, J. Memb. Sci., 105 (1995) 163. R. Govind and D. Atnoor, Ind. Eng. Chem. Res., 30 (1991) 591. J.P. Collins and J.D. Way, Ind. Eng. Chem. Res., 32 (1993) 3006. K.L. Yeung and A. Varma, AICHE J., 41 (1995) 2131. B.K. Abdalla and S.S. Elnashaie, J. Membr. Sci., 85 (1993) 229. N. Itoh, J. Chem. Eng. Jpn., 23 (1990) 81.

235 10 E. Gobina and R. Highes, Chem. Engng. Sci., 51 (1996) 3045. 11 E.A.Mason and A.P.Malinauskas, Elsevier, Amsterdam 1983. 12 S.S.Elnashaie, B.Abdalla and R.Hughes, Ind.Eng.Chem.Res.,32(1993) 2537. 13 J.G.P. Sheel and C.M. Crowe, Can. J. Chem. Eng., 47 (1969) 183. 14 J.K. Marangozis, B.G. Mantzouranis and A.N. Sophos, Ind. Eng. Chem. Prod. Res. Dev., 18 (1979) 61. 15 J. Villadsen and M.L. Michelsen, Prentice Hall, New York, NY, 1987. 16 A.C. Hindmarsh, Gear, Lawrence Livermore Lab., 1974. 17 Bend Research Inc., (1996).

236

x/

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Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

237

A SIMPLE AND FLEXIBLE MICRO REACTOR FOR INVESTIGATIONS OF HETEROGENEOUS C A T A L Y T I C GAS PHASE REACTIONS a,i,

G. Veser ' , G. Friedrich a, M. Freygangb, and R. Zengerleb a Institut f'tir Chemische Verfahrenstechnik, Universit~it Stuttgart, D-70199 Stuttgart, Germany; b Hahn-Schickard-Gesellschaft, Institute of Micromachining and Information Technology, D-78052 Villingen-Schwenningen; Germany;

1.

INTRODUCTION

Micro reactors, i.e. chemical reactors with characteristic dimensions in the sub-millimeter range, seem to hold great promise for novel chemical process routes (Lerou 1996, Wengeng 1996). Among their proposed advantages for chemical processes are: the very small thermal inertia, allowing for a very direct control of temperature as a very critical reaction parameter; the inherent safety due to both the small reactant volume being present at any time in the reactor and the well controllable reactor and reaction conditions; and their small dimensions, making it easy to integrate micro reactors into existing processes or to use them where space requirements are critical. In addition to those general reactor engineering considerations, micro reactors offer the additional advantage for heterogeneously catalysed gas phase reactions of allowing for very large surface to volume ratios in well defined reactor set-ups. These large surface to volume ratios should at least theoretically allow for an effective suppression of homogeneous gas phase reactions, since free surfaces typically are strong sinks for radical species which are required to keep homogeneous gas phase reaction alive by propagating and branching radical chain reactions. Therefore, it should be possible to conduct a heterogeneously catalysed reaction involving a mixture of potentially flammable (if not explosive) gases in a micro reactor without danger of open flames and explosions. This has two-fold advantages: from an applied point of view, the reactor becomes intrinsically safe to operate due to the suppression of explosions. From a scientific perspective, such a micro reactor makes it possible to study heterogeneously catalysed hightemperature reactions without influences by parallel homogeneous reaction pathways, making it a very valuable tool for research into this class of reactions. Very advanced micro-manufacturing methods have been developed over the past 10 years, such as laser ablation methods and X-ray based lithographic methods (LIGA and similar methods, Wise 1991, Bacher 1994, Wengeng 1996). However, these techniques are still quite expensive and require very specific equipment. In contrast to this, due to the well developed semiconductor industry silicon based micromachining is a technique which is not only readily available, but also very cheap. It seems therefore very well suited for exploratory studies at this rather early stage of development for chemical micro reactors, since it allows for a comparatively quick and cheap prototyping of such reactors. Furthermore, the good mechanical strength, high thermal conductivity, and relative chemical inertness of silicon make it an ideal candidate for this type of studies. It is for these reasons that we choose silicon as the base material for our current micro reactor design. In the following, we will describe the reactor design and fabrication in more detail, report first experimental experiences with this reactor, and discuss possible practical applications as well as next steps in the further development of the current design.

* correspondence author ([email protected]). Financial support by "Deutsche Forschungsgemeinschafl" through a fellowship to G.V. is gratefullyacknowledged.

238 2. REACTORFABRICATIONAND SET-UP The development of a micro reactor at this early exploratory stage should take into consideration an easy handling of the reactor set-up, suitability of the micro reactor for a wide range of reactions, reusability of all parts of the reactor, and an easy and cheap fabrication technology. Therefore, the decision was made for a modular system. The reactor chamber is defined through two stacked silicon chips, into which channel and hole structures are etched (fig. 1). The catalyst can be added as a wire which is then chucked into the channel structure. The reactor housing presses the pair of Si chips which contain the catalytic wire, avoiding gas leakage due to the atomic-scale flatness of the silicon surfaces. The realised system thus incorporates the following main design considerations: - easy changing of the catalyst material by changing the wire material - easy changing of the reactor wall material by different thin film deposition of the silicon - easy handling of the housing and reuse of all components - cheap batch processing of the silicon wafers. 2.1.

micro

reactor

chips

For manufacturing the reactor chips in a way as simple as possible and at least cost, we chose a process with only two masks for photolithography and only one step of wet etching with KOH. The wafers were etched from both sides simultaneously to a depth of a little more than half the wafer thickness of about 525 ~tm. This resulted in the formation of a reaction chamber of some 0.167 mm~ cross-sectional area and a length of about 20 mm. The cross sectional area of the gas inlet bore is designed to equal the gas channel in the reactor chamber. The outlet hole area is 1.5 times larger than the gas channel cross section. The exact dimensions of the channel bottle-necks in which the wire is f'Lxed are defined by means of the common self stopping mechanism of the KOH wet etch process. After etching the wafer in a fast and cheap batch process, several wafers were coated at 950~ with a silicon oxide thin film of 400 nm thickness for testing purposes. As a further modification, on some of these oxidized wafers LPCVD silicon nitride was deposited at a thickness of about 110 nm. The nitride films were then tempered for one hour at 1100~ to reduce intrinsic layer stresses.

Fig. 1: The etched Si-chips with the micro-reaction channel at the center. The holes for the electrical contacting of the catalytic wire can be seen on the bottom half of the reactor, the top half contains the holes for the gas inlet and outlet. The catalyst is introduced into the micro channel in the form of a wire. This has several advantages over a catalytic coating of the channel walls: the catalyst can be exchanged very easily

239 between experiments, and can also very easily be examinated after the reaction. The same micro chips can be re-used many times for different catalysts. Most metals are readily available in the form of wires with different diameters, and can also be handled easily in this form. By varying the wire diameter, the total catalytic surface area as well as the surface-to-volume ratio can easily be changed without the necessity for any changes in the micro reactor itself. Furthermore, in one of the few previous studies with a catalytic micro reactor in which the micro channel walls were catalytically coated with platinum, the authors reported degradation problems at temperatures above 700~ probably due to thermal stresses and thermal expansion mismatch between the Si-based substrate and the Pt film (Srinivasan 1997). This should not be a problem using this simple catalytic wire set-up. However, the wire set-up has also obvious disadvantages in comparison with catalytically coated walls. First of all, the heating of the wire through resistive heating - while very simple - does not allow for an exact heating of particular zones of the catalyst, but always leads to a "global" heating over the whole length of the reactor. Furthermore, the positioning of the wire in the micro channel, particularly at elevated temperatures is much less well defined than a precise coating of the channel walls. And finally, the catalytic wire in the center of the reaction channel leaves the large area of the channel walls unchanged, w h i c h - depending on the chemical nature of the wall material - could catalyse side reactions in some applications. While this should not be a problem for the type of reactions discussed in this paper, it cannot be generally excluded a priori for other reactions. Nevertheless, the wire set-up forms a very simple and particularly flexible configuration for catalytic micro reactor studies.

2.2. reactor housing The design goal for the micro reactor housing was to ensure an easy and flexible handling of the micro reactor itself, i.e. a quick and easy access to the micro reactor as well as an easy connection to standard (macroscopic) laboratory equipment. Furthermore, the equipment (i.e. the micro reactor as well as the housing) should be re-usable many times to keep development cost low. The reactor housing is therefore built completely from stainless steel (fig. 2) and is comparatively large with a diameter of about 30mm and a thickness of about 7.5mm.

Fig. 2: The stainless-steel micro reactor housing (diameter 50 mm, thickness 7.5 mm) with the two gas inlets (lower right) and the gas exit (upper left). The inner surfaces of the housing are manufactured by spark-erosion to produce a flat, smooth surface contacting the silicon micro chips (fig. 3). This not only helps to ensure a well defined positio-

240 ning of the chips in the housing, but also allows for a gas-tight connection between the micro chips and the housing at the gas inlet and outlet without the necessity for additional sealings.

Fig. 3: The open reactor housing with the top and bottom parts of the micro reactor. In the bottom part of the micro reactor (left), the catalytic Pt wire is visible. The catalytic wire in the micro channel is contacted electrically through small metal screws, which are insulated against the steel housing by ceramic shafts. This allows a direct, resistive heating of the catalyst for the purpose of igniting or thermostating the reaction. The reaction gases can be fed separately through two steel tubings with an inner diameter of 0.5 mm and outer diameters of 1.5 mm. Connections to the peripherals can be made through standard Swagelock-type connectors, allowing for an very easy connection to standard laboratory equipment. The feed gases are mixed by perpendicular impingement at the reaction channel entrance. This mixing right at the beginning of the catalytic reaction zone helps to avoid possible homogeneous reactions between the reactants in the feed tubing which can pose a problem for the type of high temperature oxidation reactions studied here. 3. EXPERIMENTALMEASUREMENTS

3.1. The H2+O2 reaction We have so far tested and validated the applicability of this simple micro reactor set-up using the Pt-catalyzed H2+O2-reaction (H2 + 0.5 02 = H20). This reaction constitutes a particularly interesting test case for the current system for several reasons: The reaction is known to proceed very fast both homogeneously in the gas phase and heterogeneously catalyzed by platinum. The reaction is characterized by particularly wide flammability limits (about 3 vol% - 75 vol% H2 in air) and very high flame velocities, i.e. strong explosions (Lewis 1987). While the reaction is strongly exothermic (AH -- -240 kJ/mol) and is therefore of interest as a possible energy (heat) source, this danger of strong explosions restricts its use for practical applications severely in that the reaction is typically conducted with mixtures below the lower flammability limit to ensure safe operation under all conditions. Therefore, if it would be possible to effectively suppress homogeneous reactions (and thus open flames and explosions), the potential of this reaction could be much more efficiently exploited. (This can be demonstrated with a simple estimate of the adiabatic temperature rise of the H2+O2reaction: while this temperature rise at the lower flammability limit in air (3 vol% H2 in air) is only about ATad= 250K, under stoichiometric conditions in air it increases to about 1600K. This value could even be further strongly increased by reducing the nitrogen content, i.e. replacing air by pure oxygen. Although this calculated temperature rise constitutes a theoretical value which obviously cannot be achieved in a practical set-up due to heat losses and limitations on reactor materials, it nevertheless emphasises the potential of this reaction as a heat source). To test our micro reactor set-up using this reaction, a Pt-wire (diameter 150 m m, purity 99.99%) was used as catalyst in the reaction channel. The reaction gases H2, 02, and N2 were fed through standard mass flow controllers. A thermocouple (K-type) was positioned right at the micro channel exit of the micro reactor to measure the temperature of the effluent reaction gases. The effluent gas

241 stream was then fed to a cryostat to remove the water vapor formed in the reaction, and the remaining gases were fed to a flow meter to measure the total flow rate and to a O~ analyzer to measure the remaining oxygen content. Since the investigated reaction follows a very simple reaction path, the extent of the reaction (i.e. the hydrogen and oxygen conversions) could be calculated from both measures and were found to agree within experimental error (< 5%). For ignition of the reaction, synthetic air was introduced at a flow rate of about 0.1 slpm (standard liters per minute), and then H2 was added to the feed gas stream. In contrast to our experience in 'macroscopic' fixed-bed catalytic reactors, the reaction did not ignite at room temperature at any H2/O2 ratio. At this point, we do not yet know whether this is due to finite size effects in the micro channel, or whether the heat losses in the micro reactor are simply too high for ignition of the reaction at room temperature conditions. Further experiments and simulation studies will be conducted to elucidate this point. The catalytic wire is therefore preheated resistively in an atmosphere about 5% H2 in air in order to ignite the reaction. At an exit gas temperature of about 100~ ignition of the catalytic reaction is observed. The temperature measured at the reactor exit jumps up to about 300~ (i.e. close to the adiabatic reaction temperature) within a few seconds, indicating ignition of the reaction. The whole ignition process takes place within about 4-8 s. This start-up time is currently limited by the resistive power input to the catalytic wire since at too high currents the wire tends to burn through. 3.2. Qualitative tests To test the thermal stability of the silicon micro chips under reaction conditions, first some test runs with only the silicon micro chips, i.e. without housing, were conducted. In these experiments, the hydrogen content in the feed stream was increased to stoichiometric ratios after ignition the reaction, and then the N2 content was slowly reduced. No temperature measurements could be done under these conditions, but it could be seen by simple visual inspection that the chip was glowing bright red/orange near the entrance zone of the catalyst. We estimate the local reactor temperature therefore to be at least around 1000~ Interestingly, the spreading of this hot spot over the whole length of the catalyst and finally a shifting towards the exit zone could be nicely observed under these experimental conditions during increase of the total flow rate through the micro channel. Following each experiment, the micro reactors were inspected for possible deterioration due to the exposure to the very high temperatures and reactive atmospheres. For none of the tested reactor materials (pure Si micro chips, as well as oxide and nitrite surfaces) any degradation was observed. Furthermore, no sintering of the top and the bottom parts of the micro reactor was observed either, making an exchange of the catalyst between runs a rather easy and quick task. Finally, no differences in the reactor behavior or the conversions were found between the different reactor materials, further strengthening the point that at least for this reaction the micro channel walls seem to be completely inert. The Pt wire roughened visibly over the course of several experiments which is a known phenomenon in Pt catalyzed oxidation reactions (Satterfield 1991). No evidence for a net loss of Pt was found in our experiments, however, in some runs a breaking of the catalytic wire occurred in the micro channel. We do not attribute this to the roughening process of the catalyst but rather to local thermal stresses at points where the hot catalytic wire touches the micro channel walls due to the considerable thermal expansion of the wire at these high reaction temperatures. While this does not pose a problem during the reaction, it means that after the end of that particular experiment the catalytic wire has to be exchanged since obviously no resistive heating is possible any more for the ignition of the reaction. A well defined positioning of the catalytic wire and/or an additional means of pre-heating the micro channel will therefore be an important consideration in a further development of this reactor. In summarizing, these first tests showed quite convincingly that all three S i-based materials used for the micro reactor are well suited for the application in these high temperature reactions, with respect to their thermal stability as well as their reactive inertness.

242

3.3. Quantitative measurements For some first quantitative measurements of exit gas temperatures and hydrogen conversion during the reaction, the micro reactor was set into the reactor housing. As before, the reaction was ignited by heating the catalytic wire resistively in a hydrogen lean N2/Oz/Hz-atmosphere. After ignition, the hydrogen feed gas stream was varied and gas exit temperature, total gas flow at the reactor exit, and oxygen concentration of the exit gas stream were measured. Increasing the hydrogen content of the feed gas stream after ignition closer to the stoichiometric ratio increases the temperature of the effluent gases drastically. Figure 4 shows typical curves for the temperature of the effluent gases for a feed gas stream of 1 slpm synthetic air with varying H2 content. One can see that even without any oxygen enrichment of the synthetic air an exit gas temperature of close to 700~ is reached at a stoichiometric feed of H2 and 02. Figure 5 shows the H2 conversion and oxygen content for the same experimental conditions. Over a rather wide range of H2/O2-ratios (between about 1.0 and 2.0), complete conversion of H2 is achieved, while below an Hz/O2-ratio of 1.0 (an Hz flow of 0.2 slpm) the hydrogen conversion drops almost linearly. This behavior is also reflected in the change in slope in the temperature curve in fig. 6 around an H2 flow rate of 0.2 slpm. Above this point the observed temperature increase flattens slightly, since complete conversion of the hydrogen content has been reached and any further increase is therefore only due to the increasing Hz content in the inlet gas stream and a correspondingly increased oxygen conversion in the reactor. To avoid possible problems due to overheating of the steel housing, we did not attempt to decrease the Nz content in the feed gas in these first experiments. In general, these experiments show a reactor behavior for this reaction set-up which shows no obvious differences from the behavior in "macroscopic" catalytic fixed-bed reactors. Although the studied reaction system is a very simple one, this can be taken as an indication that the current micro reactor set-up should indeed be well suited for investigations of more complex catalytic reaction gystems such as partial oxidation reactions.

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[slpm]

Fig. 4: Exit gas temperature at the microreactor Fig. 5:1t2 conversion and 02 content of the exit exitfor increasing H2 content in 1 slpm air. gas stream for increasing 112 content in 1 slpm air.

243 4. POSSIBLEAPPLICATIONS It seems remarkable that in the above experiments, full conversion could be achieved at flow rates in excess of 1 slpm. This total through-put is surprisingly high for such a small reaction channel, indicating that actually quite reasonable amounts of chemicals can be converted in such tiny reactors at least for very fast high-temperature catalytic reactions (i.e. many catalytic oxidation reactions). Under our experimental conditions, a power of about 72 W can be produced in a single micro channel at the stoichiometric point for the He+O2-reaction. By further increasing the flow rate and/or reducing the nitrogen content of the gas stream, it should be possible to increase this value by at least a factor of 2. This potential could for example be used to preheat an automotive car exhaust catalyst. This preheating is currently the focus of intensive research, since 80-90% of the emissions during typical test cycles are produced in the first 60s, i.e. in the cold start-up phase of the catalyst. Increasingly stringent requirements of ULEV and EZEV legislation can only be met if a substantial improvement in this start-up phase can be achieved. Currently, three different concepts are favored in the literature(Kirchner 1997): 1. Preheating the exhaust gases through an electrically heated pre-catalyst. This concept can be shown to be effective, however, it requires a rather complicated two-staged pre-catalyst upstream of the main catalyst, and the power requirements for the heating of the catalyst are currently still pushing the limits of typical car batteries (Otto 1995, Kirchner 1996). 2. Preheating the main catalyst through an external gasoline fired burner. While this concept allows for a very fast heating of the main catalyst, the increased hydrocarbon emissions in the initial seconds due to incompletely burned gasoline from the burner poses a severe problem for this concept (Kollmann 1994). 3. The most promising concept might be a chemical pre-heat of the exhaust catalyst through a controlled feed of hydrogen to the exhaust gas stream (Kirchner 1997). This concept has the strong advantage that the chemical heat is released right on the catalyst surface, and it produces no additional emissions. However, due to safety considerations, the hydrogen content in the exhaust stream has to be kept well below the lower flammability limit. Another problem of this concept is the fact, that the H2+O2reaction is blocked at temperatures below about 150-200~ by other components in the exhaust stream (mainly CO and unburned hydrocarbons), so that a pre-heating phase before engine start would be required for this concept. Conducting the H2+O2 reaction in a micro reactor could allow a new set-up, which can be seen as a combination of the three above concepts without their respective inherent problems: the Pt wire in the micro reactor needs a rather low energy input to preheat it to ignition temperature for the H2+O2 reaction (P= 4 W for a single wire in our experiments). Thus, the power requirements are kept very low and restricted to only a very short initial time (a few seconds). Reacting an H2/air mixture in a micro reactor would now allow to convert a stoichiometric H2 content, thus preheating the effluent gases of the micro reactor to temperatures around 700~ as seen in the above experiments. This preheat would be similar to the burner concept, although of course at much lower flow rates. However, this "micro-burner" would produce zero emissions (other than water vapor) unlike the gasoline fired burner. Dosing the H2 slightly over-stoichiometrically, it would be additionally possible to retain a low H2 content (e.g. 3 vol.%) in the hot effluent gases, thus again operating in the safe, nonflammable regime once outside the micro reactor and also taking advantage of the chemical pre-heat concept. Furthermore, since the gases would already be at elevated temperatures, blocking effects due to strongly adsorbing exhaust components should no longer pose a problem in this set-up. Therefore, this set-up would combine the most important aspects of each of the three above discussed pre-heating concepts in a way that would circumvent each of their individual disadvantages. Based on an analysis by Kirchner (1997), one can estimate that even without further optimization less than 100 micro channels would suffice to feed enough heat to the catalyst to pre-heat it within about 10s to the required operation temperatures. Combining for example as few as five parallel micro

244 channels into one micro reactor, only 20 micro reactors would be needed in this concept. Due to their small space requirements - which is a major concern in automotive applications - these micro reactors could easily be positioned on the outside of the car exhaust pipes. Obviously, a more detailed engineering of this problem would further focus onto an optimal positioning of these micro reactors along the circumference of the exhaust catalyst - or maybe even implanting the micro reactors in the monolithic catalyst itself. Another area in which the current set-up seems to have an interesting potential is the study of hightemperature catalytic partial oxidation of hydrocarbons. These reactions have been shown to proceed very selectively to syngas and olefins over noble metal catalysts at high temperatures (about 8001200C) and very short contact times (about 1-50 ms) (Schmidt 1995). From the channel geometry and the total flow rate we can calculate extremely short contact times of about 0.1 ms and less for the experiments shown above, demonstrating that it is possible to get into contact time regimes well below those of previous experiments. This range of extremely short contact times seems to hold potential for very interesting new reaction pathways in the complex network of parallel and consecutive reactions typical for oxidation reactions (see for example Goetsch 1996 and Witt 1997). However, due to the high temperature requirements for these process routes and therefore the inherent danger of homogeneous side reactions (and thus flames or explosions), it has been studied only very little so far. As we could show using the example of the H2+O2 reaction, homogeneous reactions can be very effectively suppressed in a micro reactor. This indicates that this configuration might also be a very good candidate for the study of this class of reactions, in which the micro channel could then be regarded as an isolated, idealized single monolith pore.

5. FURTHER DEVELOPMENTS

From the above discussion of the first tests of the Si-based catalytic micro reactor it becomes apparent that the current set-up is in principle well suited for the intended goal of studying hightemperature catalytic oxidation reactions. Therefore, further developments will now focus on refining the current, very simple design without abandoning the simplicity and ease of use of the current design. The most important next step in the reactor design will be the implementation of local heater elements on the micro chips. This will allow to pre-heat the catalyst to ignition temperatures without the necessity of a resistive heating, thus circumventing the problems that occur when the catalytic wire breaks during an experiment. Furthermore, this should allow for a zone-wise heating of the reaction channel, thereby for example imposing a temperature profile over the reactor length, allowing for an 'engineering' of the reaction conditions in the micro reactor Such local heaters can relatively easily be implemented in the current design by thin film design as already used for example in the design of chemical sensors (Semancik 1993). When not being used as heaters, such elements also allow to measure the temperature profiles along the reactor axis yielding additional important information about the local reaction extent. Furthermore, this configuration would make the catalytic micro reactor a very flexible and well defined tool for catalytic ignition-extinction studies, which in turn can yield important 'finger-print' information about essential reaction steps in these types of reactions (Veser 1996, Ziauddin 1997).

245 6.

SUMMARY

The system of building up a micro reactor with simple reusable silicon chips has been shown to be a cheap and flexible way to investigate the advantages and disadvantages of a micro reactor for the study of high temperature catalytic oxidation reactions. We realized such a silicon based, very simple reactor in which the design emphasis has been on the ease of handling of the micro reactor. The reactor set-up was tested using the Pt-catalyzed H2+O2 reaction. We could demonstrate that the current reactor design is well suited for the intended task, withstanding very high temperatures without any degradation and very effectively suppressing homogeneous flames or explosions in this reaction system. Very high flow rates in excess of 1 slpm, corresponding to catalytic contact times of about 100 m sec, could be achieved at complete conversion of stoichiometric H2 feeds in air. We see a promising potential use of this set-up in the application of the H2+O2 reaction in a micro reactor for pre-heating of an automotive exhaust catalyst. Due to the effective suppression of flames in a micro channel, the strong exothermicity of this reaction near stoichiometric conditions can be put to use in a micro reactor set-up. A further use of the current reactor set-up can be seen in the study of high-temperature catalytic partial oxidation of hydrocarbons. 7..REFERENCES

W. Bacher et al. (1994), Naturwissenschaften 81,536-545 D. Goetsch and L.D. Schmidt (1996) Science 271, 1560-1562 T. Kirchner and G. Eigenberger (1996), Chem. Eng. Sci. 51, 2409-2418 T. Kirchner (1997), PhD thesis, University of Stuttgart, Germany. K. Kollmann, J. Abthoff, and W. Zahn (1994), SAE Papers 940469. J.J. Lerou et al. (1996), in: qVlicrosystem Technology for Chemical and and Biological Microreactors', DECHEMA, Mainz. B. Lewis and G. von Elbe (1981), 'Combustion and Flames and Explosions of Gases', Academic Press, New York. E. Otto et al. (1995), MTZ Motortechnische Zeitschrift, 56, 488-498 C. Satterfield (1991), 'Heterogeneous Catalysis in Industrial Practice', McGraw-Hill, New Yore L.D. Schmidt, M. Huff, and S. Bharadwaj (1995), Chem. Eng. Sci. 49, 3981-3994 S. Semancik and R.E. Cavicchi (1993), Appl. Surf. Sci. 70/71,337-346 R. Srinivasan et al. (1997), AIChE J. 43, 3059-3069 G. Veser and L.D. Schmidt (1996), AIChE J. 47, 1077-1086 R.S. Wengeng, C.J. Call, and M.K. Drost (1996), presentation at 1996 spring national meeting of AIChE. K.D. Wise and K. Najafi (1991), Science 254, 1335-1342 P. Witt and L.D. Schmidt (1998), J. Catal., in print M. Ziauddin, G. Veser, and L.D. Schmidt (1997), Catal. Lett. 46, 159-167

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Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

T h e u s e of a c a t a l y t i c reactions

wall reactor

for s t u d y i n g

247

highly

exothermic

B. Amon, E. K l e m m and G. Emig L e h r s t u h l fiir Technische Chemie I, University of E r l a n g e n - N u r e m b e r g , Egerlandstr. 3, 91058 Erlangen, G e r m a n y

Abstract A Catalytic Wall Reactor operated under isothermal conditions was used to study both the steady state kinetics as well as the deactivation kinetics of the vapour phase hydrogenation of nitrobenzene to aniline using a p a l l a d i u m supported on (z-alumina catalyst. The activity of the catalyst declined with time on s t r e a m due to parallel coking. Axial coke profiles were m e a s u r e d by total carbon analysis. S e p a r a t i n g the initial reaction rate from the rate of deactivation, the kinetic p a r a m e t e r s of the overall reaction could be derived. The results show t h a t the hydrogenation of nitrobenzene follows a Langmuir-Hinshelwood approach considering the reaction between one adsorbed nitrobenzene molecule and one adsorbed hydrogen atom as the rate determining step. By fitting calculated to m e a s u r e d coke profiles an adequate deactivation model could be derived.

1. I N T R O D U C T I O N Kinetic m e a s u r e m e n t s for highly exothermic reactions represent a difficult problem in reaction engineering, since very often isothermal reaction conditions cannot be ensured. In integral fixed-bed-reactors e.g., reproducible e x p e r i m e n t a l conditions can only be adjusted for weakly exothermic or endothermic reactions (AHR < 60 kJ/mol) [1]. It is true t h a t differential recycle reactors show isothermal behaviour, but they are often very complicated in use. The Catalytic Wall Reactor (CWR) combines the a d v a n t a g e s of the fixed-bed reactor and the differential recycle reactor [2]. Because of the excellent t r a n s p o r t of h e a t directly t h r o u g h the wall, the catalyst surface t e m p e r a t u r e can be easily controlled leading to isothermal conditions inside the reactor. The t r a n s p o r t of m a s s between the gas bulk inside the reactor and the catalyst on the wall of the reactor can be improved by an inert bed. Therefore, the kinetic m e a s u r e m e n t s are not limited by m a s s t r a n s f e r in the film. The CWR can be used to m e a s u r e not only steady-state, but also unsteadystate kinetics such as the kinetics of catalyst coking. By s e p a r a t i n g the initial reaction rate from the time dependent one the m e c h a n i s m of both the deactivation and the m a i n reaction can be modelled. The essential a s s u m p t i o n

248 for this approach is t h a t the reaction rate obtained after extrapolation towards time zero is the rate for the m a i n reaction on the fresh catalyst. Information about the developing coke profiles along the reactor can be obtained by stopping the reaction after certain times on s t r e a m and scratching off the coked catalyst from the reactor wall in intervals. The carbon content of the coked catalyst samples m a y e.g. be determined by total carbon analysis (TC). With the knowledge of the deactivation kinetics based on the partial pressures of the r e a c t a n t s and the reaction t e m p e r a t u r e , it is possible to make predictions about the reactor behaviour by m e a n s of appropriate simulation studies. The hydrogenation of nitrobenzene to aniline is deactivated by coking of the catalyst and has a high heat of reaction (AHR=-443 kJ/mol). Therefore, this reaction system was chosen for the kinetic m e a s u r e m e n t s in the CWR.

2. EXPERIMENT 2.1. E x p e r i m e n t a l Set-up The experiments were performed in a fully a u t o m a t e d experimental set-up. It was possible to feed nitrogen as inert gas, hydrogen as hydrogenating gas, oxygen to b u r n off the coke and m e t h a n e as internal s t a n d a r d for the gas c h r o m a t o g r a p h as well as the fluids nitrobenzene and aniline. The flow rates of both the gases and the fluids were controlled by mass flow controllers (Bronkhorst Hi-Tec). Gaseous and fluid reactants were mixed and the fluids were vaporised in a commercial evaporater (CEM; Bronkhorst Hi-Tec). The bypass line allowed controlling the adjusted flow rates and starting the reaction at a certain point in time by switching the two three-way-valves. A small p a r t of the gas composition was analysed online by a gaschromatograph (Chrompack CP9001) equipped with a CP-SIL 5CB fused silica column and a flame ionisation detector.

2.2 Assembly of the catalytic wall reactor and preparation of the catalyst Fig. 1 shows the CWR used for the experimental investigations. The CWR consists of several reactor segments with varying length and an inner d i a m e t e r of 10 mm. The segments possess screw threads on both ends so t h a t they can be interlinked by a nut and sealed by graphite gaskets. The individual tube s e g m e n t s are m a n u f a c t u r e d out of V4A-steel. The inner wall of the reactor segments was coated with catalyst by m e a n s of a suspension process. An industrial Pd-A1203 (1.1 wt%) catalyst in the form of spheres with a diameter of about 5 m m was ground to particles smaller t h a n 250 ~m. The catalyst powder was suspended in a solvent. While the closed tube s e g m e n t was rotated the catalyst suspension was injected into the tube t h r o u g h a hole in one of the covers. The solvent evaporated and the inner tube wall was coated with the catalyst. Finally the catalyst was calcinated at a t e m p e r a t u r e of 250~ for 8 h.

249 The reactor was h e a t e d by external h e a t i n g devices. To ensure the isothermal operation of the CWR the t e m p e r a t u r e very close to the catalyst surface was m e a s u r e d by thermocouples positioned in radial holes in the wall of the reactor. The thermocouples were plugged to a computer-controlled data logging system which read and stored the t e m p e r a t u r e every 60 s. The t e m p e r a t u r e in the gas bulk (axial direction) was determined by a movable thermocouple in the centre of the tube. The t r a n s p o r t of m a s s between the gas bulk and the catalyst inside the reactor was improved by inert particles.

2.3 E x p e r i m e n t a l c o n d i t i o n s All experiments were carried out at atmospheric pressure. The t e m p e r a t u r e was Figure 1. Catalytic wall reactor varied between 275~ and 425~ the inlet partial pressure of nitrobenzene between 1.7 k P a and 6.8 kPa. The molar ratio of hydrogen to nitrobenzene was adjusted between 3:1 and 24:1. The influence of aniline was investigated at p a r t i a l pressures of 1.7 kPa, 3.4 k P a and 6.8 kPa. Tube segments with a catalyst coating between 2 cm and 12 cm in length were used for the experiments. The catalyst weight varied between 0.15 g and 1.8 g. To activate the catalyst, 10 vol.-% hydrogen was fed to the reactor at reaction t e m p e r a t u r e , before the hydrogenation reaction was started. The reaction was stopped after I h, 2 h, 5 h and 10 h, respectively. To remove sorbed components from the catalyst, the reactor was flushed with nitrogen for 4 h at reaction temperature. 2.4 D e t e r m i n a t i o n of the coke profiles a l o n g the r e a c t o r After every experiment the reactor was cooled down and the reactor segments were dismantled. The coked catalyst was scratched off the tubes in intervals of I cm and the carbon content of each probe was determined by TC analysis (C-Mat 550; StrShlein Instruments). The TC was calibrated with a mixture of CaO/CaCO3 containing I wt% carbon.

3. E X P E R I M E N T A L R E S U L T S To exclude the possibility of reactions at the free m e t a l surface of the reactor, at the inert bed or the catalyst carrier (A1203), experiments were performed

250 under reaction conditions without the active catalyst component Pd. The conversion of nitrobenzene was always less than 1% and therefore negligible in the hydrogenation experiments. An adulteration of the kinetic data by limitation of inner and outer mass transfer could be excluded by appropriate experiments. The measurements of the conversion in dependence on the bulk velocity and catalyst thickness showed no significant change in conversion. The selectivity to aniline was always higher than 99.7 %. No by-products except of coke on the catalyst could be quantitatively determined.

3.1 Steady-state behaviour Figure 2 depicts the measured 3.0 conversion of nitrobenzene depending on the reaction time at a reaction r--,-i temperature of 275~ and a constant ~_. 2.0inlet partial pressure of hydrogen of rn 10.2kPa. The length of the catalyst • Z 1.0 coating amounts to 2 cm which corresponds to a catalyst weight of 0.0 I I I 0.15 mg. The curves clearly show the decrease in the conversion of 0 150 300 450 600 nitrobenzene due to the loss of activity t [min] by coking of the catalyst. The higher the Figure 2. Conversion of inlet pressure of nitrobenzene is, the nitrobenzene vs. reaction time higher is the initial conversion and the (pNB,0= 1,7kPa , ; pNB,0= 3,4kPa ~; stronger the loss in activity. The results for determining the pNB,0= 6,8kPa -) initial conversions and the initial reaction rates, respectively, at different reaction conditions were obtained by extrapolation of the conversion curves to time t = 0 min. The reaction conditions have already been described in Section 2.3. Figures 3 and 4 show the initial conversions in dependence of the partial pressures of nitrobenzene (Figure 3) and hydrogen (Figure 4), respectively. The curves in both Figures show an inhibition of the initial reaction rate by both nitrobenzene and hydrogen. The inhibition by nitrobenzene is only considerable at low hydrogen partial pressures (Figure 3), whereas the inhibition of the reaction by hydrogen is more pronounced because of the strong adsorption of hydrogen molecules at the active sites of the catalyst (Figure 4).

251 4.0 1:3..

5.0

3.0

.0

..=. ,~ 2.0

X

m

Xz 2 . 0

Z

1.0 0.0

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

3.0.............

i.............

i .............

i .............

....

iI . . . .

,I . . . .

,I . . . .

1.0,,, 0.0

,

,

,

2.0

4.0

6.0

u.u 8.0

0.0

12.5

PNB,O [kPa]

Figure 3. Initial conversion at pH2,o= 10,1 k P a . , pH2,O= 20,2 k P a and pH2,O= 40,4 k P a 9

25.0

37.5

50.0

PH2,0 [kPa]

Figure 4. Initial conversion at pNB,O= 1,7 k P a . , pNB,O= 3,4kPa ~ and pNB,O= 6,8 k P a 9

E x p e r i m e n t s in which aniline was fed to the reaction mixture showed t h a t aniline has neither a noticeable influence on the initial reaction rate nor on the deactivation. W h e n the reaction t e m p e r a t u r e was varied b e t w e e n 275~ and 425~ the initial conversion increased only slightly.

3.2 Deactivation by Coking For the d e t e r m i n a t i o n of the u n s t e a d y - s t a t e kinetics the same reaction conditions were chosen as for the experiments described in Section 3.1. The only difference was the length of the coating. In the e xpe rime nts p r e s e n t e d in this section the length of the catalyst coating was 8 cm to resolve the coke profile. Figure 5 depicts the axial profiles of 1.0 the coke content depending on time at a reaction t e m p e r a t u r e of 275~ The p a r t i a l p r e s s u r e of nitrobenzene was pNB,0---- 3 . 4 k P a and t h a t of hydrogen pHe,0=20.4kPa. As expected, the coke content increases in time. Besides, the coke content decreases from the inlet to the outlet of the reactor which indicates t h a t nitrobenzene is responsible for the deactivation by coking. Figure 6 shows the coke profiles at different reaction t e m p e r a t u r e s . With increasing t e m p e r a t u r e the coke content increases noticeable.

0.8 o

0.60.49 0.2

....

0.0

'i ....

2.0

I . . . .

4.0

i

6.0

....

8.0

z [cm]

Figure 5. Time development of the coke profile ( l h .; 2h ~; 5 h - ; 1 0 h - ; 20h ~)

252 1.5 1.2

0.7 0.6

-

o~ ~, 0 . 9 0.6

,..____, o

0.5-

0.40.3-

-

0.3

-

.... 0.0

I .... 2.0

0.2

i .... 4.0

6.0

z [cm]

Figure 6. T e m p e r a t u r e dependence of coke profiles after 10h (275~ .; 325~ ~; 375~ .; 425~ , )

, , , , i

0.0

I ....

2.0

l i l l l l i t l l l

4.0

6.0

8.0

z [cm]

Figure 7. Coke profiles at different hydrogen partial pressures after 10h (PS2,0-- 10.1 k P a ~; pH2,O= 20.2 k P a -; pH2,O= 40.4 k P a . )

In Figure 7 the coke profiles at different hydrogen partial pressures are shown. It can be seen t h a t hydrogen counteracts the deactivation by coking. The higher the hydrogen content in the feed the lower the coke content on the catalyst. Also the activity after a certain time is r e m a r k a b l y higher at higher hydrogen p a r t i a l pressures. The experiments with constant hydrogen partial pressures show exactly the reverse results. The higher the nitrobenzene content in the feed the higher the coking rate and the lower the activity of the catalyst. Additional TPO m e a s u r e m e n t s of coked catalyst samples showed two peaks belonging to coke on the active sites (C-a) and coke on the support (C-s). 4. K I N E T I C M O D E L L I N G

4.1 D e t e r m i n a t i o n of t h e rate e q u a t i o n Following the approach of Szepe and Levenspiel [3] the overall reaction rate r(t) is deemed separable and can be written as r(t) = ro(t = 0). a(t) where r(t) is the overall reaction rate, r0 the initial reaction rate and a(t) the activity of the catalyst. With the a s s u m p t i o n t h a t only coke on the active sites is responsible for the decline in activity, a(t) can be expressed as a function of the coke content on the active sites (a(t) = f(wc-a)). Using the same approach for the formation of coke on the active sites (C-a) as for the m a i n reaction following equation can be stated. d(wc_a)

dt

= rc.a, o (t

= 0)- a c-~(t)

253 Also the activity function ac-a(t) can be expressed as a function of the coke content on the active sites (ac-a(t) - f(WC-a)). The rate of formation of coke on the support is assumed to be not deactivated and can be w r i t t e n as d(wc_s) = rc_s,o dt For the a d j u s t m e n t of the model p a r a m e t e r s the software tool Simusolv was used. A n u m b e r of models for both the reaction and the deactivation functions were tested using a one-dimensional, isothermal, pseudo homogeneous, plug-flow model to describe the fluid dynamics of the CWR.

4.2 Steady-state rate equation The best description of the initial reaction rate r0(t=0) was achieved by means of a Langmuir-Hinshelwood approach. Nitrobenzene and hydrogen adsorb at the active sites and the rate determing step is the surface reaction of one nitrobenzene molecule with one dissociatively adsorbed hydrogen atom. According to this model the initial reaction rate can be w r i t t e n as follows

5.0

4.0 ~"

~

-

:

.......-.........

.......

:-

- r/

/ ~..o .....

i +20~

z

:. .......

:

3.0

.....

i,o!

2.0

......

1.0

.-

0.0

0.0

k. K ~ . KH2.p~-P~i~ r0(t = 0) = ( I + K ~ "pNB + K.2 9p~52)2

--i

1.0 2 . 0

3.0 4.0

5.0

PNB,exp[kPa] Figure 8. Parity plot for the initial rate equation

The t e m p e r a t u r e dependence of the reaction rate obeys the Arrhenius law. Figure 8 shows the parity plot of calculated and experimental partial pressures of nitrobenzene.

4.3 Unsteady-state rate equation Taking into account the observed effects described in Section 3.2. following deactivation model could be obtained by fitting the calculated to the m e a s u r e d coke profiles. Formation of coke on active sites: d(w C-a ) __ k d--------~

C-a

p ~ "exp(-k ~ des'C'a W

C-a

)

254

2.0

Formation of coke on support: d(wc-~) dt

1.6

ml

= kc_ s

Psu p~ ,.___,

Deactivation function for main reaction: a(Wc-a ) = exp(-kdes, Wc_ a )

The t e m p e r a t u r e dependence of the kinetic p a r a m e t e r s kc-a und kc-s obeys the Arrhenius law. Figure 9 shows the comparison of calculated and measured coke contents of the catalyst.

1.2

E o 0.8

0.4 0.0 0.0 0.4 0.8 1.2 1.6 2.0 WC,exp [%]

Figure 8. Parity plot for the coke formation

5. C O N C L U S I O N S The experiments showed that a Catalytic Wall Reactor is an appropriate tool to perform both steady-state and unsteady-state kinetic measurements for highly exothermic reactions under isothermal conditions. By separating the initial reaction rate from the deactivation rate it was found that the hydrogenation of nitrobenzene follows a Langmuir-Hinshelwood approach. Descending axial coke profiles measured by total carbon analysis identified nitrobenzene as coke precursor whereas high partial pressures of hydrogen suppressed the formation of coke. Aniline had no influence neither on the main reaction nor on the deactivation by coking. TPO measurements of the coked catalyst showed two peaks belonging to coke on the active sites and coke on the support. The adjustment of calculated and measured coke profiles confirmed the TPO measurements. The best adjustment of the kinetic parameters was obtained using a deactivation model in which coke on active sites was distinguished from coke on the support. This deactivation model which is based on partial pressures and reaction temperature can be used to simulate the reactor behaviour of industrial fixed-bed reactors in the hydrogenation of nitrobenzene. 6. R E F E R E N C E S [1] P. Kripylo, K.P. Wendtland and F. Vogt, Heterogene Katalyse in der chemischen Technik, Dt. Verlag f. Grundstoffindustrie GmbH, Leipzig; 1993. [2] T. Baron, W. Manning and H.F. Johnstone, Chem. Eng. Prog.; 48 (1952) 125. [3] S. Szepe and O. Levenspiel, Catalyst deactivation, in Chemical Reaction Engineering, Pergamon Press, Oxford,1971.

Reaction Kinetics and the Developmentof Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

255

A N o v e l R e a c t o r to A c t i v a t e C h r o m i u m - C a t a l y s t s H. Sch6nfelder a, M. KS_mmerer b, W. de Lange a bELENAC GmbH, G e r m a n y aBASF AG, G e r m a n y 1. A b s t r a c t C h r o m i u m - C a t a l y s t s are widely u s e d for the p r o d u c t i o n of polyethylene (PE) in g a s p h a s e a n d slurry reactors. They are prepared as p a r t i c u l a t e solid material w h i c h is injected into the polymerization e n v i r o n m e n t where the solid p o l y m e r is growing a r o u n d a catalyst particle. The p r e p a r a t i o n of the catalyst is carried out in two steps. First, the catalyst carrier is i m p r e g n a t e d with a solution containing a C h r o m i u m c o m p o u n d . In order to obtain the active material, the C h r o m i u m h a s to be oxidized to the sixth oxidation stage (Cr-VI) in an activation p r o c e s s which is the topic of the p r e s e n t paper. In this process, called activation, the catalyst is c o n t a c t e d for several h o u r s with oxygen at t e m p e r a t u r e s between 500 a n d 900 ~ d e p e n d i n g on the type of PE to be produced. Proper control of the residence time a n d t e m p e r a t u r e d u r i n g the activation process is a c c o m p l i s h e d by use of a fluidized bed reactor which is o p e r a t e d batchwise. In this paper, the u n i q u e features of the Elenac reactor design (LUPOCAT C | Technology) a n d their impact on the process c h a r a c t e r i s t i c s a n d fluid m e c h a n i c s are described. 2. I n t r o d u c t i o n The yearly increase in Polyethylene (PE) p r o d u c t i o n from now until the year 2000 is e s t i m a t e d to be 4.8 %. Based on the c u r r e n t p r o d u c t i o n rate of 37.2 million tons of approximately equal proportions of high a n d low density PE (HDPE a n d LDPE), this e q u a t e s to 47.2 millions tones of PE in the year 2000 [1, 2]. This m a k e s Polyethylene the m o s t i m p o r t a n t plastic. Besides conventional Ziegler catalysts a n d the more recently developed Metallocene catalysts there is an industrially i m p o r t a n t class of C h r o m i u m type c a t a l y s t s to produce m e d i u m a n d high density polyethylene in fluid bed gas p h a s e a n d slurry loop reactor systems. Elenac as major E u r o p e a n PE p r o d u c e r h a s developed a n d m a n u f a c t u r e s a proprietary class of C h r o m i u m catalysts, called LUPOCA T C | Commercial catalysts u s u a l l y contain u p to 2 % C h r o m e on an inorganic, inert carrier. Prior to polymerization, it is n e c e s s a r y to activate the catalysts at high t e m p e r a t u r e s in an oxidative a t m o s p h e r e , in order to modify the C h r o m e as surface C h r o m e (VI). Either d u r i n g the polymerization or a preceding step, the C h r o m e (VI) will be reduced to a lower oxidative stage (cf. Figure 1).

256

H

H

0

0

!

I

carrier

O

O

%//

CrO 3 , AT

0

/

Cr

\

or Or-bonds carder which can react to Cr(VI )-oxide

/Cr\E

"El" 0

O = C

0

0

ll" n (C=C)

carder

0

/

/PE

Or

\

0

carrier

Figure 1. Schematic of oxidation stages of Chrome on the catalyst surface

The catalyst activity and the melt index of the PE-polymer increase with a decreasing Hydroxy population. Everything, which i n f u e n c e s the Hydroxy population will therefore influence both the activity and the molecular m a s s of the polymer. Through calcination p r o c e s s e s (e.g. decreasing Hydroxyl concentration), the activity and melt index can be modified significantly [3]. Scope of the p r e s e n t p a p e r is to p r e s e n t the advanced activation of chrom i u m catalysts in a fluidized bed reactor. The p a p e r gives a general idea of how to operate s u c h an activator u n d e r consideration of material selection, safety and environmental issues. Special attention is given to the a d v a n t a g e s of the extraordinary design of the r e a c t o r s and experimental results concerning the flow s t r u c t u r e in a large scale u n i t are presented.

3. A c t i v a t i o n p r o c e s s

The activation is carried out in a p a t e n t e d u n i t t h a t is schematically depicted in Figure 2. The s y s t e m consists of a fluidized bed reactor, a gas b u r n e r - t o supply heat, a cyclone to s e p a r a t e particles from the off gas stream, a venturi w a s h e r to remove fines from the off gas and a settler to c o n c e n t r a t e the fines in a slurry t h a t can be discharged. In the course of an activation, the vessel is loaded with C h r o m i u m containing material and then slowly h e a t e d u p to the desired activation temperature. Therefore, the gas b u r n e r is u s e d to feed hot flue gases to the

Figure 2. S c h e m a t i c flow sheet of a catalyst activation u n i t

257 h e a t i n g j a c k e t of the fluidized bed reactor. Air is fed to the activator as fluidizing gas. The fluidization of the particles results in vigorous mixing a n d ins u r e s t h a t only m i n o r t e m p e r a t u r e gradients are found in the reactor a n d a s m o o t h h e a t i n g p r o c e s s is achieved. During the process, small particles are continuously e n t r a i n e d from the fluidized bed a n d carried to the reactor outlet. The cyclone at the top of the reactor removes m o s t of the particles from the outlet gas a n d r e t u r n s t h e m to the fluidized bed. The d u s t - l a d e n gas from the cyclone is lead to a venturi w a s h e r where it is cleaned. Finally, the off-gas is cooled down and s e n t to the e x h a u s t . The venturi water is treated in a settler to c o n c e n t r a t e the rem a i n i n g c a t a l y s t particles before discharge. After a few h o u r s of t r e a t m e n t at a c o n s t a n t t e m p e r a t u r e , the catalyst b a t c h is cooled down by reducing the t e m p e r a t u r e of the b u r n e r ' s flue gases. The fluidizing gas is switched from air to nitrogen. Finally, when the temp e r a t u r e is close to a m b i e n t conditions the (Cr-VI) catalyst is discharged. Conventional fluidized beds u s e d as activators often have specific problems d u r i n g the discharging process. It is difficult to completely discharge a catalyst b a t c h t h r o u g h s e p a r a t e openings in the gas distributor plate. Significant a m o u n t s of catalyst may r e m a i n on the grid a n d e n d a n g e r the quality of the s u b s e q u e n t batch. This may either be due to the n o n u n i f o r m residence time distribution or --even worse - due to a mix of catalysts with different specifications. Cleaning p r o c e d u r e s recl~uire the u s e of w a s h i n g water a n d are time c o n s u m i n g . Elenac's L U P O C A T C ~' Technology m a k e s u s e of a very special kind of reactor, the bottom of which is equipped with a cone type gas distributor to solve this problem.

4. R e a c t o r g e o m e t r y The p r e s e n t activators are equipped with a u n i q u e type of gas d i s t r i b u t o r at the reactor inlet. W h e r e a s m o s t fluidized beds are s u p p o r t e d by a gas distributor grid t h a t is designed as a plate with simple bore holes, bubble caps or similar openings [e.g. 4], the p r e s e n t d i s t r i b u t o r consists of a conical section of the reactor wall. At the lower end of this section, the inlet gas e n t e r s the reactor t h r o u g h a single opening a n d forms a low velocity jet. As depicted in Figure 3, particles sliding along the inclined reactor walls are continuously fed to the central jet where they are c a u g h t by the gas stream and entrained upwards. An i m p o r t a n t a d v a n t a g e of the conical bottom section for the present application of a fluidized bed is the ease of the discharge process. Figure 3: Principle flow s t r u c t u r e in Simply by reducing the gas flow a fluidized bed with a conical bottom t h r o u g h the central opening, partisection.

4----F4-".

..

.

9

9

9

.

9

o

.

9

.

.

9

.

9

258 cles slide d o w n to the discharge valve from where they fall into t r a n s p o r t bins. No fluidizing gas is required to s u p p o r t the u n l o a d i n g of the reactor, w h i c h is especially favorable for the u s e of the cyclone. D u r i n g fluidization of the bed, t h e efficiency of the cyclone collapses as soon as the s u r f a c e of the bed d r o p s below the lower e n d of the cyclone's r e t u r n leg. In t h a t case, the p r e s s u r e seal m e c h a n i s m of the r e t u r n leg does not work a n y m o r e a n d fluidizing gas will e n t e r the cyclone t h r o u g h the r e t u r n line from u n d e r n e a t h res u i t i n g in significant loss of particles. A f u r t h e r a d v a n t a g e of the conical gas d i s t r i b u t o r is its ability to cope with cohesive p a r t i c l e s w h i c h are difficult to fiuidize (Geldart C type particles [5]). As h a s b e e n s h o w n by V e n k a t e s h et al. [6], the conical gas d i s t r i b u t o r is especially well s u i t e d for difficult to fluidize particles t h a t t e n d to g e n e r a t e c h a n n e l s a n d s t a g n a n t regions w h e n being a e r a t e d with a s t a n d a r d gas distributor. With a well designed conical gas distributor, c h a n n e l s formed in the s y s t e m will collapse quickly b e c a u s e particles will not form s t a g n a n t regions on the inclined wall. Particles u s e d as C h r o m i u m c a t a l y s t s belong to the G e l d a r t A (good flowability a n d fluidization properties) or C group. The forgiving c h a r a c t e r i s t i c s of the conical gas d i s t r i b u t o r provides s o m e e x t r a confidence in the c a s e of Geldart C type catalysts. Two p o t e n t i a l d i s a d v a n t a g e s of the conical gas d i s t r i b u t o r c a n be identified: 9 G a s / s o l i d c o n t a c t efficiency is low in the e n t r a n c e region b e c a u s e m o s t of the gas f o r m s a gas jet with low solids content. 9 Particles n e a r the wall are poorly (if at all) fluidized a n d m a y easily be exp o s e d to c o n d i t i o n s of o v e r h e a t i n g or overcooling. Little gas is available at the wall of t h e conical section.

Figure 4: E x p e r i m e n t a l s e t u p for investigation of fluidized b e d flow s t r u c t u r e

In the p r e s e n t activation process, the first c o n c e r n is of m i n o r i m p o r t a n c e , since the r e a c t i o n s to o c c u r are slow. The s e c o n d d i s a d v a n t a g e , on the o t h e r h a n d , m a y lead to lower p r o d u c t quality. It is therefore crucial to m i n i m i z e the time particles reside at the wall in the p a c k e d b e d state. An extensive series of e x p e r i m e n t s with small a n d m e d i u m scale s y s t e m s w a s conducted to e n s u r e the p r o p e r choice of the cone ape r t u r e angle. The detailed analysis of the r e s u l t s h a s s h o w n t h a t this angle is of major i m p o r t a n c e for the resuiting m o v e m e n t of p a r t i c l e s along the wall. Neither very steep n o r very fiat g e o m e t r i e s yield optimal flow behavior.

259

5. E x p e r i m e n t a l study 5.1. Equipment The i n c r e a s i n g d e m a n d for a very large p r o d u c t i o n scale of p o l y m e r i z a t i o n u n i t s also r e q u i r e s large activators to provide sufficient a m o u n t s of catalyst. Therefore, t h e q u e s t i o n arises as to how big a n a c t i v a t o r with a conical section c a n be built w i t h o u t significant fluid m e c h a n i c a l p r o b l e m s . W h e r e a s , t h e favorable p r o p e r t i e s of activators with a d i a m e t e r of 0.5 m are k n o w n , no i n f o r m a t i o n is available on s y s t e m s with a d i a m e t e r of 1 m or more. T h u s , it w a s decided to erect a large scale cold m o d e l to investigate the fluid m e c h a n i c a l p r o p e r t i e s of a 1 m ID fluidized b e d with a conical gas dist r i b u t o r . The u n i t built is depicted in Figure 4. In o r d e r to d e t e r m i n e the quality of fluidization, t h r e e different f e a t u r e s were u s e d : 9 F o u r Plexiglas windows allowed visual o b s e r v a t i o n of the particle m o v e m e n t . Black tracer particles were a d d e d in small a m o u n t s (< 1 wt.-%) to improve visibility. 9 Differential p r e s s u r e m e a s u r e m e n t s at v a r i o u s l o c a t i o n s along the c o n e height. 9 M e a s u r e m e n t s with a fiber optical reflection probe. Two p o r t s ( u p p e r a n d lower) allowed i n s e r t i o n of the optical p r o b e t h a t w a s u s e d to q u a n t i t a t i v e l y describe the flow s t r u c t u r e in the s y s t e m . The principle of t h e fiber optical probe is depicted in Figure 5. It s h o w s t h a t the p r o b e

Figure 5. Principle of fiber optical probe. c o n s i s t s of two i n d e p e n d e n t s e n s o r s , w h i c h have n u m e r o u s fibers to s e n d light into the s y s t e m . Particles in the vicinity of the p r o b e tip r e f e c t the light, w h i c h is collected by o t h e r fibers guiding the light to d e t e c t o r s , g e n e r a t i n g a voltage o u t p u t . The h i g h e r the voltage, the h i g h e r the c o n c e n t r a t i o n of re-

260 flecting particles at the probe tip. In the case of gas b u b b l e s located at the probe tip, the signal p l u m m e t s to small values. The probe signal was recorded digitally on a PC, in order to calculate bubble void fractions from the resulting signals. The s a m p l i n g rate was set to 500 Hz a n d records of 4 seconds were taken, resulting in 2000 single d a t a points per record. In order to g u a r a n t e e statistically correct results, 8 records were t a k e n at each position. For a s s e s s i n g the local volume fraction of voids (bubbles or jets), a threshold value was d e t e r m i n e d m a n u a l l y for each d a t a set to be analyzed. All d a t a points below this value were considered as ,void". The volume fraction of voids is then defined as the ratio of m e a s u r i n g points classified as ,void" to the total n u m b e r of points.

5.2. S o l i d s p r o p e r t i e s The solids u s e d in the study were a catalyst carrier m a t e r i a l with properties listed in Table 1. Since typically only small a m o u n t s of active s u b s t a n c e s are required for the catalytic active material, the fluid m e c h a n i c a l properties of the carrier a n d the catalyst are very similar. Table 1. Properties of solids u s e d in fluid mechanical study Bulk d e n s i t y Sauter mean diameter Q3 (200 tim) Qa (125 tim) Q3 (90 tim) Qa (63 tim) Q3 (45 tim) M i n i m u m fluidization velocity (ambient air)

225 78 99 90 88 2 0,3 1,3

kg/m 3 tim wt.-% wt.-% wt.-% wt.-% wt.-% mm/s

5.3. O p e r a t i n g c o n d i t i o n s The u n i t was filled with 300 kg of the carrier and was fluidized with dried a n d oil-free air at superficial gas velocities, u, of 6.4 c m / s a n d 4.4 c m / s in the cylindrical section. The t e m p e r a t u r e and p r e s s u r e were n e a r a m b i e n t conditions.

5.4. Visual o b s e r v a t i o n of flow The particle m o v e m e n t as observed through the Plexiglas section was excellent a n d far more vigorous t h a n in any smaller a p p a r a t u s . Though a l m o s t no voids are p r e s e n t at the walls, m o v e m e n t of particles at t h a t position does occur in the downward direction. Note, t h a t the particles moving d o w n w a r d s close to the wall are probably not in the fluidized state. Particles resting at the wall were rarely detected and this condition never lasted longer t h a n a couple of seconds. It could be easily seen t h a t a reduction of the gas velocity from 6.4 to 4.4 c m / s results in significantly less particle motion at the wall. From t h e s e observations it can be deduced t h a t scale-up of this kind of gas

261 d i s t r i b u t o r d o e s n o t c a u s e p r o b l e m s with the particle m o t i o n a t t h e wall a n d d e a d z o n e s if t h e r e a c t o r g e o m e t r y a n d gas velocity are p r o p e r l y c h o s e n .

5.5. Pressure profile A m a j o r topic for t h e d e s i g n of t h e activator is the e x p a n s i o n b e h a v i o r of t h e c a t a l y s t d u r i n g fluidization. In o r d e r to get i n f o r m a t i o n on this p r o p e r t y of t h e s y s t e m , differential p r e s s u r e m e a s u r e m e n t s were c a r r i e d o u t d u r i n g t h e fluidization e x p e r i m e n t s . A s s u m i n g complete fluidization of t h e particles, t h e pressure gradient at a given p o s i t i o n is a m e a s 2000 u r e of t h e m e a n d e n s i t y E1800 * of t h e bed. It h a s to be c o n s i d e r e d g 1600 t h a t c o m p l e t e fluidization d o e s n o t o c c u r in 14oo mll-- Uthe cone, b e c a u s e s o m e "E 1200 ~ u=4.4 cm of t h e p a r t i c l e s slide along t h e vessel wall a n d 1000 will t h e r e f o r e n o t con800 t r i b u t e to the p r e s s u r e d r o p of t h e gas. T h u s , ~= 600 the differential p r e s s u r e :6 400 measurement in the cone yields v a l u e s of t. , . , . 9 ,o "~ 200 , ' . solids c o n c e n t r a t i o n t h a t '~ 100 110 120 130 140 i 50 are slightly too low. The axial profiles of apparent bed density [kg/m s] the a p p a r e n t b e d d e n s i t y Figure 6. Profile of a p p a r e n t b e d d e n s i t y in coas m e a s u r e d d u r i n g the ne e x p e r i m e n t s are s h o w n in Figure 6. It c a n be s e e n t h a t the axial profile is s u r p r i s i n g l y fiat in the cone region. Above t h e edge b e t w e e n the cone a n d the cylindrical p a r t of the vessel, a d i s t i n c t d e c r e a s e of the solids c o n c e n t r a t i o n t a k e s place. This finding is in c o n t r a d i c t i o n to the m e a s u r e m e n t s of o t h e r a u t h o r s w h o h a v e observed a c o n t i n u o u s i n c r e a s e of the particle c o n c e n t r a t i o n with i n c r e a s i n g d i s t a n c e from the gas inlet [7]. It h a s to be considered, t h o u g h , t h a t the particles u s e d in this s t u d y belong to the Geldart A g r o u p w h e r e a s in m o s t o t h e r i n v e s t i g a t i o n s m u c h c o a r s e r particles (Group D), w h i c h are typical for s p o u t e d b e d s , are u s e d .

~

5.6. Optical probe signals The m e a s u r e m e n t s with the optical probe yielded raw signals, e x a m p l e s of w h i c h c a n be s e e n in Figure 7. The plots c o n t a i n the original signals of b o t h fiber p a c k a g e s i n s t a l l e d in one probe. The c h a n n e l s h a v e different c h a r a c teristics, so t h a t b o t h signals c a n be plotted in one c h a r t w i t h o u t interference. It c a n clearly be s e e n from the u p p e r plot in Figure 7, t h a t a high v o l u m e fraction of voids is p r e s e n t in the c e n t e r of the cone. The signal gives a good

262 indication of the vigorous m o v e m e n t and mixing in the vessel. A c o m p a r i s o n of the two c h a n n e l s shows how both fibers ,see" rising bubbles s i m u l t a n e ously - b u t with a slight time difference of milliseconds. The gas jet g e n e r a t e d at the cone inlet obviously does not extend to an axial length, which is high e n o u g h to detect the jet by the probe. Instead, the behavior of the g a s / s o l i d flow a p p e a r s to be t h a t of a bubbling system. The central plot of Figure 7 shows the d a t a obtained at a radial position that is only 10 cm away from the wall. C o m p a r e d with the u p p e r plot less turbulence is observed. However, there is still a 1,2 considerable fraction of voids. Thus, it can be 1,0 , ,,,,-r'concluded t h a t even at 0,8 r a t h e r small d i s t a n c e s from the wall the parti0,6 . , . , . , . , cles can be considered to 1,6 0 500 1000 1500 2000 be in the fluidized state. M e a s u r e d d a t a from the position at the wall (lowest plot of Figure 7) shows a completely dif--'m 1,0 ] location: upper port, 100 mm from wall ferent behavior of the r u,c~ . , . , . , . , system. No voids c a n be r 500 1000 1500 2000 seen in the plots. Howr 1. ,_6 ]0 ever, the likelihood of voids is not zero. By vis1'4 l ual observation t h r o u g h 1,2 the Plexiglas windows, 1,0~ one can occasionally location: upper port, at the wall detect rising bubbles, 0,8 which m a y occur once a 0 m i n u t e at a given locameasuring time [2 ms] tion.

1,4]

Figure 7. Example of signals from optical probe c m / s . C h a n n e l s 1 and 2 are shifted by 0.2 V to ease reading

u=6.4

5.7.

Local

void

frac-

tions

By evaluating all records taken, the distribution of the local volume fraction of voids can be determined. Figure 8 shows the result of all m e a s u r e m e n t s carried out at the superficial gas velocity of 6.4 c m / s . The local void fraction is plotted as a function of the d i m e n s i o n l e s s radial position for the u p p e r and lower m e a s u r i n g ports. Two major conclusions can be drawn: 1. The d i m e n s i o n l e s s void-profiles are similar at both axial positions. 2. Voids at the lower position occupy less volume although the superficial gas velocity is higher t h a n at the u p p e r position An explanation for the first finding is that the jet generated by the central gas inlet is more quickly split into bubble-like s t r u c t u r e s t h a n expected from the literature [e.g. 8] b e c a u s e of the use of Geldart A particles. These are

263

0,4 0,35 7',

0,3

II ~

~

upper port I

o 0,25 C

=,..

O .m 0,2 w- 0,15 ~" O

O

0,1 0,05 t

0

0,2

I .......

0,4

i_

0,6

I. . . .

0,8

rlR [-]

Figure 8. Radial distribution of void fraction in cone at both m e a s u r i n g levels (u=6.4 c m / s ) .

1

known to exhibit m u c h smaller effective viscosities [9] t h a n c o a r s e r particles, so t h a t voids are less stable. The second conclusion from the d a t a is also surprising since the gas superficial velocity dec r e a s e s with height. It may be the conversion of jet-like, stretched voids with high local gas velocities into more spherical bubbles, which is responsible for the p h e n o m e n o n .

6. Conclusions The technology to produce Elenac's LUPOCA T C | catalyst m a k e s u s e of a special type of fluidized bed reactor, which has unique advantages. 9 The s e c u r e a n d complete discharge process using the conical reactor section maximizes the yield a n d i n s u r e s product quality. 9 Down-time due to the cleaning procedure is minimized a n d no liquid is required for this purpose. 9 C h a n n e l i n g a n d similar fluidization problems i n f u e n c i n g the p r o d u c t quality are avoided by u s e of the conical gas distributor. S u m m a r i z i n g the information obtained by the m e a s u r e m e n t s with the optical probe, it can be stated t h a t scale-up of the unique conical gas distributor for c a t a l y s t s does not create problems concerning the flow structure. The fluidization properties do not s e e m to differ significantly from conventional fluidized bed. This is confirmed by the operational experience t h a t the catalysts p r o d u c e d from a u n i t with a cone type distributor are at least as homogeneous as those from fluidized beds with conventional gas distributors.

7. R e f e r e n c e s 1 2 3 4 5 6 7 8 9

Nachr. Chem. Techn. Lab, 44 (1996), 202 Kunststoffe, 85 (1995) 1532 M.P. McDonald, Adv. Catal., 33 (1985), 47 E.U. Hartge a n d J. Werther, Fluidization IX (1998), 213 D. Geldart, Powder Technol., 7 (1973), 285 R.D. V e n k a t e s h , J. Chaouki, D. Klvana, Powder Technol., 89 (1996), 179 P.G. Romankow, N.B. Rashkovskaya, A.D. Gol'tsiker and V.A. Seballo, Sov. Chem. Ind. 5 (1970), 62 T. Djeridane, F. Larachi, D. Roy, J. Chaouki, R. Legros, Can. J. of Chem. Eng., 76 (1998), 190 J. Werther, Chem.-Ing.-Tech. 49 (1977), 193

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Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

Kinetics and reactor simulation for reactions

267

polyethoxylation and polypropoxylation

E. Santacesaria, M. Di Serio and P. Iengo Dipartimento di Chimica, Universitb. di Napoli "Federico II", Via Mezzocannone, 4 (80134) Napoli, Italy. Abstract

The kinetics of, respectively, polyethoxylation and polypropoxylation reactions in the presence of a hydrophobic starter and a basic catalyst are presented with a general approach based on the reaction mechanism. A general kinetic model and a simplified model useful to simulate the alkylene oxide consumption and the oligomers distribution in well-mixed semibatch reactors are presented. Kinetic laws and related parameters have been tested by simulating a spray tower pilot plant reactor with two alternative models considering the liquid in the drops emerging from the spray nozzle to be either stagnant or well mixed. The later possibility seems to be the correct one. 1. INTRODUCTION Polyethoxylation is normally performed at industrial level in industry to prepare non-ionic surfactants and polymers. The reaction is exothermic and requires an efficient heat exchange [1 ]. The reaction can be promoted by both acid and basic catalysts, but the basic ones are usually preferred, because acid catalysts give place to undesired by-products. In spite of the industrial importance of the mentioned reactions very few papers have been published on the kinetic aspects of ethoxylation and propoxylation. The early works on the subject have been reviewed by Schick [2] and more recently by Nico M. van Os [3] and Santacesaria et al. [4] Ethoxylation and propoxylation of hydrophobic starters promoted by alkaline catalysts, as generally accepted, occur through a nucleophilic substitution SN2 [5]. Therefore, the nucleophilicity of the anion plays a key role in promoting the oxirane ring opening and the differences observed in the activities in the presence of alkoxide, phenoxide and carboxylic anions are dramatic [6-8]. In particular, the carboxylic anion is not nucleophilic enough to open the oxirane ring and the reaction initially occurs with another mechanism favoured by the acidity of the starter [8]. Another peculiarity of fatty acid substrates is that ethoxylated molecules undergo transesterification reactions occurring together with alkoxylation. Alkylene oxide consumption is not affected by the intervention of these reactions because the number of mobile hydrogens and of growing chains does not change. In this case, the kinetic model is complicated by the development of three different distributions, that of monoesters, diesters and polyglycols. More details about this particular system are described elsewhere by Di Serio et al. [8]. The acidity of the starter is another important factor in these reactions influencing the proton transfer or chain transfer equilibria. Then, it must be recognized that a basic catalyst dissolved in a highly hydrophobic starter promote the alkoxylation reactions acting as an ionic pair [6]. As matter of fact, the kinetic behaviour abruptly changes with the kind of alkaline catalyst used. We can roughly say that the large is the ionic radius of the cation, the more active is the catalyst [9]. Besides, activities can also be affected by treating an alkaline catalyst with a crown ether sequestering the cation [ 10]. The influence of the ionic pair is not restricted to the overall catalytic activity, because, by using alkaline catalysts of the second group, narrower distributions of the ethoxylated homologues are obtained, probably as a consequence of the shifting of proton transfer equilibria [ 11].

268 Ethoxylation and propoxylation are moderately fast reactions having a Hatta number near 1 [12]. Therefore, gas-liquid mass transfer and reaction can be considered consecutive processes and the steady state assumption can reasonably be introduced to estimate the eventual mass transfer contribution. However, by using a well-stirred (1500-2000 rpm) laboratory reactor, the mass transfer contribution can normally be neglected. At last, on the basis of the information collected on the reaction mechanisms and of the kinetic results for each hydrophobic starter, we will here report a general approach to the kinetic model for ethoxylation and propoxylation promoted by alkaline catalysts. This kinetic model is able to reproduce, for any mentioned situation, the evolution with time of alkylene oxide, the substrate consumption and oligomer distributions, in well-stirred semi-batch reactors where alkylene oxide is dispersed as bubbles into the liquid phase. This model considers also the increase with time of the volume of the reaction mixture and the change of alkylene oxide solubility in the evolving reaction mixture. Density and solubility parameters have been determined by the authors interpreting experiments performed independently of the kinetic investigations and are reported elsewhere [ 13,14]. 2. REACTION MECHANISM AND KINETIC MODELS Ring opening in the presence of an alkaline catalyst occurs as a consequence of a nucleophilic attack to one of the carbon atoms of the oxirane ring:

/o N RX-

+

CH2

.,-'oN CH

, R'

'

-~

..~ R X C H 2 - - - - C H O

I

--- CH2----- CH

(-)

R'

When R ' = H, the two carbon atoms are identical, otherwise carbon atoms are not equivalent. However, it has been shown that in the last case the attack to the methylenic carbon atom is largely prevalent [13,15]. Therefore, the difference between ethylene and propylene oxide, for example, is that in the first case we always obtain primary alcohols, while in the second we always obtain less active secondary alcohols. Reminding the effect of the ion pairs, we can write the following reaction mechanism for alkoxylation:

RXH + MOH RX-M + + AO

~

, =' RX'M + + H 2 0 I

k0

~ RXAO -M §

=

R X ( A O ) i M + + AO

RX(AO)i+I-M +

Catalyst formation Initiation

Propagation

Kci

RX-M* + RX(AO)iH

RXH + RX(AO)i'M*

Proton transfer

269 corresponding to the most general mechanism for the anionic polymerization of a living polymer. As the alkoxylation reactions occurs with a SN2 mechanism, we can write the following kinetic laws: Initiation ro : k o [ ~ - M

+] ) [ A O ]

(1)

Propagation r~ k~[RX(AO)~. M+][AO] =

i = l , ......

(2) The concentration of charged species appearing in the kinetic laws can be calculated by combining equations of proton transfer equilibria: _ [RXH][RX(AO)~. M +] K~, - [R X - M § ][RX(AO)~H]

(3)

with mass and charge balance equations. By assuming with a reasonable approximation [RXH] + [RX-M +] _=[RXH] [RX(AO),H] + [RX(AO); M +] _--[RX(AO),H] we obtain:

[mC(AO),-M +] = X~ [ ~ - M +][RX(AO)~#] ,

[~H]

(4)

and: [R X - M +] =

[RXH]B~

(5)

[RXH] + s Ke,[RX(AOljH] j=l

where B ~ is the overall catalyst concentration. It is, therefore, possible to calculate the concentration of each ionic pair by introducing the corresponding equilibrium constant Kei and the concentrations of the oligomers. 2.1 A general model to describe a semibatch well-stirred reactor

On the basis of the described kinetic laws the performance of a well-stirred semibatch reactor can be simulated by using the following mathematical model:

270 (i)

Substrate consumption

d[t~(H] dt (ii)

= -r 0

Oligomers formation or consumption

d[RX(AO),H] act

(iii)

(6)

= ri_l - rj

i - 1..... n

(7)

Overall alkoxide consumption

-diVl - ~AO- __ v,s ~

(8)

i=0

To correctly use this model, the dependence of AO solubility and of the liquid phase density on the amount of reacted alkylene oxide must be known. 2.2 Simplified model of a semibatch well-stirred reactor

The described general kinetic model has the possibility to assume that the kinetic constants are different for each reaction step or that proton transfer equilibrium constants change by increasing the number of alkoxide adducts. Our experimental observations are consistent with the reasonable assumption that, in many cases, the chain length does not influence the rate of propagation steps and the values of the proton transfer equilibrium constants [6-8,12]. In the case of a primary fatty alcohol as a starter, for example, only one kinetic and one equilibrium parameter are enough to completely reproduce kinetic runs at a defined temperature and alkoxide pressure, also for what concerns the evolution with time of the oligomers distribution [7]. In the ethoxylation of a secondary alcohol, one equilibrium constant and two kinetic constants are necessary because in the initiation step a primary alcohol is formed being more reactive than the starter [ 13]. In the ethoxylation of nonylphenol, the initiation rate constant is different from the propagation one because of the different acidity of the substrate with respect to the ethoxylated molecule [6]. In the propoxylation of a secondary alcohol (2-octanol) we have the same situation as the ethoxylation of a primary alcohol (k0=kp) because in the initiation step a secondary alcohol is formed which has the same reactivity as the starter; in the propoxylation of a primary alcohol (1-octanol) we have ko ~kp, because in the initiation step a secondary alcohol is formed which is less reactive than the starter [13]. In all the mentioned cases, we need one proton transfer equilibrium constant only, and at least 1 or 2 kinetic constants to simulate the evolution with time of the alkylene oxide consumption and of the oligomers distribution. In the presence of this situation, the mathematical model can be simplified. The equation (5) can be rewritten as:

271 [RX-M +] =

[RXH]B~ [RXH] + K e ~ R X H ~

(9)

where [RXH ~ is the initial substrate concentration. Using equation (9) we can integrate equation (6) obtaining at any time the molar fraction of the unreacted starter: [RXH] [RXH ~

Xo = ~

(lo)

This value can be introduced in the Weibul and Nicander equations relative oligomers distribution [ 16]:

c(i-I) c~ - ~ 1)' XO-- x~ x, - (c -

.;=~

-1)In (j-

to calculate the

(11)

1)!

and the overall alkylene oxide consumption:

nAo = n ~

where:

(12)

1 t

kXoJ

,o

c = Ke kp = Ke ~ .... exp ko ko ~

RT

I = c ~ex

(13)

The system of n differential equations of the general model has been reduced to one differential equation, only, giving the same results. Examples of runs simulations performed in well-mixed isotherm reactors [12] are reported in Fig. 1 for alkylene oxide consumption and in Fig. 2 for the corresponding oligomers distributions. The parameter employed have been derived from literature and are reported in Tab. 1. The same table also shows the parameters for the simulation of other tested with similar results.

272

'~176176176 ~"i~~no, ,.,".~L=~~ ~o~oI ~ jl molar %

T=120"C PAO=2atm KOH : 2% mol

100 | 8O 60

1-octanol AO : EO

It

40 20

2-octanoi ~- T AO : EO

n,

l-oemnolAO = PO 12'

0

!

:.-60'

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

1001[~ 206013040i~18'~

1-octanol 9 AO = PO

2"~

0

,,,

~

lOO 60 40 o

0

50

100

150

200

time (minutes)

Fig. 1 Simulation of AO consumption distributions

0

AO1:34'=EO la_l 1!45' 1

,

~ i,lrn-;nn-,~,

5e 85'

.... 012345 012345 012345 012345 n AO

Fig. 2 Simulation of oligomers obtained for the runs reported in Fig. 1

Table 1 Kinetic parameters determined for both the proposed models (AHe can Starter AO ko~ 108 k. ~ 108 AEo AEp cm3tool~s~ cmLJtool-] s l Kcal/mol Kcal/mol 1-Octanol EO 8.16 8.16 13.0 13.0 1-Octanol PO 71.5 795. 15.6 19.0 2-Octanol EO 223. 8.16 16.8 13.0 2-Octanol PO 795. 795. 19.0 19.0 1-Dodecanol EO 5.95 5.95 13.2 13.2 Nonilphenol EO 1960. 29100. 18.2 19.4 *T>130~

be assumed about 0) co B Kcal/mol 2.0 0 38.9 -3.4 0.08 3.8 2.5 0 4.1 0 7.4* -1.2

273 2.2 Validation of the kinetic models by simulating a spray tower loop pilot plant reactor The use of the described kinetic model has been extended to the simulation of a spray tower reactor with liquid recirculation. These types of reactors are currently used at industrial level to perform both ethoxylation or propoxylation reactions. In this reactor, the sprayed liquid is dispersed in the form of thin liquid drops flying into the ethylene oxide gaseous atmosphere. Normally, drop flight times are long enough to achieve saturation, the rate of mass transfer being very fast. On the contrary, flight times are very short with respect to the ethoxylation rate and the extent of the reaction occurring inside the drops can be neglected. As a consequence mass transfer and reaction occur separately, in two distinct zones of the reactor. In order to describe the mass transfer zone two alternative models have been developed, the first considering the drops stagnant, as suggested by Hall and Agrawall [ 17], and the second considering flying drops inside well stirred. For the case of stagnant drops Crank [18] has given an analytical solution to calculate the concentration profiles inside the drop at different flight times. By elaborating this equation, an average alkoxide concentration can be determined with the following relation:

~Ctrt ~

Ci

C O- C~

1 = 1--~ ~-~i-exp(-4Dn2~r2ts / d2 )

(14)

n=l

where C m is the average concentration in the drops, Ci the initial concentration, Co the equilibrium concentration, D the diffusivity, d the drops diameter, tf the mean flight time. In the case of well mixed drops the evolution with time of the alkoxide concentration is easier to calculate as:

Co-C _ 6kl ~ C o - ~ = expl---ff-tl)

(15)

where C is the concentration in the bulk of drops and kl~ is the mass transfer coefficient that can be calculated using the relation proposed by Srinivasan and Aiken [ 19] In order to do calculations in both cases, we must know the mean diameter of the drops and the averaged flight time. The mean diameters of the drops have been determined experimentally through data obtained whit a laser scattering technique. A mathematical procedure has been developed to evaluate the mean flight time based on the knowledge of the geometrical amplitude of the cone of drops at the spray nozzle, the distances between the cone and the wall of the reactor and the free liquid standing in the reactor, the liquid recirculation feed rate and the pressure drop of the spray nozzle, respectively. The evolution of the reaction, occurring in the liquid phase, performs a plug flow behaviour, from the top to the bottom of the liquid column. Using the developed kinetic model the alkylene oxide concentration and the temperature profiles in the liquid column can be evaluated. Drops falling on the liquid free surface form a layer with a definite alkoxide concentration decreasing from the top to the bottom of the liquid column as a consequence of the reaction. As the reaction is exothermic, temperature increases. To evaluate the internal profiles, we must solve the differential equations:

274

d[AO]_ M+]) dz - A[AO](k~[RX_M+ -~L ]+ kp~-"j[RX( AO)~. dT dz

AH R d[AO] C pp dz

(16) (17)

where A = cross-section of the reactor, FL= recirculating feed rate of liquid, z- distances from the top of the reactor, Cp= specific heat, 0 = liquid density, AHR= heat of reaction. taking account of the fact that the liquid at the bottom of the column will have the same composition fed to the spray nozzle. On the contrary, temperature is corrected to the desired level by a heat exchanger. We observed that the model considering well mixed drops gives better results than that considering drops stagnant in predicting the total amount of alkylene oxide consumed. 3. ACKNOWLEDGEMENTS

Thanks are due to Pressindustria SpA and Scientific Design for the financial support.

4. REFERENCES

1 L.E. St. Pierre in Polyethers - Part I: Polyalkylene Oxides, (N.G. Gaylord ed.), Wiley Interscience, New York, 1963. 2 M.J. Schick (ed.) Nonionic Surfactants, Surfactant Science Series, Vol. 1, Marcel Dekker, New York, 1967. 3 Nico M. Van Os (ed.) , Nonionic Surfactants: Organic Chemistry, Surfactant Science Series, Vol. 72, Marcel Dekker Inc., New York, 1997. 4 E. Santacesaria, P. Iengo, M. Di Serio, Catalytic and Kinetic Effect in Ethoxylation Processes in Annual Surfactant Review, D.R. Karsa (ed.), Vol. 2, Acc. Press, Sheffield, in press 5 R.E. Parker and N.S. Isaacs, Chem. Rev., 59 (1959) 737. 6 E. Santacesaria, M. Di Serio, L. Lisi, D. Gelosa, Ind. Eng. Chem. Res., 29 (1990) 719 7 E. Santacesaria, M. Di Serio, R. Garaffa, G. Addino, Ind. Eng. Chem. Res., 31 (1992) 2413. 8 M. Di Serio, S. Di Martino, E. Santacesaria, Ind. Eng. Chem. Res., 33 (1994) 509 9 C.L. Edwards,, Distribution of polyoxyethylene chains, in <>, (Nico M. Van Os ed.), Marcel Dekker, New York, 1997.. 10 J.A. Orvik, J.Am.Chem.Soc 98,(11) (1976) 3322. 11 E. Santacesaria, M. Di Serio, R. Garaffa, G. Addino, Ind. Eng. Chem. Res., 31 (1992) 2419. 12 J.C Charpentier,. Advances in Chemical Engineering, Vol 11, Academic, New York, 1981. 13 M. Di Serio, G. Vairo, P. Iengo, F. Felippone, E. Santacesaria, Ind. Eng. Chem. Res., 35 (1996) 3848. 14 M. Di Serio, R. Tesser, F. Felippone, E. Santacesaria, Ind. Eng. Chem. Res., 34 (1995) 4092. 15 G. Gee, C.E. Higginson, P. Levesley, K.J. Taylor, J. Chem. Soc., (1959) 1338. 16 H. Weibull, K. Nicander, Acta Chem. Scand., 8 (1954) 847. 17 C.A. Hall, P.K. Agrawall, Can. J. Chem. Eng., 68 (1990) 104. 18 J. Crank, Mathematics of Diffusion, Oxford University Press, Oxford, 1956. 19 V. Srinivasan, R.C. Aiken,, Chem. Eng. Sci., 43, (12)(1988) 3141.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

275

On the Use of Response Reactions in the Kinetic Modeling of Complex Heterogeneous Catalytic Reactions I. Fishtik and R. D a t t a *l Department of Chemical and Biochemical Engineering, The University of Iowa, Iowa City, IA 52242-1219, USA * Department of Chemical Engineering, Worcester Polytechnic Institute, Worcester, MA 01609-2280, USA

I. Introduction

Multiple catalytic reactions usually proceed through an exceedingly complex network of elementary surface steps involving the reactants, intermediates, and products. Thus, reaction mechanisms proposed in the literature are based mostly on shrewd guess work. A knowledge of the actual surface mechanism involved in a catalytic reaction system is, however, crucial to the determination and detailed understanding of its kinetics. It is, therefore, of fundamental importance to have a systematic and rigorous way of determining all of the possible mechanisms for a given catalytic reaction as the starting point. Recently, we have shown [1, 2] that the fundamental equations of chemical thermodynamics have the property of being decomposed into a linear sum of contributions associated with a unique class of reactions referred to as response reactions (RERs). This so far unnoticed feature of chemical thermodynamics has been shown by us to be a powerful tool for a more profound understanding of multiple chemical reaction systems. Here we describe application of the concept of RERs in determining mechanism and kinetics in heterogeneous catalysis. 2. Response Reactions

Consider the general case of a reaction system that contains n distinct chemical species (reactants, intermediates, and products) B1, B2, ..., Bn. Let these species consist of r elements El, E2, ..., Er, where "dements" are understood to be the minimal set of stoichiometrically appropriate atomic, molecular, or ionic species that may be used to "construct" the species, and a~ ( l = 1 , 2 . . . . . r ; i = 1 , 2 . . . . . n) is the number of elements E 1 in the species B i. The stoichiometric relations referred to as the formation reactions are:

1Corresponding Author. E-mail: [email protected]

276

B 1 =

~11E1

B 2 =

0~12E 1 +

B r =

r

Br+ 1 =

B n =

+

1 +

o ~ 2 1 E 1 4- ...

4-arlE

a22E 1 +

+

r

1 +

O~l,r+lE 1 +

alnE

1 +

...

...

+

ctrzE 2

...

(1)

O~rE 2

o~2,r+lE 1 +

a2nE 1 +

2

+

...

+

O~r,r+lE 2

amE 2

From stoichiometric considerations, it is apparent that in a given chemical reaction among species involving r elements, there exists a minimum number of species that must be involved in it; that is, removing any of the chemical species from the reaction will violate the mass balance. Let m be the number of linearly independent reactions among species B l, B2 ,..., Bn. Then, it may be shown, that in the absence of special stoichiometric constraints [3], the minimum number of species that must be involved in a chemical reaction is equal to r + 1. The the reaction that involves r+ 1 species, say the first r+ 1 ones, B1, B2,...,Br+I, may be derived by a special linear combination of the set of r+ 1 formation reactions, eq. (1), for these species so as to eliminate all the dements El, E2, ..., Er, resulting in [4]: all

(X12

...

O~lr

1

0~21

0~21

9

a2r

0 B 1 +

O~r+l, 1 Ofr+l, 2 . . .

ar+l, r 0

0~11

(~12

...

O~lr

0

0~21

(g21

"'"

O~2r

1

O~r+l, 1 a r + l , 2 . . .

B2

ar+l, r 0

(2)

+

...

+

O[11

(3112

"'"

(glr

0

O[21

O[21

...

O[2r

0

ar+l, r

1

a r + l , 1 Ofr+l, 2 . . .

Br+ 1 = 0

which is a so called response reaction (RER) denoted by p(B1,B 2..... B r+1)" The complete set of RERs for a given system is obtained by considering all of the combinations of r+l species from the total of n. The RERs are not all linearly independent. They are, however, unique by virtue of the uniqueness of the formation reactions. The definition of RERs given above is purely stoichiometric. The origin of RERs is, however, in chemical thermodynamics. Thus, the general thermodynamic relations of chemical thermodynamics can be always represented as a unique sum of contributions associated with RERs [ 1, 2].

277 3. RERs and Generation of Elementary Reaction Mechanisms A literature review has shown that virtually all of the overall and elementary intermediate reactions used in the kinetic analyses of various heterogeneous catalyst systems are RERs. Thus, the concept of RERs can be used formally to generate a comprehensive set of elementary reactions and mechanisms. The starting point for this is a list of species, i.e., reactants, intermediates and products. Thus, a mechanism may be derived using a purely stoichiometric algorithm [4]. This algorithm is next exemplified for the C-H-O-S system, where S is an active site on the catalyst surface. A RER for this quaternary system (r = 4) is defined by r+l = 5 species BI=C~ 11 H %2 O. 13 S C~14,B 2 =C~H. 22 O. 23 S~4,B3 -C %1 H. 32 OCt33 8%4 , B 4 = C ~41 H~42 O.43 S 1~44' and C~ H%O%S%. Thus, the equation ofa KER from eq. (2) is given by (3114 1

O~11 O~12 (3113 0114 0

O~21 a22 a23 a24 0

0121 {3f22 0[23 (9[24 1

all

Of 12

0113

Gf31 0[32 0[33 0[34 0 C~H~ 12 O~13 S %4 + 0131 0[32 0[33 O~34 0 C~H,~22 O~23 S (X24 0141 0[42 0[43 0[44 0

0141 0[42 O~43 0[44 0

0~51 0f52 0[53 0[54 0

O[51 0[52 O~53 (:1154 0

0112 (:1113 a14 0

~{3fll (:1112 0113 (:1114 0

a21 0[22 0[23 a24 0

0121 0[22 a23 0[24 0

a31 0[32 0[33 0[34 1 C%,H%~0%3S~3, +

a31 0[32 0[33 0[34 0 C~4,H %0% S %

0~41 a42 0~43 0[44 0

0141 (3[42 a43 (3[44 1

0~51 0[52 a53 0~54 0

O~Sl 0[52 O~S3 a54 0

all

O~ll a12 a13 O[14 0 a21 0[22 0[23 0[24 0 a31 0[32 0[33 0[34 0 0141 (:1142 a43 0[44 0 aS1 a52 0[53 0[54 1

For example, the surface species S, HS, OHS, HCOS and HCOOS define the following RER:

278 00

0 0s+

0 0 0 10 0001~ ~ 1 0 11 ~101 11 1 0 H S + 1111

0 0

1111 1 0 12 1 0

1

1

i 11

0010 oo 1010 1 1 1 1 0 HCOS+ 1 1111 1 1210 1

~ i

OHS

11111 121 lO 10 1 0 HCOOS - 0

10 11

which is HCOS + OHS = HCOOS + 2HS The formalism discussed above was applied to generate mechanisms for a large number of heterogeneous catalytic reaction systems such as the partial oxidation and steam reforming of hydrocarbons and oxygenates, methanol synthesis, Fischer-Tropsch synthesis, oxidative coupling of methane, and NO selective catalytic reduction. All of the elementary steps in mechanisms proposed in the literature under different conditions and catalysts for these reactions were thus generated. Further, a large number of alternative and entirely plausible reactions and mechanisms overlooked by other investigators were additionally generated. 4. D i r e c t M e c h a n i s m s f r o m G e n e r a t e d E l e m e n t a r y R e a c t i o n s

Response reactions are also intimately related to the theory of reaction routes [7], which was originally introduced from purely stoichiometric considerations [5, 6], bur now appears to have a rigorous thermodynamic background [7]. The main purpose of the theory of reaction routes is to determine a unique set of ways (direct mechanisms) [8] to determine how a unique set of overall reactions involving only terminal species (direct overall reactions [9]) can be generated from elementary steps. Thus, the problem is to enumerate possible ways in which the elementary steps in a mechanism can be linearly combined so as to eliminate the intermediates and to arrive at a set of overall reactions. Assigning the symbols s l, s2, ..., Sp to label the elementary steps in a chemical reaction system a mechanism M p is defined as [9] p Mp = E ojsj (3) j=l where a 1, cq ..... % is a set of real numbers called the stoichiometric numbers, which describe the rate of occurrence of the step sj in the overall reaction. In terms of the approach based on RERs, the problem is analyzed as follows. The total number of steps p in a mechanism is equal to or greater than the number of stoichiometrically independent reactions (SIRs) m and, hence, one may always select representative m independent steps from the total o f p based on a variety of considerations. Let these chosen steps be

279

S 1"

vl~B ~ + v12B 2 + ... + v l . B . = 0

S 2"

v 2~B 1 -I" v 2 2 B 2 + -.. + VznB. = 0

(4)

...

v~aB 1 + v ~ B 2 + ... + v.=B. = 0

Sm~

where v~ ( j = l , 2 ..... m; i=1,2 ..... n) are stoichiometric coefficients. These reactions can be always combined so as to eliminate m-1 species, say Bi ,Bi2 . . . . . B i _ .

Then, following a very similar

procedure with that described above, one can define the reaction

n E

]

Vli

Vli

! V2il

V2i

i~1

2 2

...

Vli

" " " V2i

m-i m-l

Vli V2i

B i -

0

(5)

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

]Vmi

Vmi

2

...

Vmi

m-I

Vmi

in which the species B i , B i ..... Bi_~ are not involved (notice, whenever i---i 1, i= i 2. . . . , i - / m - l , two columns of the determinants in eq. (5) are equal and, hence, the stoichiometric coefficients of the species Bi ,Bi2 . . . . . Bi_~ are equal to zero, i.e., these species are eliminated from the reaction). Now, if the species ]3i ,Bi2 . . . . ,]3i _~ are intermediates, then the resulting reaction is nothing but a direct overall reaction and, concomitantly, is a RER [7]. This direct overall RER is produced by the corresponding direct mechanism

M

p

Vl,i I

V l , i 2 - - - 'Vl,i._~

S 1

V2,il

V2,i

$2

. • • V2,i_ l

(6)

. . . . . . . . . . . . . . . V n t i I Vnti 2 . . -

Vnti_l

Sm

Equation (6) may be alternatively written in the form of Eq. (3) with o] = ( - 1 ) m÷, D ~ • j = l , 2 ..... m where Dmj is the minor of the determinant Mp obtained by deleting its j-th row and m-th column. Thus, within the RERs approach, the stoichiometric numbers can be given in an ana/yt/cal form. To derive all of the possible direct mechanisms, one has to consider all of the possible combinations of the m independent elementary steps from the total ofp. It is clear that as far as the RERs are unique, the direct mechanisms are also unique.

280 5. Discrimination Among Mechanisms Based on Reaction Energetics

It is assumed that the following information is available (either from experimental data or theoretical estimations): a) a list of reactants, intermediates and products; b) a set of plausible elementary reactions; c) the enthalpy changes (AH~) and the activation energies for forward (Ejf) and reverse (E~) elementary reactions [10, 11 ]. With this information in hand, it is natural to try to find energetically the most favorable pathways (lowest activation energies of the elementary steps) by which the reactants can be transformed through intermediates into reactants. This may be done, for instance, by plotting and comparing the respective energy diagrams. The procedure is explained next as applied to the mechanism of the water-gas-shift reaction (WGSR) on Cu. From the literature, the following information is available [ 10]

an;

G

s12:

H20 + S = H2OS CO + S = COS H2OS + S = OHS + HS OHS + S = OS + HS COS + OS = CO2S + S COS + OHS = COx S + HS COS + HS = HCOS + S HCOS + HS = H2COS + S COS + OHS = HCOOS + S HCOOS + S = HCOS + OS HCOOS + HS = HECOS + OS HCOOS + S = COx S + HS

-14 -12 24 -5 -17 -22 24 -20 -20 39 19 -2

0 0 26 16 11 0 24 0 0 39 19 26

14 12 2 21 28 22 0 20 20 0 0 2

s13:

C O 2 5 = CO2 + S

5

5

0

Sl: s 2: $3: $4: $5~

$6: $7: $8: $9: s10: Sll:

2HS = H2 + S 8 15 7 The rank of the stoichiometric matrix of this subset of elementary reactions is equal to 10. Thus, any 10 (or fewer) linearly independent elementary reactions can define a direct mechanism. As an example, consider the following set of 10 linearly independent elementary steps {Sl, s2, s3, s4, s7, ss, Sl 1, s12, s13, s14}. There are 10 intermediate surface species that need to be eliminated from this subset of elementary reactions to obtain a mechanism for the WGSR. It is enough, however, to eliminate only 9, since the last one will be eliminated automatically by virtue of the material balance. These 9 intermediate surface species may be chosen arbitrarily without any effect on the overall RER and the respective direct mechanism. Thus, according to eq. (6), the respective direct mechanism is s14:

M = S1 + S2 + $3 + $4 + $7 + S8 - S11 + S12 + S13 + S14 Similarly, the complete set of direct mechanisms for the WGSR may be obtained

281 M1 =s1 + $ 2 + s 3 + s 4 + s 5 + s 1 3 + s 1 4 M2=s1 +s 2+S 3+S 6 +s13+s14 M3=s1 +$2+$3+$9+s12

+S13 +S14

M4 = Sl + $2 + $3 + s4 + $7 - SlO + s12 + s13 + s14 M 5=S 1 +s 2+S 3+S 5-S 7+S 9+S10+S13+s14 M6=s1 +$2+$3+s4+s7+$8-Sll

+s12+s13+s14

From this complete set, one may select the energetically most favorable routes. For the WGSR there are two such mechanisms:

a)

AH~' E~f

o]

H20 + S = H2OS CO + S = COS H2OS + S = OHS + HS COS + OHS - CO2 S + HS

-14 -12 24 -22

CO2S = CO2 + S

5

2HS = H2 + S

8

0 0 26 0 5 15

1 1 1 1 1 1

b)

AH~'

Ejf

o]

H20 + S = H2OS CO + S = COS H2OS + S = OHS + HS COS + OHS - HCOOS + S HCOOS + S = C02 S + HS

-14 -12 24 -20 -2

0 0 26 0 2

1 1 1 1 1

CO2S = CO2 + S

5

5

1

2HS = H2 + S

8

15

1

Net:

Net:

H 2 0 + CO = CO2 + H2

H 2 0 + CO - CO2 + H2

All of the other possible routes that lead to the WGSR are not expected to be significant due to higher values of the activation energy barriers in some of the elementary steps. The above two routes determined are equivalent to those proposed in [10]. From this analysis it appears that the dissociative adsorption of water is the rate determining step in either mechanism. This conclusion is also supported by a recent microkinetic analysis of the WGSR [12]. Once the plausible mechanisms have been thus identified, one may, of course, derive corresponding rate equations using standard LHHW or quasi-steady state approximation formalisms [ 13].

282 6. Concluding Remarks From the abbreviated analysis presented above, it follows that the concept of uniqueness of chemical reactions in a multiple chemical reactions system may be fruitfully utilized to rationalize several important aspects of heterogeneous catalysis. Thus, the concept of RERs provides a simple and systematic algorithm to generate comprehensive reaction steps and a unique and finite number of direct overall reactions and corresponding mechanisms. When combined with reliable estimations of the energetic characteristics of a set of elementary steps, the RERs approach enables a straightforward derivation of the probable reaction pathways and kinetics. As a result, one arrives at a quite manageable set of overall reactions and mechanisms that can be used for a quantitative kinetic analysis.

References 1. Fishtik, I., Gutman, I. and Nagypal, I.,. J. Chem. Soc. Faraday Trans., 1996, 92, 3625. 2. Fishtik, I., Gutman, I. and Nagypal, I., Z. Naturforsch., 1996, 51a, 1079. 3. Smith, W. R. and Missen, R.W., 1982, Chemical reaction equilibrium analysis: theory and

algorithms, John Willey, New York. 4. Fishtik, I and Datta, R., 1998, submitted 5. Horiuti, J., Ann. N.Y. Acad. Sci., 1973, 213, 5. 6. Temkin, M.I., Adv. Catal., 1979, 26, 173. 7. Fishtik, I and Datta, R, 1998, submitted 8. Milner, P.C.J. Electrochem. Soc. 1964, 111, 228. 9. Happel, J. and Sellers, P.H., Adv. Catal., 1989, 32, 273. 10. Shustorovich, E and Sellers, H. Surf. Sci. Reports, 1998, 31, 1. 11. Bell, T. A. In Metal-Surface Reaction Energetics Shustorovich, E., Ed.;VCH Publishers, Inc. New York, 1991: p. 191. 12. Ovesen, C. V., Stoltze, P., Norskov, J. K., Campbell, C.T.J. Catal. 1994, 146,1. 13. Froment, G. F., and Bischoff, K. B., 1990, Chemical Reactor Analysis and Design, Willey, New York.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

283

K i n e t i c E f f e c t s of C h e m i c a l M o d i f i c a t i o n s of PMo12 C a t a l y s t s for t h e S e l e c t i v e O x i d a t i o n of I s o b u t a n e M. Sultan, S. Paul and D. Vanhove ~ Laboratoire de G~nie Chimique et d'Automatique, Ecole Centrale de Lille et Ecole Nationale Sup~rieure de Chimie de Lille, BP 48, 59651 Villeneuve d'Ascq, France

Abstract A kinetic approach of the screening of HPA-catalysts for the direct transformation of isobutane into methacrylic acid is carried out. PMo12 HPA have been modified by substitution of the countercations by NH4 § and/or Cs § and insertion of V in the Keggin structure. The rate of consumption of isobutane is well represented by a Mars & Van Krevelen model and a single simplified reaction scheme can be proposed for all the catalysts tested. The study of the rate constants obtained underlines quantitatively the role of each modification on the catalytic performance. Hence, V stabilises the catalysts and enhances the selectivity into desired products but the main effects are encountered w h e n NH4 § and Cs* countercations are added to the formulation. NH4* lowers the direct degradation of isobutane whereas Cs § strengthens the activation of isobutane while reducing in parallel the degradation of products.

1. INTRODUCTION

Because of the global abundance of liquefied petroleum gas (LPG), interest in the potential use of propane and butanes as sources of corresponding alkenes or their derivatives is increasing [1,2]. In the last decade much progress has been made and various kinds of catalytic reactions and processes have been proposed, particularly for the selective oxidation of light alkanes with molecular oxygen in gas phase. The most successful process for the oxidation of n-butane has been industrialised [3-5]. In a same approach, isobutane could be used in a near future to produce methyl methacrylate, an important monomer of resins. Industrial production of this methacrylate is traditionally achieved by the acetocyanohydrin process [6-8]. However, this process uses the dangerous hydrogen cyanide and overproduces solid ammonium bisulphate. Recently, alternative methods - the methylation of propionaldehyde and the oxidation of isobutene - have been developed but these processes still have problems using high-price feedstocks and consist in two steps synthesis [6-9]. Therefore, direct synthesis of methacrylic acid via the oxidation of isobutane looks more promising. Obviously, this reaction needs a multifunctional catalyst since the reaction is a multielectron oxidation. Many patents and papers have already been published concerning this reaction [10-15] and the catalysts used are mostly Keggin-type heteropolycompounds (HPA) containing phosphorus as central element and molybdenum as peripheral atom and modified by the addition of v a n a d i u m in the primary structure and of different metal ions in cationic position. At present, however, the achievements reported are not good enough to be industrialised reflecting the high difficulty of the effective activation of isobutane over solid surfaces. Corresponding author- E-mail : [email protected]

284 In most of the works, the catalysts were evaluated by direct comparison of the conversions and selectivities obtained in standard operating conditions (contact time, temperature, partial pressures). This approach makes it difficult to understand the effects of formulation modifications due to the strong dependence of selectivities on isobutane conversions and temperature. Moreover, the important thermal effects observed at high conversions cause a drastic decrease of selectivity and activity [16]. The purpose of this study is, therefore, to trace a new route towards formulation of more active, selective and stable catalysts for the reaction. To achieve this objective, a quantitative evaluation of the effects of each modification in composition on the isobutane activation, the reoxidation of the catalyst but also on the selectivities obtained has been carried out.

2. EXPERIMENTAL 2.1. P r e p a r a t i o n of the catalysts (NH4)3PMo120,0 synthesis has been described in [17]. The desired quantities of ammonium molybdate and phosphoric acid were dissolved in hot water, and then 10ml of conc. HNO 3 were poured into the solution in order to precipitate a yellow compound. Dried solid was calcinated at 350~ for 5h and used as such for the catalytic tests. H4PMol~VO,0 was prepared by a method derived from Courtin works [13]. It consists in the preparation of three solutions: i) Sol.[A] : 0.1 mole of NaVO 3 was dissolved to 500 ml of boiling water, then 0.1 mole of Na2HPO4.2H20 was added and the solution obtained was cooled at room temperature; ii) Sol.[B] : 1.1 mole of NaMoO4.2H20 was dissolved to 500 ml of water at ambient temperature; iii) Sol. [C] : 410 ml of conc. HC1 (37%). Sol.[A] was acidified rapidly by a fraction of Sol. [C]. Then Sol. [B] was added dropwise and finally the remaining of Sol. [C]. A red orange solution was obtained and cooled to ambient temperature. H4PMo~VO40 was extracted by diethyl ether, then a quantity of w a t e r equivalent to half of the volume of the organic phase was mixed to it. After the evaporation of ether, the remaining aqueous solution was placed at 4~ to crystallise. Ammonium and caesium salts of H4PMo~IVO,0 were precipitated from mixing 50ml of 0.2M chloride salt solution with 20 ml of 0.005M H4PMollVO40 solution. The precipitates were washed four times by centrifugation to eliminate the unreacted compounds and then dried at 50~ Mixed salts were prepared by dispersing required ratio of the insoluble salts in 20 ml of water. After stirring, the suspension was dried at 50~ The synthesis of the caesium, ammonium coprecipitate catalyst (catalyst F in Table 1) is described in detail in [18]. All the catalysts tested (Table 1) were dried at 120~ overnight in an oven to evaporate the w a t e r of crystallisation. Moreover, as heteropolyanions are very sensitive to temperature, a thermal pre-treatment has been done at 360~ for 5h under a nitrogen flow in a view to stabilise their catalytic performance.

285 Table 1 List of the catalysts tested . . . . . . . . . . . Catalyst ......Reference ......Preparation . Method H4PMollVO40 A Crystallization (NH4),PMol~VO40 B Precipitation (NH4)3PMo 12~ 40 C Precipitation Cs1.2Ho.35(NH4)2.45PMollVO40 D Mixture Cs175Ho.6(NH4)l.65PM~ E Mixture CSl .~(NH,)~.~PM011VO,~ F Coprecipitation

2.2. A p p a r a t u s

The experimental investigations were conducted in a t u b u l a r fixed-bed reactor described in [18]. In order to have an isothermal catalytic bed, differential conditions of conversion were m a i n t a i n e d and dilution with SiC powder (250 l~m) was used. The fixed bed consists in three 5 cm high layers : the catalytic bed made of 3 ml of catalyst (3.7 g) diluted in SiC (1:1 by volume) was sandwiched between identical pure SiC layers. The reactor was fed with a mixture of isobutane (0.09-0.26 atm), oxygen (0.060.20 atm), w a t e r (0.12 atm) and nitrogen at a total flow rate of 3N1/h. All the experiments were carried out at 340~ and i atm. 2.3. A n a l y s i s o f t h e r e a c t a n t s a n d p r o d u c t s

The concentrations of each component (except water) at the inlet and outlet of the reactor were determined by on-line gas chromatography. Two i n s t a n t a n e o u s mass balances, based on the conversions of the reactants (isobutane and oxygen) and the yields in oxidised products, were calculated (see [18] for detail). 2.4. C a l c u l a t i o n s m e t h o d

The carbon and oxygen balances are usually close to 100%. However, the sum of the selectivities is very sensitive to the fluctuations of C balance especially at low isobutane conversion. This is essentially caused by the lack of accuracy in the estimation of isobutane conversion in differential conditions. Thus, in order to avoid erratic results during the determination of kinetic parameters, the isobutane conversion used for modelling was t a k e n as the sum of the products yields. The reaction rate for each reactant or product can be evaluated by the global balance on the catalytic bed : (p~ - Pi)V~a,. r~ =

t cmRT

-

Fi~

pjVcata F/BuYj

rj - t ~ m R T -

m

m

i = iBu or 0~. j = MACO or MAA

286 3. R E S U L T S A N D D I S C U S S I O N We have recently proposed a more rational method for catalyst screening based on a kinetic study [18]. Actually, the rate of disappearance of isobutane on heteropolyanionic type catalysts has well followed the redox kinetic model of Mars and Van Krevelen (MVK). This model is based on the redox dynamics of the catalyst sites, reduced by reaction with hydrocarbons coming from the gaseous phase and further oxidised by the gaseous oxygen, as follows : iBu + Cata-O

~- Products + C a t a

Cata + 02

,-"

iBu + 0 2

Cata-O

~ Products

The balance on catalytic sites at steady state leads to the following equation" r ~

kr "ko " PiB, " Po2 N s kr " PiB, + k o P o 2

where N s cannot be easily determined. Consequently the values of the e s t i m a t e d kinetic p a r a m e t e r s are the specific rate constants : kr.N s and ko.N s. Since the p a r t i a l pressure of isobutane is very high in comparison to t h a t of i n t e r m e d i a t e s at low isobutane conversions, their reaction rates have not been included in the sites balance in order to avoid the excess of kinetic parameters. Therefore, the corresponding t e r m s for products do not appear in the denominator. This a p p r o x i m a t i o n has been verified a posteriori as being founded and it permits us to write a r a t e equation i n d e p e n d e n t of the products kinetic terms. These kinetic p a r a m e t e r s were d e t e r m i n e d by a non-linear regression method (Marquardt's m e t h o d [19]) based on the sum of the squared differences between e x p e r i m e n t a l and calculated values of outlet isobutane partial pressure expressed by

o.f .- ~(PiBu -/3iBu) 2

where

22.4-tc.

m . k o 9k

i=1

o ~) iBu = P iBu - -

Vcata

r -Po2 " -PiBu

"3.6" (k,.-Pibu - ~ k o "Po 2 )

N s

In order to d e t e r m i n e well-defined rate constants, large ranges of isobutane and oxygen feed concentrations have been used.

287

j

0.30 0.25

. . . . . oA

aB

+D

0.20

~0.15 ~0.10

0.05 0.00 0.00

0.05

0.10

0.15

0.20

0.25

0.30

Experimental partial pressure of iBu (atm)

Figure 1. Calculated vs. experimental values of isobutane partial pressure. Figure 1 shows the excellent agreement obtained between experimental and calculated values of isobutane partial pressure. This means t h a t the MVK model remains applicable for all the HPA tested in this work. Table 2 Redox rate constants Catalyst

Reference

H4PMo11VO40 A (NH4)4PMo11VO40 B (NH4)3PMo~2040 C Cs 1.2Ho.35(NH4)2.45PMo11V040 D C s 1.75Ho.6(NH4)1.65PMO 11VO40 E .......................Cs:,:~(NH,)2.~PMo,:V0,n ................................................F ............

ko.N "10 .3 3.3 5.1 3.3 16.3 21.9 11.0

kr.N ~ "10 .3 2.6 1.3 O.8 3.4 7.6 2.4

ko/kr 1.3 3.9 4.1 4.8 2.9 4.6

Preparation Method C P P M M CoP

W h a t e v e r the catalyst, the results in the Table 2 show t h a t the value of ko.N s is higher t h a n kr.N, the activation of isobutane is therefore always the limiting step. Moreover, a strong influence of the composition modifications is observed. The total substitution of the protons by NH4 § ions increases twice k o.N~ values whereas kr.N ~ is decreased in the same proportion (catalysts A & B). This behaviour can be a t t r i b u t e d to the reduction of the acidity of the HPA. When a V atom is introduced in the Keggin structure of (NHt)4PM01204o (catalysts B & C), the stability of the catalyst is significantly improved. Moreover, the rates of activation of isobutane and reoxidation are identically enhanced (constant ko/kr). Pure Cs3HPMo11VO40 and KtPMo11VO40have proved to be completely inactive for the reaction studied. The partial substitution of NH4 § with Cs § during the coprecipitation of the salt (catalyst F) leads to a further twice increase of both rates of reoxidation and reduction. The mixtures of catalysts (D & E) showed very interesting results. They are actually the most active of the set of catalysts tested. Both values of ko.Ns and kr.N ~ are increased in the same proportion for the catalyst D whereas the best results are obtained for the Cs175Ho.6(NH4)1.65PMo~1040 formulation (E) for which the rate of isobutane activation is more enhanced t h a n the reoxidation one (lower ko/kr). This result shows the i m p o r t a n t role of the Cs ratio in the catalyst. It is proposed t h a t the Cs ratio allows to control the acidity of the HPA. If it's too weak the activation is

288 impossible, on the contrary, if it's too strong the degradation is important. A good balance have then to be found to keep both activity and selectivity. The catalyst E gives a significant increase in MACO and MAA yields under standard conditions (Table 3). In varying operating conditions, this yields reaches up to 5.3%. It has been noted t h a t stability is relatively long to be attained with this kind of catalysts. It has to be underlined at this point that the study of the Table 3 on the only basis of selectivities and conversions would place the catalyst E in a bad position because of the apparent low selectivity in desired products. The total yield is a better feature of its performance. Table 3 Catalysts performance - reaction conditions: isobutane 0.26 atm, O 5 0.13 atm, H~O 0.12 atm, T-340~ tc=3.6S ~P = l atm. Catalyst Ref. ~su Xo~ SMAA S MACO S C O Total (%) (%) (%) (%) (%) Yield (%) A 2.5 2.3 25.0 38.9 26.8 1.6 H,PMo11VO,o B 2.3 11.4 49.4 32.2 12.5 1.9 (NH,),PMox~VO,o C 1.5 7.9 31.9 43.8 18.4 0.7 (NH,)3PMo~.O,o D 6.1 36.3 45.6 14.8 29.7 3.7 Cs1.~Ho.35(NH,)~.,sPMol~VO,o E 10.3 72.4 32.1 8.1 45.5 4.1 Cs~ ~sH06(NH,)~6~PM~ ,0 Cs, .~(NH,)2.~PMo,1VO,o F 4.1 26.0 44.6 23.3 ......25,6 . 2.8 ............ Furthermore, the following simplified reaction scheme is proposed for studying the effects of catalyst's composition on the various steps of the reaction: KI

Isobutan!

Methacrylic Acid + Methacrolein ~K3 ~ DegradationProducts

The rate of product formation can be derived from this scheme as 9 dP iBu

k ok r -ffiBuff o,

dt

k o P o~ + k r P iBu

~ m

dPAMACO dt

NS

klkoPiBuPo2 - k3koPo, PAuaCO koPo2 + krP~u

N s with k 1 = Kl.k~ ,k 3 = IZ~.k~and Kx+ K~ =1

After integration and simplification, these equations lead to 9

o

PiBu P AMACO =

(K,-l)

K 3 -1

The kinetic parameters were determined by the above discussed non-linear

289 regression method.

Table 4 Relative rate constants Catalyst H4PMollVO40 (NH4)4PMollVO40 (NH4)3PMo12040 Cs, 2H035(NH4)~45PMo,~VO40 Cs175H06(NH4)~.~sPMo~VO40 .......... Cs,.~(NH,)~:~PMo,,VO,o

Reference A B C D E F

K~

K3

0.75 1.0 0.9 0.73 0.6 ~ 0.76

13.0 19.0 25.0 5.6 7.8 9.0

Preparation Method C P P M M CoP

The results in the Table 4 show t h a t I~ is always much greater t h a n K1, i.e. the desirable products react so fastly compared to isobutane t h a t a good yield could not be obtained at high conversions. As compared to the other HPA tested, the pure ammonium salts give a negligible initial degradation of isobutane (I~ ~ 0) but the further degradation of required products is rapid (K 3 - 20-25). The introduction of a V atom in the Keggin structure has slightly increased the selectivity in valuable products by reducing K 3 (catalysts B & C) but the more significant effects in this way is the presence of Cs in the formulation. All the catalysts containing Cs (D, E, F) have actually much lower K 3 constants t h a n the others. The exact role of Cs is not totally elucidated but it is expected to play a role in the enhancement of the rate of transformation of MACO into MAA avoiding hence the degradation of the intermediate aldehyde.

3. C O N C L U S I O N This kinetic approach of the formulation of HPA catalysts for the selective oxidation of isobutane into MACO and MAA seems to be promising. Actually, it allows to underline the importance of V in the Keggin structure, stabilising the solid and leading to more selective catalysts. Moreover, NH4 § cations presence leads to very selective catalysts at low conversion but their weak acidity prevents a good isobutane activation. In this way their coexistence in the formulation with Cs § cations is very important because it gives more active catalysts leading to good selectivity and therefore to higher yields. The role of Cs § cations is not determined precisely but an effect on the rate of transformation of MACO into MAA is expected.

Abbreviations and Notations MAA : methacrylic acid MACO iBu

: methacrolein 9isobutane

F"i : inlet molar flow rate of reactant i (mol/h) k o : turnover frequency for the reoxidation step (mol/atndsite/h) k r : turnover frequency for the reduction step (mol/atm/site/h) m : weight of the catalyst sample (g)

290 N " concentrations of sites in the catalyst (sites/g) Pi: inlet partial pressure for component i (atm) Pi " outlet partial pressure for component i (atm) Pi" mean partial pressure in component inside the reactor (atm) /~i,/~j" calculated outlet partial pressure for reactants or products (atm) o

ri, rj" mean reaction rate for reactant i or product j (mol/h/g) S 9selectivity for products t c 9contact time (s) Vcata" volume of catalyst (ml) X~" conversion of reactant i Y 9yield in products

References

1. Y. Moro-oka and W.Ueda, Catalysis, Vo1.11 (Royal Society of Chemistry, London, 1994), 223. 2. F. Cavani and F. Trifiro, Catalysis, Vol.ll (Royal Society of Chemistry, London, 1994), 247. 3. F. Cavani and F. Trifiro, Chemtech 24, (1994); G. Busca, G. Centi, F. Trifiro and V. Lorenzell, J. Phys. Chem. 90 (1986) 1337. 4. G. Centi (Ed.), Vanadyl Pyrophosphate Catalysts, Catal. Today, 16 (1993). 5. The Chemical Engineer, 13 sept. (1990) 3. 6. H. H. Kung, Adv. Catal. 40, 1 (1994); D. Artz, Catal. Today 18, 173 (1993). 7. R. A. Sheldon, Dioxygen Activation and homogeneous catalytic oxidation, 573, Elsevier, Amsterdam (1991). 8. M. Misono and N. Nojiri, Appl. Catal., 64 (1990). 9. G. Centi, Catal. Lett. 22,53 (1993). 10. H. Krieger and L. S. Krich, US Patent 4, 260, 822 (1981), assigned to Rohm & Hass Company. 11. S. Yamamatasu and T. Yamaguchi, Jpn. Patent 02-042,032 (1990) assigned to Asahi Chemical Industries Company. 12. T. Kuroda and M. Okita, Jpn. Patent 04-128,247 (1991), assigned to Mitsubishi Rayon Company. 13. F. Cavani, E. Etienne, M. Favaro, A. Galli and F. Trifiro, Catal. Lett. 32 (1995) 215. 14. N. Mizuno, M. Tasaki and M. Iwamoto, Appl.Catal. A 128 (1994) L1. 15. N. Mizuno, M. Tasaki and M. Iwamoto, J. Catal. 163 (1996) 87. 16. D. Vanhove, Appl. Catal. A, 138 (1996) 215-234. 17. P. Courtin, Rev. Chem. Min., t8 (1971) 75. 18. S. Paul, V. Le Courtois and D. Vanhove, Ind. & Eng. Chem. Res., 36 (8) , 1997, 3391-3399. 19. D. W. Marquart, Soc. Ind. Appl. Math. J.,11 (1963) 431.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

291

Kinetic B a s e d D e a c t i v a t i o n M o d e l l i n g o f an Isothermal Propane Dehydrogenation Reactor E Hugh Stitt, S David Jackson & Frank King Synetix, Research Technology and Engineering Group, PO Box 1, Billingham, Cleveland TS23 1LB, United Kingdom

Abstract

Deactivation, due primarily to the deposition of carbonaceous species onto the surface is well known in the catalytic dehydrogenation of light alkanes. Catalyst regeneration is by burning off the "coke". This paper describes the experimental measurement and modelling of the kinetics of deactivation of a propane dehydrogenation catalyst. Stable isotope tracer experiments [8] showed that it is possible to neglect the adsorption terms. Thus a simplified kinetic model is derived based on a set of pseudo-reactions, and the changes to the kinetic parameters with deactivation studied. A simple, intrinsic deactivation model is derived and fitted based on the accumulated carbonaceous species on the catalyst surface. Finally, using a dynamic reactor simulation it is shown that the model can be successfully used to predict the catalyst activity change and performance of a given isothermal reactor.

1.

INTRODUCTION

The deactivation of catalysts during the dehydrogenation of light alkanes due to the deposition of carbonaceous species onto the catalyst surface is well known. Regeneration of the catalyst is periodic and by oxidation. The designs of commercial reactors for this process take different approaches both to managing the deactivation and to the regeneration. Examples of actual and proposed schemes include: 9Lummus / Houdry "Catofin" process [ 1] Heat of regeneration is fuel supplemented and used to regeneratively heat the endothermic dehydrogenation reaction by cycling the feeds through a series of 3-4 fixed, adiabatic beds, over a 15-25 minute period. 9UOP Oleflex process [2,3] Deactivation suppression by hydrogen co-feed. Heating is by fired inter-heaters between three adiabatic moving beds. Catalyst cycle times are 3-5 days 9Phillips STAR process [4] - where deactivation is suppressed by a high molar ratio co-feed of steam in conjunction with a Sn promoted Pt catalyst in fired tubular reactor. Cycle times are in the order of 6-8 hours 9BASF/Linde Process[5] - As above, a fired tubular reactor is used, but here deactivation suppression is by hydrogen partial pressure. On-stream times are 6-8 hours 9Snamprogetti / Yarsintez Process [6] - this utilises twin fluidised beds, one each on process and regeneration duty with catalyst cycling between them. Catalyst cycle times are not clear but must be relatively short as regenerative heat transfer is claimed.

292 Operation is generally at 550 - 650~ and 0.3 - 3 atm. Each of these processes operates with a changing catalyst activity and with different temperature profiles. The understanding of deactivation, the effects of temperature and cycle averaged effects is thus vital to the design or selection of the optimum configuration. This paper describes the derivation of an intrinsic kinetic and deactivation model for a chromia catalyst. This model was then used to determine the optimum design for a dehydrogenation reactor [7].

2.

PROCESS CHEMISTRY AND CATALYSIS

2.1. Reactions 1) C3H8 r

C3H6 + H2

The propane dehydrogenation reaction is strongly endothermic (111 kJ/mol) and is equilibrium limited; equilibrium conversion at 600~ and 1 atm is 54.6 %, Fig. 1. At such elevated temperatures alternate reactions occur, notably cracking to lower hydrocarbons (methane, ethane, ethene), the formation of "coke" (viz. polyaromatics / carbon deposits on the catalyst surface), and the dehydrogenation of propene. Experimental work with stable isotope tracers indicates that adsorption terms are not necessary in the macroscopic modelling of the overall process [8]. Thus, the competing reactions can be represented using a set of pseudo reactions: 2) Calls ~ 3C + 4H2 3) C3H8 ::~ C + 2CH4

4) C3H8 ::> CH4 + C2H4 5) C2H4 + H2 <::>C2H6 The approximation of coke by carbon is questionable but allows a major simplification in kinetics and the eventual model. The model neglects formation of coke from ethane, ethene and propene, an assumption supported by experimentation with these components over the catalyst at the highest anticipated temperatures. The dehydrogenation of propene to propyne and propadiene was also neglected. The products were not detected in experiments and equilibrium conversion of propene at the operating temperatures is known to be low, Fig. 1.

2.2. Catalyst & Experimental The catalyst used in this study was a modified alumina supported chromia, prepared by an impregnation route, involving drying and calcination. For a more detailed discussion of the catalyst and experimental methods refer to Jackson et al [8]. Two isothermal reactor systems were used in this study. Pulsed reaction studies were performed in a dynamic mode using a pulse-flow microreactor system with on-line GC. Continuous flow reaction studies were performed in a 1 atm microreactor with the gas stream exit the reactor being sampled by on-line GC. The catalyst was reduced by heating to 600~ in a stream of hydrogen, and then the flow was switched to propane.

3.

EXPERIMENTAL DATA AND KINETIC MODELLING

Experimental data for the reaction product distribution, measured using continuous flow experiments, are given in Fig.2, showing variations of performance as a function of

293

temperature and space velocity. Simple rate equations, first order in the alkane were assumed for all of the reactions. Where applicable, equilibrium constants were calculated from first principles (Gibbs free energies). Evaluation of rate constants was by fitting against experimental data. Data from pulsed experiments at high GHSV were preferred as this approach minimises the effects of catalyst deactivation. These data were plotted using an Arrhenius type relationship to obtain the activation energy and the pre-exponential constant, shown for the main reaction in Fig.3.

4.

CATALYST DEACTIVATION AND MODELLING

Experimental data, such as those shown in Fig.4, demonstrate significant catalyst deactivation over the experimental period (< 1 hour). It is also evident that selectivity improves with time on stream, or as deactivation increases. Activity is restored by coke "bum-off' followed by reactivation (reduction with hydrogen). It will be a requirement to process the catalyst through a dehydrogenation - regeneration - reactivation cycle. Process and reactor design will require dynamic modelling of the catalyst through deactivation. Regeneration and reactivation were also modelled, although this aspect of the work is not presented herein. A simple non-mechanistic model was required to allow reactor modelling

294

and design with a predictive capability. Deactivation must be related to the surface effects; the accumulated coke deposits on the catalyst. This will enable dynamic modelling of the reactor. It has been recognised for over 30 years that the deactivation due to the deposition of carbonaceous species can be modelled using an expression of the form [9] : d~= exp(-a.Cc)

(1)

where d~ is the activity relative to flesh catalyst, Cc is the coke concentration and c~ is a constant. This approach has been successfully applied to dehydrogenation of butene-1 [ 10]. A number of detailed mechanistic studies have followed on from this utilising more sophisticated approaches, but confirming the validity of the general approach, see for example Acharya & Hughes [ 11] or various papers presented at the 7th International Symposium on Catalyst Deactivation [12]. Some have attempted to interpret coke profiles within catalyst pellets. The studies with butene- 1 [ 10, 11 ] do not consider the formation of lower hydrocarbons but simply the selective dehydrogenation reaction and formation of coke from the reactant and product. Both reports note that the value of a is constant; independent of the reaction as well as of temperature and partial pressure. It is specifically concluded from this that the alternate reactions occur on the same active sites. The objective of the present study was however to to take advantage of the conclusions from the stable isotope tracer work, and develop a semi-empirical model for deactivation that could be used predictively to consider optimum designs for a reactor; and one that could be derived with the minimum of experimental data. Learning from previous work was therefore taken back to the bare essentials and Eqn. 1 was utilised in its glorious simplicity, with the deactivation parameter (c~) being fitted empirically. The kinetic equation set described above allows local rates of carbon deposition to be calculated and, through combination with the deactivation equation based on Cc the accumulation of surface deposits and their effect on activity is intrinsically included within a predictive dynamic reactor model.

295

4.1. Deactivation Model Description The rate of reaction i at time t is given by: R it = t~.R io

(2)

where R~0 is the initial rate (fresh catalyst) and ~ is an activity factor given by Eqn 1. The deactivation parameter, tx was, in contrast to previous work [9, 10], found not to be constant but rather, a function of temperature and different for each reaction, discussed further below.

4.2. Deactivation Model Fitting Rearranging and integrating the deactivation model equations (1) and (2) gives: 1 / R , = 1~Rio + c~.t

(3)

Hence tx and R~0 can be evaluated for each reaction by the relatively simple plot of 1/R vs. time. A sample data fitting plot is given in Fig.5.

The parameter ct, determined from plots such as shown in Fig.5, was found to be a function of temperature. The best fit of tx as a funtion of temperature for a given "reaction i" was obtained using an expression of the form: cti = Qi. e x p ( q , / R T )

(4)

where Qi and qi are constants, and qi is positive. This is shown in Fig. 6 for the main carbon formation reaction. The values of constants Q~ and qi, and thus a , are different for each reaction. This indicates that different types of site are responsible for the dehydrogenation and non-selective reactions and that they deactivate preferentially. This is in contrast to previous studies with butene-1 [ 10, 11 ], but is consistent with the observation in this study

~ t h s ~ ~ F ~ F a t l # M a h C l r b o n ~ ~ ~

+

3

C

%

h

m

*

4

~

Madd

Coko (F@ 7a: Accaunukd )

-6:

~ b t o f ~ v k h k d d d ~

m-mob: m a p d w r d i f R m n t ~ ~ ~ /

--

;

ltmiFmdT-(lflc) (ThnmlE3) mPrdPr

C

*is'

3

AccmmkdC~)*bqnd.V.bw MO ssO600658

p-a.8: ~ c e c t i o a ~ a s M o n o f - - Ti~l#IOC:CB1Sv~IIr

wlAao'd CllsbDD W F - 1

51

Fig.9: ~ 0 1 1 R s b s s r ~ F ~ o f ~ F i i l O : P m d u d ~ V . l u c r

C

o

o

10

mA%b

R-pp+H2

1

r

TL-B#IOC:

n

~

20 30 40 ~ i m on e Shim (mh)

RilblAxb

R&X+a

O

f

SO

iioO

O H S V ~ ~ l h f

-* h*C+2k

UObl-

*

~ t hCt pdw b w of ~ k p a~h e d d Model ""

:OPLPV4OWlwST~~C

N8dalV.LUr(d* 3oJar

m

l2mh

*

21&

*

Mlnhl

297 that selectivity increases with deactivation; presumed due to the active sites that promote the cracking reactions becoming fouled. As noted above, previous workers [9-11 ] have found to be constant with respect to temperature. That this is not so in this study is attributed to the simplifications and non-mechanistic nature of the kinetic equations used.

5.

DEACTIVATION MODEL VALIDATION

The combined kinetic and deactivation model was used for simulations in the form of a FORTRAN dynamic model based on sequential steady state solutions. This was used in isothermal mode to simulate the experiments and thus obtain model validation. A comparison of predicted and measured accumulated carbon deposits given in Fig.7. This shows a good correlation and confirms the validity of the modelling approach. The coke deposition rate has thus been fitted for the experimental range, and then the rate equations fitted, to give an inherent rate of coke deposition, see Fig 8. The model fit, shown in Figs 8-10 is good for the main reaction, although scatter is evident on the side reactions involving the production of ethane and ethene. This is attributed to inaccuracies in determination of the kinetic parameters caused by the low concentrations of the C2 species. These simulation runs yielded also reactor profiles as a function of time. Examples of results from a single simulation are given in Figs. 11. These are consistent with observations in the literature [ 10, 11 ]. Fig. 11a shows that the coke laydown is strongest at the front of the bed, where the propane concentration is highest, and that the coke concentration profile retains this gradient throughput the run, Fig. 11b. The overall effect of this on propane conversion rate, Fig. 11 c, is seen mainly at the front of the bed where the deactivation leads to significant loss of rate. The rate at the end of the bed is little affected; a balance of the increasing in propane concentration due to the deactivation of the upstream part of the bed and the loss of activity in the rear. The reaction selectivity, Fig. 11 d, increases with time as non-selective sites become fouled, but declines through the length of the reactor for the duration of the run. The former has previously been discussed, and the latter is due to the relative decrease in the dehydrogenation rate caused by the effect of the equilibrium constraint. The model was later used in a non-isothermal, non-adiabatic design program to consider reactor design [7]. 6.

REFERENCES

1 RG Craig & DC Spence, Catalytic Dehydrogenation of LPG by the Houdry Catofin and Catadiene Processes, in RA Meyers, Handbook of Petroleum Refining Processes, McGraw Hill (1986) 2 BV Vora, PR Pujado & RF Andesrson: Oleflex: C2-C5 Dehydrogenation Updated, Energy Progress, 6(3), (1986), 171 - 176. 3 FP Wilcher, CP Luebeke & PR Pujado: Productions of Light Olefinsfrom LPG, Hydrocarbon Technology International, (1992), 93-102. 4 RO Dunn & RL Anderson : "STAR"- The Phillips Steam Active Reforming Process for Light Paraffin Dehydrogenation, AIChE Summer Meeting, San Diego CA, (Aug 1990) 5 H Bolt & H Zimmerman : Dehydrogenation Process for Propane and Isobutane, Hydrocarbon Technology International, (1992) 149-151.

298 Fig. 11 9 Results from Simulation of Isothermal Reactor with Catalyst Deactivation Temperature 600 C, GHSV 4031 per hr Fig. 11 a 9 Local Coke Deposition Rate

Fig. 1 l b 9 Build up of Carbon Deposit

~8 L)

2.5

....., O

.~ 2

~7 "-6 ~,6

"~1.5 5 n~Q

40.5

_~3. r~

r,.) 0.2

0.4

0.6

m

m

0.8

3 mins

m

m

m

m

m

m

n

m

m

m

Normalised Length

Normalised Length Start

i

0

12 mins

30 mins

21 mins

iiiiiii

Fig. 1 l c

Local Propane Reaction Rate

Fig. 11 d 9 Local Reaction Selectivity

120

~ 96

100

~ 95

80

~ 94

Z"~.'-~-.~. ........ m

~

~

~

~

i~mm~ ~

o

60

....... ~.2:.._~ :.-~ ........ . . . - . . . ~ . . . ~

40 20

6 7 8

93 .o 92

o

o14

o16

NormalisedLength

91 0

'

012

'014

'

016

'018

'

1

Normalised Length

F Buonomo, G Donati, G Fusco, F Galimberti, I Miracca & L Piovesan : Fluid Bed Dehydrogenation of Mixed Paraffin Feedstock, AIChE Ann. Meeting, Chicago, (1996). EH Stitt, SD Jackson, F King & DG Shipley : "Modellingfor Design of a Deactivating Non-Isothermal Propane Dehydrogenation Reactor", to be published. SD Jackson, J Grenfell, IM Matheson & G Webb : Modelling of Alkane Dehydrogenation

Under Transient and Steady State Conditions over a Chromia Catalyst Using Isotopic Labelling, Int Symp Reaction Kinetics and the Development of Catalytic Processes, Brugge, 19-21 April 1999. 9 GF Froment & KB Bischoff: Non-Steady State Behaviour of Fixed Bed Catalytic Reactors Due to Catalyst Fouling, Chem.Eng.Sci., 16, 189-201 (1961) 10 FJ Dumez & GF Froment : Dehydrogenation ofl-Butene into Butadiene. Kinetics, Coking, and Reactor Design, Ind.Eng.Chem. Process Des.Dev, 15, (1976) 291-301. 11 DR Acharya & R Hughes : Modelling of Butene-1 Dehydrogenation in a Fixed Bed Reactor - Bed and Pellet Profiles, Can.J.Chem.Eng., 68(2), (1990) 89-95. 12 CH Bartholomew & GA Fuentes (Eds): Catalyst Deactivation 1997, Studies in Surface Science and Catalysis, Vol 111, Elsevier, 1997.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

299

Development of kinetic models for reactions of light hydrocarbons over ZSM-5 catalysts. Experimental studies and kinetic modelling of ethene transformation and deactivation of HZSM-5 catalyst D. B. Lukyanov Centre for Microporous Materials, Chemical Engineering Department, UMIST, PO Box 88, Manchester M60 1QD, United Kingdom

Abstract Ethene aromatization reaction was studied at 400~ over HZSM-5 zeolite at different space velocities and time-on-stream (TOS). Consequently, the experimental d a t a were obtained corresponding to the reaction proceeding over fresh (TOS = 5 min) or deactivating (TOS up to 4 hours) catalyst. The reaction studies were supported by investigation of coke formation by using in situ EPR method. As a result of this work, a detailed kinetic model for ethene aromatization reaction over HZSM-5 zeolite under conditions of catalyst deactivation was developed.

1. I N T R O D U C T I O N Conversions of various feedstocks over ZSM-5-based catalysts can be represented schematically [1,2] as a sequence of three main stages: (i) conversion of the feed molecules into olefins; (ii) olefin interconversion; and (iii) aromatization of olefins. Every stage is a complex catalytic reaction involving numerous chemical transformations of reagents. Additionally, catalytic processes are complicated by coke formation and subsequent catalyst deactivation. In a number of our previous papers [3-7] an approach to the construction of a general kinetic model for aromatization of different feedstocks over ZSM-5 catalysts was developed. This approach was based on splitting of the full kinetic model into four separate models: (1) Feed conversion into olefins (e.g., paraffin cracking, methanol-to-olefin reaction); (2) Olefin interconversion; (3) Aromatization of olefins; (4) Coking and catalyst deactivation. Such a splitting allowed to move from relatively simple kinetic model for olefin interconversion over HZSM-5 zeolite to much more complex kinetic models for light olefin and paraffin aromatization reactions over HZSM-5 and GaZSM-5 catalysts [5-7]. The latter models have provided a detailed description of the aromatization processes over fresh catalysts, but have not described coke formation and catalyst deactivation. Therefore, the present investigation was undertaken in order to obtain experimental information on coke formation in the course of ethene conversion over HZSM-5 zeolite and to

300 use this information for construction of a kinetic model for ethene conversion under conditions of catalyst deactivation.

2. EXPERIMENTAL Kinetic investigation of ethene conversion over HZSM-5 zeolite ( S I O 2 / A 1 2 0 3 = 240) was carried out in a quartz flow reactor at 400~ and ethene concentration of 40 mol.% in N2. The initial catalytic activity and product selectivities were characterised by the data obtained at catalyst time-on-stream of 5 min. Formation of coke and its effect on the catalyst activity were studied in situ by using EPR method with on-line GC analysis of the reaction mixture. In these experiments ethene concentration was varied between 5 and 40 mol.% in Nz, and the reaction was carried out in the reactor of small volume (-3 ml) placed in the resonance cell of an EPR spectrometer (Rubin). For recording EPR spectra of the deactivated catalysts at room temperature a Bruker ER 200E spectrometer was used.

3. RESULTS AND DISCUSSION 3.1. Coke formation

At the beginning of this investigation, catalyst samples deactivated in the course of ethene transformation were studied by EPR method at room temperature. EPR spectra of the coked HZSM-5 samples showed a singlet line with a g-factor of 2.0025 and a width (AH) of 8-10 G. The intensity of the signal did not change upon adsorption of oxygen. This indicated that the coke formed in the catalyst had no delocalized electrons which could interact with adsorbed oxygen, i.e., coke with a pseudographite structure was not generated. Similar EPR spectra were observed [8] with HZSM-5 samples containing small amount of coke (-1 wt%). With increasing coke content in the samples, the spectrum changed to a superposition of two signals, one broad (AH = 6-8 G) and one narrow (AH = 1-2 G) [8]. The intensity of the narrow signal decreased upon oxygen adsorption as a result of its broadening. This result was explained [8] by the presence of two types of coke in the zeolite: multi-nuclear condensed compounds on the outer surface of the zeolite (the narrow signal) and large nondesorbing molecules (monoaromatic compounds with branched side-chains) in the zeolite channels (the broad signal). In the present study the narrow signal was not observed and, therefore, it was reasonable to assume that the coke was formed in the channels of the HZSM-5 catalyst. To develop the kinetic model for ethene conversion under conditions of catalyst deactivation, quantitative data on the rate of coke formation in the course of the reaction are necessary. Such data were obtained by using in situ EPR method. This investigation was performed at 400~ with the mixtures of ethene (5-40 mol.%) in N2, and the reaction was carried out in a differential mode. Some of the results obtained are shown in Figure 1, which indicates that two processes with different characteristic time occur during coke formation. The first process is observed during the initial period of the catalyst operation (10-20 min). In our opinion, this process corresponds to the formation of coke nuclei, while the second process, observed during several hours, corresponds to their growth. The latter process of

301 80 t 960 I

c

Ethene concentration 9 5 mol.% 9 21 mol.% 9 40 mol.%

a) '- 40 w,

2r E E

20

0

50

100

150

200

250

300

Time, min

Figure 1. Effect of time-on-stream and ethene concentration on the intensity of the EPR signal of the paramagnetic coke formed in the HZSM-5 catalyst during ethene conversion at 400~

80

o~ ~

E 60 0

9 WHSV=6.2 h -1 9 WHSV=8.8 h -1

E

> 40 e~ 0 0 eI~ 2 0 eW

0

50

100

150

200

250

Time, min Figure 2. Deactivation of the HZSM-5 catalyst in the course of ethene transformation at different weight hour space velocities (WHSV). Symbols = experimental data; Curves = kinetic modelling results.

302 coke formation is apparently responsible for catalyst deactivation (see Figure 2) and, therefore, its kinetic description is of particular interest. Based on the results obtained, it was established that the rate of growth of paramagnetic coke (ro) at a constant initial ethene concentration could be described by the following equation: (1)

ro : kcKa[C2:]ot/(1 + Ka[C2:]) = kr

where ko and Ka are constants, [C2:] is the mole concentration of ethene in the gas phase, [C2:Z] is the concentration of adsorbed ethene, and ot is the fraction of nondeactivated acid sites responsible for ethene adsorption and coke formation. The value of ot was determined from the equation c~ = (Ir '*,- Ir162*'

(2)

where L is the intensity of the EPR signal at any given time, and Io" is the intensity of the signal at t --> oo. To establish a relationship between the paramagnetic coke and all coke formed in the zeolite, we collected data on carbon dioxide which was formed during catalyst regeneration. These data are included in Table 1 which shows that the ratio between the number of carbon dioxide molecules (Nco2) and the number of paramagnetic sites (Np.s.) remains approximately constant for samples with different coke content. This result indicates that the intensity of the EPR signal can be used to characterise the amount of coke formed under conditions of ethene transformation. Table 1

Nco2/Np.s. ratio as function of the intensity of the EPR signal of the paramagnetic coke formed on HZSM-5 catalyst during ethene conversion at 400~ Ir rel. units

Nco2/Np.s. 10.3

23.5

32.7

47.4

58.0

1.5

1.3

1.7

1.6

3.2. Kinetic model for ethene transformation over H Z S M - 5 zeolite

Based on the results obtained, a kinetic model for ethene conversion under conditions of HZSM-5 catalyst deactivation was constructed. This was done by insertion of deactivation function (~) in the rate equations of the kinetic model for ethene aromatization over fresh HZSM-5 catalysts [5]. The procedure and principles of the kinetic model formulation were discussed in detail previously [3,5], and, as a consequence of this, are not considered in this paper. Brief description of the model is given below. To decrease the number of the reaction species in the system, all isomers of the same carbon number were lumped into a single component. In this way, 42 components, which represent all reaction species involved in ethene aromatization reaction, were formed: 9 olefins Cf-C10=, 10 paraffins C1-C10, hydrogen, 7 dienes D4-D10, 5 alkylcyclohexenes X6-X10, 5

303 alkylcyclohexadienes Y6-Y10, and 5 alkylbenzenes A6-A~0 (subscript index denotes the number of carbon atoms in the component molecule). The reaction scheme used for the development of the kinetic model for ethene aromatization reaction includes 31 steps of hydrocarbon adsorption and 228 steps of hydrocarbon transformation on zeolite catalytic sites (Z). All adsorption and reaction steps, which were used previously in the kinetic model for light olefin aromatization reaction over fresh catalysts [5], are shown below: 1. Hydrocarbon adsorption on zeolite acid sites (31 steps): ]X'al

C,= + Z +-) C,=Z

(2__n_< 10)

(3)

(4 < n < 10)

(4)

(6 _
(5)

(6 _
(6)

(6 _ n _ 10)

(7)

/k'a2

D, + Z ~-~ D,Z /~'a3

X, + Z ++ X,Z J~'a4

Y, + Z ~-~ Y,Z ~-'a5

A, + Z ~ A,Z

2. Olefin oligomerization and cracking on zeolite acid sites (16 steps): koL(n,m) C.= nt- Cm=Z ~-~ C.+m=Z kcR(n,m)

(2_
(8)

3. Diene formation via hydrogen transfer between two olefin molecules on zeolite acid sites (63 steps) kDF(m)

Cn=+Cm:Z ----) DnZ+Cm

( 4 < n < 1 0 ; 2 < m < 10)

(9)

4. Interaction of olefins with dienes on zeolite acid sites to form higher molecular weight dienes (15 steps): kDOL

Cn--+DmZ ~

Dn+mZ

(2
10)

(10)

kDCR

5. Diene cyclization on zeolite acid sites (5 steps) k,c(n) D,Z --+ X,Z

(6_
(11)

304 6. Aromatics formation via hydrogen transfer between olefins and cyclic olefins (or cyclic diolefins) on zeolite acid sites (90 steps): k~(m) X,+Cm:Z ---> Y,Z+Cm

( 6 < n < 1 0 ; 2 < m < 10)

(12)

~(m) Yn+Cm=Z --> AnZqt-Cm

(6
(13)

7. Cracking of C5+ paraffins on zeolite acid sites (39 steps): kcH(n)

C,+Z Cn -~-Z

--> C,:Z+H2 kcc(n,m) ~ Cm:Z-[--Cn.m

(5_
(14) (15)

Formation of coke and deactivation of the catalyst are described in the kinetic model developed in this work by the following equations: C,:Z --> coke;

dc~/dt = -koaY~[C,:Z]

(16)

The equations for the rates of the reaction steps were derived on the basis of the mass action law [9] (details are given in Ref. 5), and the deactivation function was introduced by multiplying these rates by ~. Thus, the kinetic model for ethene aromatization under conditions of catalyst deactivation consisted of a set of 42 equations describing transformation of 42 components in 228 reaction steps plus equation (16) describing deactivation of the zeolite catalytic sites. For kinetic modelling the plug flow reactor model was used, and for numerical integration of the system of differential equations special computer programme was developed. Estimation of the rate constants was done step by step, as reported previously [5]. First, the rate constants were estimated for the reaction over fresh catalyst, and the experimental data, obtained at time-on-stream of 5 min, were described. During the second stage, calculations were performed for different time-on-stream and some of the results obtained are shown in Figures 2, 3 and 4. Comparison of these kinetic modelling results with the experimental data shows that the kinetic model constructed in the present work, describes quantitatively ethene transformation over HZSM-5 zeolite under conditions of the catalyst deactivation. Interestingly, but not unexpectedly, Figures 3 and 4 show that both the space velocity and time-on-stream have the same effect on distribution of the products of ethene conversion over HZSM-5 zeolite.

4. ACKNOWLEDGEMENTS I thank Vadim Shtral for development of the programme for computer simulations and Tatiana Vajnova for her assistance in experimental work and preparation of this paper. Also, I

305

o~

5

Time

~4

C4 =

A

WHSV

C 3-

E

O :,~ 3 L 4,,.I

E ~2 0 E 0 1

o

0

~ ,IJ~"~

0

,

-

I

10

-

,

I

,

20

C6 = I

,

30

40

Ethene conversion, % Figure 3. Concentrations of the main products of ethene transformation over HZSM-5 catalyst as functions of ethene conversion varied by change in space velocity (open symbols) or during deactivation at WHSV = 8.8 h 1 (solid symbols). Temperature = 400~ Symbols = experimental data; Curves = kinetic modelling results. 20 Time

~

C3

="

C6 =

15

E o

10

L

c o c o

5

/

o

0

C 2 -C a

0

,rom ,cs , 4o

'

6o

'

8'.

|

100

Ethene conversion, % Figure 4. Concentrations of the main products of ethene transformation over HZSM-5 catalyst as functions of ethene conversion varied by change in space velocity (open symbols) or during deactivation at WHSV = 6.2 h -1 (solid symbols). Temperature = 400~ Symbols = experimental data; Curves = kinetic modelling results.

306 gratefully acknowledge helpful discussion with Aleksey Slinkin of the results obtained in the EPR experiments.

5. REFERENCES 1 W.O.Haag, Proc. 6th Intern. Zeolite Conf., Butterworths, Guildford, UK, 1984, 466. 2 P.B.Weisz, Ind. Eng. Chem. Fund., 25 (1986) 53. 3 D.B.Lukyanov, V.I.Shtral, V.I.Timoshenko, S.N.Khadzhiev, Proc. 10th International Congress on Catalysis (Budapest, 1992), Elsevier, 1993, 2391. 4 D.B.Lukyanov, V.I.Shtral, S.N.Khadzhiev, J. Catal., 146 (1994) 87. 5 D.B.Lukyanov, N.S.Gnep, M.R.Guisnet, Ind. Eng. Chem. Res., 33 (1994) 223. 6 D.B.Lukyanov, N.S.Gnep, M.R.Guisnet, Ind. Eng. Chem. Res., 34 (1995) 516. 7 D.B.Lukyanov, Proc. 1lth Intern. Zeolite Conference (Seoul, 1996), Stud. Surf. Sci. Catal., 105 (1997) 1301. 8 L.N.Kurina, A.V.Demidov, N.G.Kalinina, A.A.Davydov, L.M.Koval, in "Chemical Synthesis Based on Single-Carbon Molecules", Nauka, Moscow, 1984, p.63 (in Russian). 9 M.I.Temkin, Adv. Catal., 28 (1979) 173.

Reaction Kinetics and the Developmentof Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999Elsevier Science B.V. All rights reserved.

307

Two-dimensional reactor modelling for the pure dehydrogenation of methanol to formaldehyde St. Ruff b, S. Schunk', G. Emig ", Th. WebeF, S. Braun c, G. BrenneF, F. Durst c "Lehrstuhl fiir Technische Chemie I, Universit~t Erlangen-Nfirnberg, Egerlandstr. 3, D-91058 Erlangen, Germany b Aventis Research & Technologies GmbH & Co KG, Industriepark HSchst, G 811, D-65926 Frankfurt am Main, Germany c Lehrstuhl fiir StrSmungsmechanik, Universittit Erlangen-Ntirnberg, Cauerstr. 4, D-91058 Erlangen, Germany

Abstract The simultaneous calculation of reaction kinetics and fluid dynamics is a topic of growing interest. Despite the increase in computer power, the derivation of an elementary reaction scheme, including the detailed 2D heat and mass transfer in a chemical reactor, is still time consuming. Therefore, a method has been developed in order to derive a reaction scheme applying 1D- and 2D-simulation tools. Following this method, a reaction scheme was derived which describes the uncatalysed dehydrogenation of methanol to anhydrous formaldehyde. 1

INTRODUCTION Anhydrous formaldehyde is produced conventionally from methanol in a two-step process by oxydehydrogenation (1) with subsequent separation of water and formaldehyde (2). An interesting alternative is the pure dehydrogenation catalysed by sodium aluminate at 1173 K yielding anhydrous formaldehyde directly. conventionally CH3OH + 8902 CH20-H~O

- CH20-H20 > CH20 + H20

pure dehydrogenation (1) (2)

CH3OH CH~0

> CH~O + H 2 > CO + H2

(3) (4)

Under these reaction conditions methanol is also converted to formaldehyde (3), carbon monoxide and hydrogen (4) without a catalyst. The uncatalysed reaction is superposed on the catalysed conversion, which was recently found not to be a heterogeneous reaction [1]. Elemental sodium catalyses the reaction homogeneously in the vapour phase after leaving the solid catalyst. The appropriate experiments were performed in a laminar flow tube reactor. A reactor simulation of this homogeneously catalysed reaction allows one to estimate the potential of a technical realisation. Unfortunately, there are no kinetic data available for either the uncatalysed or the sodium-catalysed conversion of methanol to anhydrous formaldehyde. The strategy applied here focused first on the simulation of the uncatalysed reaction in order to extend the reaction scheme and subsequently to describe the action of the catalyst. The appropriate simulation of the system requires an at least two-dimensional modelling of the heat and mass transfer processes involved as well as a reaction scheme including adequate kinetic

308 parameters. With respect to computer memory and calculation time, the number of species and reactions were limited, aiming at a reduced reaction scheme. The procedure was carried out as follows. The reaction scheme was fitted to experimental data by selecting sensitive elementary reactions using kinetic parameters from the literature [2]. In order to reduce computing time, the derivation was performed using the software package HomRea, neglecting heat and mass transfer phenomena in the reactor. This reaction scheme was used for the comprehensive two-dimensional reactor simulation with the software package Fastest 2D [3]. For a satisfactory agreement of experimental data and numerical results, only one collision factor k ~had to be adjusted. 2

EXPERIMENTAL SET-UP AND MODELLING TOOLS

2.1

Experimental set-up

The experimental investigations to study the uncatalysed decomposition of methanol were carried out in a corundum tube reactor (12 x 8 x 1000 mm) with a reaction zone of 450 mm centred in the axial middle of the tube. The wall temperature zone was kept at 1173 K, using an oven and a heat-pipe. Methanol was diluted with nitrogen generating 10% and 20% (mole fractions) feed streams. Methanol and a various amounts of nitrogen were led directly into the reaction zone through an inner tube, while a constant stream of 1000 ml/min nitrogen flowed through the annular gap between the corundum wall and the inner tube. The latter was installed in order to provide elemental sodium as catalyst for further experiments. The mean inflow velocity was varied between 0.75 and 1.5 m/s. The product stream was analysed by on-line gas chromatography (Hewlett Packard 5890 series II plus, Poraplot Q and molecular sieve columns) detecting methanol, formaldehyde, hydrogen and carbon monoxide as by-product and nitrogen as internal standard [4].

2.2 Modelling tools 2.1.1 H o m R e a

The software package HomRea (HOMogeneous REAction, Prof. Warnatz, Heidelberg) is able to compute time-dependent homogeneous reaction systems. It is possible to simulate systems with user-specified time-dependent values of pressure, volume and temperature generally used for combustion chemistry. Transforming time into position allows the one-dimensional modelling of a tube reactor, however, without considering radial gradients. Regarding reaction engineering aspects, HomRea does not consider the heat of reaction, different heat capacities of the feed stream with varying compositions and different flow velocities as a result of both changing gas temperatures (gas volume) and changing numbers of particles. HomRea offers the possibility of analysing pathways in a reaction scheme by listing the percentage of total formation and decomposition of all species in each reaction. This tool is very helpful for identifying sensitive reactions. The rate is given in form of a first-order power law as it is typical for free-radical reactions using mole fractions to define input compositions. For the calculation of

309 the rate coefficients, an extended Arrhenius equation (5) is used, while the Arrhenius parameters of a reverse reaction are calculated through thermodynamics. EA

rj = k 0i -e

R.T

Eduk~

9W a 9H x i

(5)

2.1.2 F a s t e s t 2 D

The two-dimensional numerical simulation was performed using an upgraded version of the CFD software package Fastest 2D. This CFD code is a multi-block finite volume code in cartesian]cylindrical coordinates, which solves the transport equations for mass, momentum, heat and chemical species on a non-staggered grid. The SIMPLE method [5] is adopted to calculate the pressure correction. The program handles solid blocks. Conjugate heat transfer is included. The approximation of the convective terms in the transport equations is obtained via a first-order upwind differencing scheme. The code is parallelized using message passing for distributed memory parallel computers. All calculations within the present work were carried out on a 24 processor PowerXplorer (Parsytec) or on a 10 processor CC system (Parsytec). The following steady-state, laminar transport equations have been solved (6)-(10): Continuity equation: 0 (pv~)+ 1 c)

Tzz

(6)

Vrl+OI !

Momentum equations:

r c}r

-2rll'~r

-~z pv-~ - Z /.l --~z

-~z pv'vz-11 -~r+-~-z

+ r-~r

prvr vz - r~l ~r + ~z

= Dr

(7)

=

(8)

c)z

Energy equation:

r~'/prv.T - r~-pc)"~)+~zz{pv.T- ~p"~zJ = ~l"rjAH,.j

(9)

Chemical species equations:

c) (prvrYi)+ ~zc) ~v zYi)-'7~" prDi'm rl bk

+

+

POi'm"

+

+ Ri

(10)

310 3

P R O C E D U R E AND R E S U L T S

Since the derivation of a reaction scheme and model discrimination is a timeconsuming procedure, short calculation times are desired. To achieve this, HomRea was used to develop the reaction scheme without considering any transport phenomena. Subsequently, the resulting mechanism was used for two-dimensional studies with Fastest 2D.

3.1

Derivation of the reaction scheme

An important aspect of the kinetic simulation using HomRea is the user-defined temperature-time profile with isothermal sections. In contrast to these block-profiles the laminar flow tube reactor profiles are parabolic, so that the latter must be appropriately transformed.

3.1.1 T e m p e r a t u r e p r o f i l e s a n d r e s i d e n c e t i m e

Two-dimensional reactor temperature profiles were calculated (with Fastest 2D) for a simple gas mixture of 10% methanol and 90% nitrogen at different mean inflow velocities and different wall temperatures without reaction. These twodimensional profiles were simplified to one-dimensional profiles consisting of six discrete areas corresponding to six time intervals for the HomRea calculations. The residence time was then calculated from the mean inflow velocity considering the varying reaction volume as a function of these temperatures using the ideal gas law. These time-dependent temperature profiles were used for the HomRea simulations, knowing that this is just a rough estimation. Both varying reaction volumes as a result of increasing number of particles with reaction progress and heat capacities with varying gas compositions were not considered, in addition to ignoring any kind of reaction heat.

3.1.2 R e a c t i o n m e c h a n i s m a n d k i n e t i c c o n s t a n t s

As mentioned above, the number of reactions and species has to be restricted as far as possible in order to limit the calculation time. After finding gas-phase spectroscopy in the present system to be very complex, a reaction mechanism for the uncatalysed dehydrogenation of methanol was derived from kinetic simulations with HomRea. Inspecting several potential elementary free-radical reactions for their sensitivity to the product distribution leads to a reaction system of 12 reactions involving 12 different species (Fig.l; two reactions are not shown - see below).

CH,+H20 CH20H M CH3+OH~~~~H2 ~H+M '"

CH30H

.2 ,, CH20

H 0

0 M +M

CO

H2 @ H CHO

Fig. 1. Reaction scheme

311

The dehydrogenation reaction is triggered by the homolytic cleavage of the C-O bond in methanol forming methyl and hydroxyl radicals (11). Both react with methanol forming methane, water and hydroxymethyl radicals (12, 13). The hydroxymethyl radicals are mainly formed by the reaction of methanol and H-atoms (14). Analysing the reaction pathway showed that this reaction is responsible for more than 97% of the overall methanol conversion. This coincides with the experimental findings, detecting methane and water in sm~l traces at the most. The hydroxymethyl radicals decompose to the main product formaldehyde either by collision with inert collision partners M (15) or by reaction with H-atoms (16). The same holds for formaldehyde (18,19), forming formyl radicals as intermediates, which decompose to carbon monoxide (19, 20). The recombination of two H-atoms forming hydrogen and of an H-atom and a hydroxyl radical forming water are not shown in the cycle because both reactions are of very minor interest concerning product distribution. Further reactions - e.g. the reaction of methanol and H-atoms forming methoxyl radicals - were not significant for the product distribution. Extending the scheme by the reaction of two methyl radicals recombining to give ethane led to no significant amount of ethane in the product stream. This coincides clearly with the experimental results. Several other reactions which may possibly proceed were found to be insensitive to the product distribution. Both kinetic and thermodynamic data for the different elementary reactions and substances which were likely to proceed were taken from the literature without modification. Only data for four reactions were modified within the limit of error. Table 1 shows the collision factors k ~ the temperature exponents a and the activation energies E A. Table 1 Kinetic constants

Data from literature k~ o~ E A [kJ/mol] Eqn. [cm, mol, s] [-] 404,1 (11) 9,51" 10 ~ -4,30 9,40"10 TM 0 376,0 9,00"10 TM 0 (12) 41,1 1,00"10 TM 0 7,1 (13) 4,00"10 TM 0 25,5 (14) 5,00"10 TM 0 105,0 (15) 3,00"1013 0 0,0 (16) 5,00"10 TM 0 320,0 (17) 2,30" 10 l~ 1,05 13,7 (18) 7,10"1014 0 70,3 (19) 9,00"10 TM 0 0,0 (20) (21) 4 1,80" 10 TM -1,00 0,0 (22) 5 2,20-10 = -2,00 0,0 1 Only the changed data are shown 2 Data from HomRea 3 Data from [2]

Changed data for simulation 1 k~ {X EA [cm, mol, s] [-] [kJ/mol] 1,71"10 TM

0

384,0

20,0

300,0 8,7 6,50" 1014

5

H+H H + OH

>H~.

> H20

312

3.2

R e s u l t s of k i n e t i c m o d e l l i n g

The reaction scheme and the kinetic p a r a m e t e r s were fitted to experimental 10 data for a 1173 K reactor temperature - ~ CH3OH exp. I and 10% methanol in the feed s t r e a m at 8 - * CH3OH sire.I four different mean inflow velocities. The -4- CH20 exp. I use of the kinetic data from Table 1 for - ~". . . . 9 ~ . ....... o....o-- ~ the simulation with HomRea leads to very good agreement. Fig. 2 shows the mole fraction of both methanol and formaldehyde at the reactor exit versus 2 ................ ....... ii-....... ~ the m e a n inflow velocity. o Methanol is described very well over the I 0 t ......... I four velocities whereas the simulated 0.75 1.00 1.25 1.50 values for formaldehyde are too small at mean inflow velocity [m/s] low velocities and slightly too high at the highest velocity. The results are Fig. 2. Experimental data compared satisfactory, particularly when con- with simulation results (HomRea) sidering the simplifications and the factors neglected. Similar results are obtained by simulating the dehydrogenation at 20% methanol in the feed stream and an 1123 K reactor temperature. The trend is always described well, while the differences between experimental results and simulation data are in the range of 20% absolute.

3.3

The t w o - d i m e n s i o n a l r e a c t o r s i m u l a t i o n

The three grid-level computational domain consists of 5650 grid points on the coarsest grid. It is split into four blocks and describes only one symmetric half of the whole reactor. In order to guarantee a correct representation of the experimental set-up, the following boundary conditions were used. At both inlets uniform velocity profiles resulting in the experimental mass flow rates and the mole fractions of all species were set constant. A mixture of methanol and nitrogen, led through the inner tube, is mixed with pure nitrogen from the a n n u l a r gap. The temperature at the inflow and downstream to the reaction zone was set constant at 523 K. The reaction zone itself was heated via a heat pipe at a constant temperature of 1173 K. Downstream of the reaction zone the wall temperature field was fixed according to the measured values. At the outlet boundary a zero gradient condition for all variables was used.

3.3.1 A d j u s t m e n t of the k i n e t i c d a t a On the basis of the reaction scheme derived from HomRea, two-dimensional reactor simulations were performed for four different inflow velocities at 1173 K and 10% (mole fraction) methanol in nitrogen. It was found that the simulation data for the conversion of methanol were in good agreement with the experimental data for all

313 inflow velocities. However, the calculated selectivity of formaldehyde was approximately 15% (absolute) too low in comparison with the experimental data. Since the selectivity of formaldehyde is strongly dominated by the speed of reactions (16) and (18), the influence of the collision factors of these two reactions on the conversion and selectivity was investigated. It was found t h a t a reduction in the collision factor of reaction (18) leads to an increase in selectivity, while the methanol conversion remains constant. For this reason, the collision factor of reaction (18) was set to 1.5-l0 s cm3/(s.mol) for all subsequent simulations.

3.3.2 S i m u l a t i o n w i t h the adjusted collision factor Simulations with the adjusted collision factor were performed for different inflow velocities and mole fractions of methanol at 1173 K. The experimental data and the simulation results concerning methanol conversion are in very good agreement in the case of 10% methanol in the feed stream (Fig. 3a). Although a deviation of approximately 10% absolute occurs at 20% methanol feed, the trend is found to be correct. 70.0

70.0

60.0 t-

.o

" D, " ' ,

-" -" x ( C H 3 O H ) = 20 %, exp. O - - ~ x ( C H 3 O H ) = 10 %, sim. D - - ~ x ( C H 3 O H ) = 20 %, sim.

8 o era eID

i

o- 60.0

[]

>,

50.0

50.0

{D

c

I

40.0

R 40.0

30.0

E ,- 30.0

E 20.0 10.0

0.5

o

,

, 110 115 mean inflow velocity [m/s]

20.0

0.5

oxp. a/,," w t cf ,

--- x ( C H 3 O H ) = 20 %, exp. | - - ~ x ( C H 3 O H ) ffi 10 %, sim. r~ - - -El x ( C H 3 O H ) = 20 %, sim. i

1.0

0

/

1.5

mean inflow velocity [m/s]

Fig. 3. C o m p a r i s o n of e x p e r i m e n t a l and s i m u l a t i o n results (Fastest 2D) of c o n v e r s i o n (a) and selectivity (b) In the case of formaldehyde selectivity similar results can be found (Fig. 3b). The deviation decreases with increasing mean inflow velocity for 20% methanol feed. The results for the temperature distribution and the mass fractions of methanol, fomaldehyde and carbon monoxide can be seen for a specified velocity in Fig. 4.

4

S U M M A R Y AND D I S C U S S I O N

In a first step, a scheme of elementary reactions for methanol dehydrogenation was developed using HomRea with the simplifications mentioned above. The reaction scheme and the kinetic data were fitted to experimental results obtained at 1173 K and 10% methanol in the feed stream.

314

8 7

T [K] 1173.15 1073.15

4 3 2 1

773.15 673.15 573.15 473.15

6 5

973.15 873.15

8 7

6 5

0.8

,_,E' N 0.6 I

0.7

0.5

'E' 0.6

/.-7

8 7

y (CH20) 0.017813 0.017104

8 7

0.092571 0.073429 0.065356 0.064068

4 3 2 1

0.012254 0.007312 0.001747 0.000017

4 3 2 1

4 3 2 1 1-

0.9 0.7P

y (CH3OH) 0.150000 0.130857

0.111714 0.103945

~/--2~

'i~S~e~ ,,4~ r

N 0.5~=. /r 6~-~,

~ f6"~

6 5

0.9~-

0.016364 0.015722

~

j

I I

6 5

I

y (CO) 0.027000 0.023429 0.019857 0.016286 0.012714 0.009143 0.005571 0.002000

_

0.9

8-----8--

0.8

f-,,~

7

!o.6 0.5 0-41

0

0.2

.

3

-0.004

0 . 3 ~

~ 0

rim]

0.004

-0.004

0

r[m]

0.004

02

-0.004

0

r[m]

0.004

0.2 -0.004

0

rim]

0.004

Fig. 4. P r o f i l e s of t e m p e r a t u r e and m a s s f r a c t i o n s of the m a i n s p e c i e s

The following 2D-simulations resulted in good agreement for the methanol conversion. Owing to the unacceptably high deviation of the formaldehyde selectivity, a further parameter study with Fastest 2D was performed, varying the collision parameters of reactions (16) and (18). The subsequent simulations with the adjusted factors led to good results for an 1173 K reactor temperature and 10% methanol feed. Moreover, the trend is represented correctly by the simulation for 1123 K (results not shown) and 20% methanol feed. The higher deviations can be explained by the fact that the reaction scheme and the kinetic data were fitted to experimental data from 1173 K and 10% methanol feed only. The applied procedure turned out to be very efficient; it included the 2Dcalculations of the reactor temperature profiles without any reaction, the simplification to 1D-profiles for the HomRea simulations and the use of the resulting reaction scheme after adjustment of the collision factors. The successful simulation of the pure dehydrogenation of methanol to anhydrous formaldehyde indicates how computational fluid dynamics, including the equations for species transport, heat and mass transfer, momentum transfer and mass conservation, can be used to provide an insight into chemical reactions, providing information on chemical reaction schemes. The experience with the uncatalysed system will be now extended to the catalytic case.

315 ACKNOWLEDGEMENT This work Was partially supported by the BMBF (Bundesministerium fiir Bildung und Forschung) as project 03D0012B9 and by the Bayerische Forschungsstii~ung within the frame of FORTWIHR (Bavarian Consortium for High Performance Scientific Computing). We are grateful to Prof. Warnatz, Heidelberg, for providing the software package HomRea. REFERENCES

[1] M. Bender, Doctoral Thesis, Bochum, 1997 [2] J. Warnatz, Rate coefficients in the C/H/O-System, Comb. Chem., Springer, Berlin 1994 [3] Fastest 2D, User manual V 3.52, Invent Computing GmbH, Erlangen 1997 [4] St. Ruf, Doctoral Thesis, Erlangen, 1998 [5] S.V. Patankar and D.B. Spalding, A calculation procedure of heat, mass and momentum transfer in three-dimensional parabolic flows J. Heat Mass Transfer, 15 (1972), p.1787

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Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

317

Kinetic modelling of enzymatic chiral resolution of (+)2-methyl butyric acid R. Garcia, M. Martinez, T. Garcia and J. Aracil D e p a r t m e n t Of Chemical Engineering. Faculty of Chemistry. Complutense University. 28040 Madrid. Spain. E-mail: [email protected] Abstract A kinetic study of 2-methyl butyric acid esterification reaction catalyzed by immobilized lipase (from Candida antarctica) has been made in a range of 62 - 78~ temperature. The kinetic characteristics observed in the chiral resolution of (~) 2-methyl butyric acid by esterification reaction with n-octanol were found to conform to an ordered bi-bi mechanism with competitive inhibition by reactants and products. According to the mechanism m a x i m u m reaction rate, Michaelis Menten constants and inhibition constants were determined. The results suggest t h a t lipase should recognize the chirality of 2-methyl butyric acid molecule in the binding process to the active site of lipase. A good quality of fit was observed by fitting experimental rate data to the kinetic model. 1. I N T R O D U C T I O N In recent years, the use of enzymes for preparation of optically enriched compounds has become an alternative to chemical synthesis. Lipases have been routinely used for the resolution of racemic alcohols and carboxylic acids through asymmetric hydrolysis of the corresponding esters. Immobilized lipases are particulary insteresting for industrial purposes, since they can be easily handled [1]. Unlike soluble lipases, a lipase in immobilized form is strongly bound to a solid phase, making it easy to separate the catalyst from the reaction mixture and to reuse it [2]. The current shift towards enzymatic processes in the enantioselective resolution of 2-substituted butyric acids requires the study of the kinetic behaviour of immobilized lipases and the development of general rate expressions to predict the yield of the pure enantiomer as a function of the operating variables. The use of a proper kinetic model is a basic element for quantitative description of the course of enzymatic resolution of racemates and the scale-up of esterification process. The validity of this model must then be tested for the experimental reaction data. To elucidate the mechanism of enzymatic reaction in aqueous solutions, it is necessary to consider the initial rate kinetics as one of the useful techniques. Moreover, since enzymes possess catalytic activity in organic media, there is a great demand for mechanistic information about enzyme kinetics in nonaqueous systems acting as biocatalysts [3]. However, it would seem that the application of enzymatic reaction in organic phase

318 will be at variance with established conventional Michaelis Menten kinetics for enzyme catalysis in a homogeneous aqueous system. Several kinetic studies from single power model to ping-pong bi-bi or ordered bi-bi mechanisms have been proposed [4]. However, the mechanism generally accepted for esterification reactions catalyzed by immobilized lipases involves the formation of an acyl-enzyme intermediate, followed by the incorporation of an alcohol molecule to produce the molecules of ester and water [5]. The aim of the present work is to develop a kinetic model to predict reaction rates for the synthesis of a pure isomer of n-octyl 2-methyl b u t y r a t e ester using a commercial immobilized lipase as catalysts.

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

(+) 2-methyl butyric acid and n-octanol, used as reactants, were supplied by Fluka and Henkel Iberica respectively. The catalyst used was a triacylglycerol hydrolase (E-3.1.1.3 lipase) from Candida antarctica immobilized on a macroporous acrylic resin. The commercial lipase (Novozym 435) was supplied by Novo Nordisk Bioindustrial A/S. Kinetic experiments were carried out in a batch stirred reactor. Samples were taken at regular intervals and then analyzed by gas-liquid c h r o m a t o g r a p h y and mass spectrometry. A description of the experimental a p p a r a t u s and analytical methods is given elsewere [6].

3. R E S U L T S

Kinetic experiments were carried out over a wide range of operating conditions, summarized in table 1.

Table 1 Operating conditions used for kinetic experiments. T e m p e r a t u r e (~ 62 Pressure (KPa) 3 Catalyst concentration (%) Acid/Alcohol molar ratio 1:10 1:7 1:5 1:2 0 Initial ester excess (%) 0 Initial water excess (%) Reaction time (hours) Stirrer speed (r.p.m.)

70 94.6 5 1:1 2:1 25 25 8-9 700

78 7 5:1 75 75

7:1

10:1

Figures 1 and 2 show the effect of temperature and catalyst concentration on acid conversion respectively. As expected [7], conversion increased with increasing temperature. A similar trend was observed for the concentration of lipase. Experimental reaction rates were determined for different chemical species involved in the reaction.

319

1.o XAc

1.01 XAc

0.8 I-,,-~OCl

0.81 f-'-~~

/ -='-4o/oI

0"6~ L~,~ ~

o.61

.j...J.---'-- ._._..__._.-.

0.4

~:/-~-.~

0 ~

. . . . 100

'200

300

Time (min)

.-'J/~

~.~

._.__...__.t,

0.2 ~ ~400" '500

Figure 1. Influence of t e m p e r a t u r e on acid conversion. 5% of catalyst, Ac/AI: 1/1, P : 7 1 0 m m H g .

0

0

100

200 Time (min)

300

Figure 2. Influence of catalyst concentration on acid conversion T:70~ Ac/AI: 1/1, P:710mmHg.

Reactions were carried out in excess of one of the r e a c t a n t s (octanol or 2-methyl butyric acid). For these cases an increase in the concentration of the other r e a c t a n t (2-methyl butyric acid or octanol) led to an increase in the reaction rate. However, as the minority r e a c t a n t disappeared the decrease in the reaction rate was larger t h a n expected. This fact suggested the possibility of enzymatic inhibition effects caused by the r e a c t a n t in excess. Figures 3 and 4 are plots of reaction rate vs n u m b e r of moles of the minority r e a c t a n t for different acid/alcohol molar ratios. In these plots the inhibition effect of the r e a c t a n t in excess can be observed.

XAc

1.0

1.01 XAc

I--,-~,~,.~ ]

0.8

I I

~I--

0.6.

[

"- Ac~AI:1/5

~1~,--Ac/AI:

0.4

o"8t| 1I-'e --'-'~'':~ -A~/~:211

0 61 9 I-&,-AclAI:7/I -~A~1~zoa

Ac/AI: ]J2

o. 1 t

117

----~"

0.2

,---J"

,/-"~-J ,J .j~_/~ , ./

--I-' 0

100

200 :$00 400 Time (min)

500

Figure 3. Influence of a n-octanol excess on conversion. 5% of catalyst P: 710mmHg, T: 62~

0

100

200 300 400 Time (min)

500

Figure 4. Influence of a 2-methyl butyric acid excess on conversion, 5% catalyst, T:70~ P:710mmHg.

In a similar way experiments were carried out to d e t e r m i n e how the presence of one of the products (n-octyl-2-methyl b u t y r a t e or water) in the reaction mixture affected the reaction. Results revealed a drastic decrease in conversion, indicating inhibition effects due to the reaction products.

320 4. K I N E T I C M O D E L

In order to determine the effect of the initial concentration of immobilized lipase on the reaction kinetics, acid conversion was plotted vs catalyst concentration for different reaction times, as shown in figure 5. Linear plots were obtained for all the temperature values tested, indicating a first-order relationship between reaction rate and catalyst concentration. Kinetic reaction order was also studied in previous works, founding a second order for global reaction, first order with respect to acid and first order with respect to alcohol [7]. The mechanism generally acepted for esterification reactions catalysed by immobilized lipases involves the formation of an acyl-enzyme intermediate, followed by the incorporation of an alcohol molecule to produce a molecule of ester and water. The most frequent inhibition type is the competitive inhibition. Competitive inhibition occurs when free enzyme can combine either with the substrate, to give the productive complex, or with an inhibitor, to give a nonproductive complex, but not with both. The commonest reason for this p a t t e r n is that inhibitor is a substrate analogue, and sits precisely where the substrate should be, in the enzyme's active site. In some cases there is competition for the formation of a complex with the free lipase, and in other cases the competition is given by the formation of ternary complexes with the most active complex, enzyme-acid. Therefore, the general mechanism of esterification corresponding to an ordered bi-bi system, including inhibition terms for reactants and products, can be espressed as follows: Enz + Ac

Enz-Ac

Enz-W

+

+

Alc

Est

Ac-Enz-Alc

Est-Enz-W

Inhibition steps:

Reaction steps: E + Ac ~ EAc

EstEW EW

EAc.'.

+ Alc ~

AIcEAc

KmA~,K~A~

AIcEAc

~-+ E s t E W ++ E W

Enz + W

~

E + A lc ++ E A lc .'. K iAIc

.'. K mAI~,K iAl~

9 K0q,(-rAr

rm..~,(-rA*)'

+ E s t .'. K mEs,K iEs

~-+ E + W .'. K , ~ w , K ~ w

E + E s t ++ E E s t

(1)

:. K'aE~t

EAc+

Ac ~

AcEAc.'.

EAc+

E s t ++ E s t E A c . ' .

K~A~ K'~'E~t

321 1.0

0.8

XAc

0.8

[ ,t:-~o,m |

o.e

l .,,~oi

I

0.6 XAc 0.4

~ / /

0.4

0.2

0.2 o

2

'3 . . . .

'4

0

's

'

6

Catalyst (%)

F i g u r e 5. V a r i a t i o n of c o n v e r s i o n vs % of c a t a l y s t a t d i f f e r e n t t i m e s of r e a c t i o n P: 7 1 0 m m H g , T: 7 0 ~

-

--- Calculated

8..i, :

o

loo

260

,

c

300 400 Time (rain)

9

| |~

]

soo

F i g u r e 6. S i m u l a t i o n of s y n t h e s i s n-octyl 2 - m e t h y l b u t y r a t e . 5% of c a t a l y s t , T:78~ P:710mmHg.

T h e a b o v e s c h e m e c o r r e s p o n d s to a n o r d e r e d r e a c t i o n m e c h a n i s m i n v o l v i n g two r e a c t a n t s a n d two p r o d u c t s . T h e r e f o r e M i c h a e l i s - M e n t e n a s s u m p t i o n s c a n be a p p l i e d to o b t a i n t h e e x p r e s s i o n for t h e r e a c t i o n r a t e as follows: ( - rAc ) fn,a,, ((--FAc) =

f (--rAc) .....

CwKmE . . . . K eq

)

K eq

C A

, ~] +(--rA ) ~ .CA K A, [1+ , ~/ + K iA Ic fl . . . . . . . . ~ K iAc fl

~,

. . . . . (. CAl ~ (--rA) + I , _ F A )r t 2 g I IK g [ | q- __2._7_______~_ [ -t- . . . . . . . . . . . . . . . . ~. K iAlc ) +

"x C Est C w 1

CAI

(--rA) ~ K.A K A, [1+ ..........

rAc)~,.~(C Ar Ale

. . . . . . . . . . (-t- ~,-- IAc )max k.. Act.. Alc -1- ~

f

CEtK - .....

K eq r

w ( , C A l ") ] l -t- _2__,_,_,_,_,_,_,_~_[+ ~ K iAIc )

f Ac)maxCAcCEstKmw K lAcK eq

(,

^'~ , CA 1 "1- K~A~c ) +

+ (--rA~)f.xCE~'C_w_(I+ C,A~) + _(--rac)rma_= ~C A,cCwK___mAc + K Cq \ K iAc fl K iw +

r (--rA~).....

CAicCwCEst

K iEst

q_ ( -

r

Ac) f. . . . C A i c C w C E s t K iAicK eq

(2) F o r a 1/1 m o l a r r a t i o acid/alcohol: CAcO = CAIO CAr

"'" X A c o -- XAiO

XAcO

=

CAcO. X A l o

C A c -- CAI

=

CAcO.(1

(3)

- XAc )

CEs = C w = CAco.XA~ A n d b o u n d a r y conditions" t=0

" CA~ =CA~ o .'. X A r

t=t

"

CAc

"- C A c

"

XAc

(4) =

(CAc 0 --CAc ) / CAc 0

322 By integrating the general equation the following expression is obtained: t = o;.Ln (~ -~eq -1)(XAc-(~/-Keq + l ) - - ~ e q ) + ~.Ln(X2c.(Keq - 1)-2Keq.XA~ + Keq)+ ( X A c . ( ~ --1)--.K~~,.( K ~ q +1)) + ~.Ln

(.~--Keq --1)

(.~/Keq

+

1)

+7.LnKeq +8. XAc + g. XAc

(5) Where greek symbols are the coeficients defined by m a t h e m a t i c a l relations which involve the kinetic constants. The entire development of kinetic model can be seen in previous works [8]. When the reaction takes place at reduced pressure (vaccum) w a t e r is continously removed from the medium. Then the reverse reaction (hydrolysis) can be considered as negligible compared to the forward reaction (esterification) and the process can be considered as an irreversible esterification system. In this case, general equation (eq. 2) can be simplyfied and integrated expresion involves only four p a r a m e t e r s [9].

5. D I S C U S S I O N

P a r a m e t e r s were estimated from values of experimental rates by applying numerical methods based on M a r q u a r d t algorithm. Table 2 shows the p a r a m e t e r s obtained for the three t e m p e r a t u r e values tested.

Table 2. P a r a m e t e r values obtained for different temperatures. Constant T = 62~ T = 70~ (-rAc)max/E0 (mol.l-l.min 1) 0.0098 0.0115 Kmac (mol/1) 2.5 0.77 Kraal (mol/1) 4.0 2.1 Km~ (mol/1) 2.4 0.41 Kmw (mol/1) 3.7 2.33 KiAc (mol/1) 7.5 5.9 KiA1 (mol/1) 1.11 0.09 KiEs (mol/1) 2.9 1.1 Kiw (mol/1) 1.04 0.56 K'iAc (mol/1) 0.3 0.352 K'iA1 (mol/1) 1.7 1.35 K"iEs (mol/1) 0.65 0.178

T = 78~

0.0141 0.44 0.64 0.46 0.35 6.1 0.021 0.53 0.13 0.08 0.078 1.31

Figure 6 is a plot of of experimental and calculated conversion values obtained for a representive experiment. Similar curves were obtained for all

323 the t e m p e r a t u r e s and enzyme concentration tested. The a g r e e m e n t between conversion values m e a s u r e d experimentally and those calculated from the model indicates t h a t the kinetic model proposed is adequate to describe the synthesis of n-octyl 2-methyl b u t y r a t e over enzymatic catalysts. T h e r m o d i n a m i c a l p a r a m e t e r s were also studied using single expressions like Arrhenius equation and Van't Hoff equation. Table 3 s u m a r i z e s the energetic changes and activation energy obtained from kinetic, Michaelis-Menten and inhibition constants at different temperatures.

Table 3. T h e r m o d i n a m i c p a r a m e t e r values obtained for the esterification of m e t h y l butyric acid with n-octanol. Constant (-AH ~ (Kcal/mol) (-AS ~ (cal/mol.K) KmAc 24.6 71.9 KmA1 27.1 74.5 KmEs 23.4 68.9 Kmw 33.4 96.8 Kiac 2.9 4.8 KiA1 56.0 167.7 KiEs 24.0 69.9 Kiw 29.4 87.5 K'inc 18.8 58.1 K'inl 22.0 64.8 Activation Energy: Ea=5.45 Kcal/mol Preexponential factor: k0=37.34mol/1.min.g E n t a l p y change: (AH~ Kcal/mol Entropy change: (AS~ cal/mol.K

_

Finally, an additional p a r a m e t e r was used for the stereospecificity of the used lipase. This p a r a r a m e t e r was the enantiomeric excess defined as follows: C R -

C S

ee = ~ Cs + C R

(6)

Where Cs and CR are the concentrations of 2-methyl butyric acid in the reaction m e d i u m for the isomers S and R, respectively. Obviously, in this case the m a y o r i t a r y ester formed was the S isomer and, therefore, the major isomer present in the reaction m e d i u m was the R isomer of 2-methyl butyric acid. Then, the value of ee p a r a m e t e r varies from 0% to 100% (0% for equal concentrations of R and S isomers, and 100% for pure isomer). Experimentally, best result were obtained at 78~ of reaction t e m p e r a t u r e and 7% of initial concentration of catalyst w h e n acid/alcohol molar ratio was fixed at 1.

324 At these operating conditions maximum 2-methyl-butyric acid conversion was 41% and enantiomeric excess was 34.4%. S isomer of n-octyl 2-methyl butyrate was only found as product. This result presents relevant importance because of the possibilities of S isomer of obtained ester in pharmaceutical applications.

6. N O M E N C L A T U R E Ac: F a t t y acid Al: F a t t y alcohol CA: Concentration of component A (mol/1) E: Free enzyme ee: Enantiomeric excess Es: Ester Keq: Equilibrium constant KiA: Inhibition constant for component A (mol/1) KmA: Michaelis-Menten constant for component A (mol/1) NA: Number of moles of component A (mol) P: Pressure (ram Hg) (-rA): Reaction rate for component A consumption (mol. 1-1.gcat-l.min-1) (-rn)maxf: Forward maximum reaction rate for component A (mol. 1-1.gcat1.min-1) (-rn)maxr: Reverse maximum reaction rate for component A (mol. 1-1.gcat-l.min-

1)

t: Reaction time (min) T: Temperature (~ W: Water w: Amount of catalyst (%) XA: Conversion of A component

7. R E F E R E N C E S

1 P. Eigtved, T.T. Hansen and C.A. Miller, Proceedings of the World Conference on Biotechnology for the Fats and Oils Industry. T.H. Applewhite, AOCS, 1988. 2 L.H. Porsorske, JAOCS, 61 N ~ 11 (1984) 38 3 J.S. Dordick, Enz. Microb. Technol., N ~ 11 (1989) 194 4 A. Zaks and A.M. Klibanov, Proc. Natl. Acad. Sci. USA (1985) 3192 5 F. Shiraishi, Enz. Microb. Technol., N ~ 2 (1993)150. 6 T. Garcia, M. Martinez and J. Aracil, Enz. Microb. Technol. 15 (1993) 607 7 T. Garcia, M. Martinez and J. Aracil, Trans IChemE C 71 (1993) 47. 8 T. Garcia, A. Coter6n, M. Martinez and J. Aracil, Chem. Eng. Sci. 51 (1996) 2841. 9 T. Garcia, A. Coter6n, M. Martinez and J. Aracil, Trans IChemE C. 73 (1995) 140

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

Regioselective synthesis Kinetic modelling

of

325

monoglycerides.

T. Garcia, M. Martinez, D. Garcia and J. Aracil" D e p a r t m e n t of Chemical Engineering. Faculty of Chemistry. Complutense University. 28040 Madrid. Spain. E-marl: [email protected] Abstract A kinetic study of ricinoleic acid and glycerol esterification reaction catalyzed by immobilized lipase from Candida antarctica (E.C.3.1.1.3) has been made in a range of 70 - 80~ The best o p e r a t i n g conditions were found at vacuum conditions, which permits the study of esterification reaction as irreversible process. High selectivities for monoglyceride were obtained by controlling the initial concentration of catalyst on the reaction medium. The kinetic characteristics were found to conform to an ordered bi-bi mechanism with competitive inhibition by reactants and products. Michaelis Menten constants and inhibition constants were determined according to the mechanism maximum reaction rate. A good quality of fit was observed by fitting experimental rate data to the kinetic model. 1. I N T R O D U C T I O N Nowadays there is an increasing interest in the application of polyfunctional molecules as raw materials for the Fine Chemical industry. The use of these products is only possible due to the availability of highly selective and specific catalysts, which control the reactivity of these products at a functional group level [1]. Immobilized enzymes make it possible to reach the above goals, and they can be recovered and reused at the end of the reaction, in the same way as heterogeneous catalyst [2]. Monoglycerides are non-ionic surfactants used as food, cosmetic and pharmaceutical emulsifiers and stabilizers. They are the largest single type of food-grade emulsifiers. At present, most monoglycerides are prepared by either a straightforward chemical esterification of glycerol and fatty acids or by glycerolysis of natural fats [3]. During the past few decades, lipases have also been studied for their potential use in the production of monoglycerides. Several workers have reported enzymatic catalysis in solvent-free systems with the aim of making the processes commercially feasible. These efforts not only excluded the toxicity problem of the solvent and surfactants, but also reduced many steps in the purification process. Although enzymes used in a solvent-free system have several advantages, their application is quite limited because most organic substrates have high melting points and the enzyme is not stable at a higher temperature [4].

326 Different types of reactors can be used for the process scale operations with immobilized enzymes. A mathematical model applicable to the analysis, design and simulation of heterogeneous enzymatic processes reacting in a stirred-tank reactor has been formulated. Information regarding the kinetics of the reaction is essential for understanding the reaction mechanism, as well as for a rational design of the esterification reactor for future scale-up. In this study we have examined the kinetics of the esterification of ricinoleic acid and glycerol in a solvent-free system, catalyzed by nonspecific lipase from Candida antarctica. The mechanism generally accepted for esterification reactions catalyzed by immobilized lipases involves the formation of an acyl-enzyme intermediate, followed by the incorporation of an alcohol molecule to produce the molecules of ester and water [5]. The aim of the present work is to develop a kinetic model to predict reaction rates for the synthesis of monoglyceride ester of ricinoleic acid using a commercial immobilized lipase as catalyst.

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

Glycerol of purity > 98% (w/w) was supplied by Henkel Iberica, and ricinoleic acid of purity > 98% (w/w) was supplied by Fluka. The catalyst used was an immobilized preparation of a thermostable lipase, Novozym 435, particularly useful for the synthesis of esters. This catalytic system was kindly supplied by Novo Nordisk Bioindustri S.A. The enzyme is a triacylglycerol hydrolase, E.C. 3.1.1.3, which simultaneously acts as an effective carboxylesterase. The positional specificity of this system depends on the reactants. In some reactions it shows 1,3 positional specificity, whereas in other reactions the lipase functions as a non specific lipase. In the manufacture of Novozym 435, recombinant DNA technology was used. The enzyme obtained by this procedure was immobilized on a macroporous acrylic resin. The final product consists of bead-shaped particles with a diameter in the range of 0.3-0.9 mm. The bulk density of the catalytic system is approx. 430 kg/m~. The ester synthesis activity of Novozym 435 was 7000 PLU/g, where One Propyl Laurate Unit (1 PLU) is defined as 1 ~mol of lauric acid converted to propyl laurate per minute under standard conditions. The experiments were carried out in a batch stirred tank reactor (BSTR) of 5x10 -4 m ~ volume, under fLxed conditions of pressure and temperature. Pressure, stirring speed and temperature controllers were provided. The reactor was also equipped with stationary baffles attached along the surface. A marine type propeller was employed. Impeller speed tested between 500 and 1200 rpm was fixed at 700rpm to maintain the catalyst particles in suspension and eliminate the mass transfer limitations along the reaction course. The reactor, which was immersed in a constant-temperature bath, was capable of maintaining the reaction temperature within •176 of that

327 desired for the reaction. Samples were withdrawn at regular time intervals, for 6 hours. The reactants and catalyst were added to the reactor. When the vacuum pump provided a working pressure of 8 KPa, the reaction mixture was heated to the desired temperature, then the reactants and catalyst were well mixed for 2 hours at 700 rpm. The samples taken at the end of the reaction were analysed by gas chromatography. No by-products were detected in the samples analysed. During each experiment, the following variables remained constant: reactor temperature, impeller speed, and pressure. Once the reaction was finished, the catalyst was separated by filtration, washed with a solvent and dried, so that the lipase was recovered for a next operation. The ester and the corresponding acid and alcohol were separated by gasliquid chromatography using a Hewlett-Packard (HP 5890 SII) a p p a r a t u s equipped with a flame ionisation detector (300~ and a split-splitless injector (270~ 30 s). A OV-1 fused silica capillary column (HewlettPackard 12-m length, 0.31-mm i.d. 0.17-mm film) was used with helium as carrier gas (0.65 ml min-1). The oven was maintained at 150~ for 5min, and then increased at a rate of 10~ min 1 to 270~ until all components had eluted. Qualitative data were obtained by a mass-spectrometry system (HewlettPackard, HP6890S, ionisation energy, 70 eV; scan speed, 1100 amu.s-1; mass range, 40-400 amu) coupled to the gas chromatograph and using the column and conditions described above. Quantitative data were obtained by peak areas integration using a Hewlett-Packard (HP 3396A) integrator. Ester, acid and alcohol concentrations were estimated by using an internal standard. Response factors were determined using response factor of products versus standard concentration as calibration curve.

3. R E S U L T S Experiments were carried out to determine the influence of operating variables on ester yield and to perform a kinetic study. Preliminary experiments were carried out at stirring speeds between 500 and 1200 rpm. No mass transfer limitation was detected in this interval. For this reason, the stirring speed was fLxed at 700 rpm.. Kinetic experiments were carried out over a wide range of operating conditions: Figures 1 and 2 show the effect of t e m p e r a t u r e and catalyst concentration on acid conversion respectively. As expected [6], conversion increased with increasing t e m p e r a t u r e . A similar trend was observed for the concentration of lipase. Experimental reaction rates were determined for different chemical species involved in the reaction. Reactions were also carried out in order to know the influences of both working pressure and acid/alcohol molar ratio. Figure 3 shows the influence of pressure on acid conversion. Vacuum pressures allow an increase on acid conversion due to the continuous removal of the w a t e r formed during the course of the reaction, so the reaction equilibrium is driven

328 towards the formation of products, and the system can be considered to operate under irreversible conditions. For the cases in which acid/alcohol molar ratio was studied, an increase in the concentration of the reactant in excess led to an increase in the reaction rate. However, the decrease in the reaction rate was larger than expected for the experiments carried out at high molar ratios. This fact suggested the possibility of enzymatic inhibition effects caused by the reactant in excess. Figure 4 is a plot of conversion of minority reactant vs time for different acid/alcohol molar ratios.

1)0-

80-

~3

100

i

80

,o

,3

20 o

9

I" 'cl

,. I "

75"C 70"C

o

60

2~o

o

~o

.

o

,

~o

.

,

2o0

Time (min)

%I

,oo

Time (min)

Figure 1. Influence of temperature on acid conversion. 3% of catalyst, Ac/AI: 1/1, P : 6 0 m m H g .

Figure 2. Influence of catalyst concentration on acid conversion T:75~ Ac/AI: 1/1, P:60mmHg.

A A .A

80 80 9

9

oo

9

9

9

60. Z

r~

~o

3.0

40.

+

60

40

9 380 mmHg o

400 Time (min)

Figure 3. Influence of pressure on acid conversion. 5% of catalyst Ac/AI: 1/1, T: 75~

o

I I ...... | /3 i . R M : I / S II

g 2o-m

9~om.,

20

9 RM: I/! j 9 i~/: 1/2

+

+

0

0

~

200

3o0

400

T i m e (rain)

Figure 4. Influence of Ac/A1 molar ratio on ricinoleic acid conversion, 5% catalyst, T:75~ P:60mmHg.

329 In all cases, mono and diglyceride were obtained. However, monoglyceride was always the majority product. Table 1 shows the influence of reaction variables on selectivity of monoglyceride. As can be seen, very low influences were detected and high selectivity levels were obtained.

Table 1. Influence of operating variables on monoester selectivity. Reaction time: 4 hours. .........T(~ .. ...................P ( m m H g ) ....... Cat aiYst(%)" ............Ac/AiMR [... ~electivity i%i::: 70 60 3 1:1 93.2 75 60 3 1:1 95.7 80 60 3 1:1 97.2 75 60 1.5 1:1 95.3 75 60 4.5 1:1 95.5 75 700 3 1:1 89.7 75 380 3 1:1 91.3 75 60 3 1:2 96.3 75 60 3 1:5 97.8 75 60 3 1:8 99.3 75 60 3 2:1 94.1 75 60 3 5:1 81.2 75 60 3 8:1 65.8

4. K I N E T I C M O D E L In order to determine the kinetics of reaction a first approach was made using a classic analysis. Kinetic data were obtained for all the t e m p e r a t u r e values tested, indicating a first-order relationship b e t w e e n reaction rate and catalyst concentration. Kinetic reaction order was also studied, founding a second order for global reaction, first order with respect to acid and first order with respect to alcohol. These results were in accordance with those obtained in previous work [7]. The kinetic models were as follows:

(--rAc) = 0.6318 exp(- 2843)Caly.CAc - 0.0742 exp(- 3036)CMono.CAr T T

(1)

2843. C (-r~,y) = 0 . 6 3 1 8 e x p ( - ~ ) ~,y.CAc

(2)

(rMono) =

2843. C 3036 0.6318exp(-~) ~,y.CAc- 0.0742exp(--~--)CMono.CAc

3036 (rDi) = 0.0742 exp(- ~)CMono.CAc

(3) (4)

330 In a second approach an enzymatic analysis was proposed for kinetic study. The m e c h a n i s m generally accepted for esterification reactions catalysed by immobilized lipases involves the formation of an acyl-enzyme intermediate, followed by the incorporation of an alcohol molecule to produce a molecule of ester and water. The most frequent inhibition type is the competitive inhibition. Competitive inhibition occurs when free enzyme can combine either with the substrate, to give the productive complex, or with an inhibitor, to give a non-productive complex, but not with both. The most common reason for this p a t t e r n is t h a t inhibitor is a substrate analogue, and sits precisely where the s u b s t r a t e should be, in the enzyme's active site. In some cases there is competition for the formation of a complex with the free lipase, and in other cases the competition is given by the formation of t e r n a r y complexes with the most active complex, enzyme-acid. Therefore, the general m e c h a n i s m of esterification corresponding to an ordered bi-bi system, including inhibition t e r m s for r e a c t a n t s and products, can be considered. In previous papers, we presented a kinetic model describing esterification reactions using immobilized lipases in organic solvent-free media. According to the general reaction scheme, the kinetic model can be obtained and simplified as a function of the studied s y s t e m characteristics. It was shown t h a t the highest conversions were obtained when the system operates in v a c u u m conditions. This fact allows the displacement of reaction equilibrium towards products formation. From an industrial point of view, the use of vacuum and 1/1 acid/alcohol molar ratio can be considered as d e t e r m i n a n t factors for product quality and, consequently, for i n d u s t r i a l scale processing. The general expression for the reaction rate can be simplified considering the esterification system as an irreversible process and concentration of diglyceride as neglected, as follows: (--rAr

=

(--rAr K

iAc

K tool,, "

1 -~-

C Gh,

..... C A~ C OJy

C Mono -- at- .... , K iMono

+ Co,,KmA~ "

1+

"1- C Ac

---~-+ K iGly

K

mGly

, K

1+

+

,

K + CA~

iAc

\ C Mono / + ,, K iMono

)

G~y

iMono

(5) For a 1/1 molar ratio acid/alcohol: CAcO "- C Gly0 "'" X Ac0 ---- X Gly0 C Ac0 . X Ac0 -- C AcO . X GlyO C A e "- C G I y -- C A t 0 . ( 1 - X A c ) C Mono :

C Ac0 . X Ac

(6)

331 And b o u n d a r y conditions: t--0

.'. CAr162

"'" X A r

t--t

.'. CAr

.'. XAr

(7)

(CA~o-CAr )/CAr o

By i n t e g r a t i n g the general equation the following expression is obtained: 1

1

t =or + [3XAr + 7.Ln . . . . . . +5 ......... (1--XAc) (1-XAc)

(8)

where Greek symbols are the coefficients defined by m a t h e m a t i c a l relations which involve the kinetic constants. The entire development of the kinetic model can be found in previous work [8] [9].

5. D I S C U S S I O N P a r a m e t e r s were e s t i m a t e d from values of e x p e r i m e n t a l rates by applying numerical methods based on M a r q u a r d t algorithm. Table 2 shows the p a r a m e t e r s obtained for the three t e m p e r a t u r e values tested.

Table 2. P a r a m e t e r values obtained for different t e m p e r a t u r e s . Constant T = 70~ T = 75~ (-rhc)max/E0 (mol.l-l.min -1) 0.0281 0.0351 KmAc (mol/1) 5.12 4.38 KinGly (tool/l) 5.41 4.52 Kiic (tool/l) 5.70 4.59 KiGly (mol/1) 3.75 2.67 K'ihc (mol/1) 0.832 0.651 K'iGly (tool/l) 1.204 0.986 S'iMono (mol/1) 0.719 0.582 K"iMono (mol/1) 0.381 0.264

T = 80~

0.0409 3.61 3.78 3.87 1.21 0.415 0.810 0.312 0.167

E x p e r i m e n t a l and calculated conversion values obtained for each e x p e r i m e n t were plotted. Curves were obtained for all t e m p e r a t u r e s and enzyme concentration tested. The a g r e e m e n t between conversion values m e a s u r e d experimentally and those calculated from the model indicates t h a t the kinetic model proposed is adequate to describe the synthesis of monoglycerides of ricinoleic acid over enzymatic catalysts. Residual analysis was used in order to evaluate involved errors using the proposed kinetic model. The results of this analysis indicated convenient fits, founding an average error of 9.79% and a m a x i m u m error of 20%.

332 Experimentally, best results for monoglyceride selectivity (97.2%) were obtained at 80~ of reaction temperature, 60mmHg of pressure and 3% of initial concentration of catalyst when acid/alcohol molar ratio was fixed at 1. At this point, maximum ricinoleic acid conversion was only found at 62% - 65%. Obviously, increasing immobilized lipase concentration and maintaining t e m p e r a t u r e level can increase acid conversion. However, this can produce the apparition of diglycerides, which can be translated as a decrease of selectivity and a poor product quality.

6. N O M E N C L A T U R E Ac: F a t t y acid (ricinoleic acid) CA: Concentration of component A (mol/1) Di: Diglyceride of ricinoleic acid E: Free enzyme Gly: Glycerol Keq: Equilibrium constant KiA: Inhibition constant for component A (mol/1) KmA: Michaelis-Menten constant for component A (mol/1) NA: Number of moles of component A (mol) Mono: Monoglyceride of ricinoleic acid P: Pressure (mm Hg) (-rA): Reaction rate for component A consumption (mol. 1-1.gcat-l.min-1) t: Reaction time (min) T: T e m p e r a t u r e (~ W: Water w: Amount of catalyst (%) XA: Conversion of A component

7. R E F E R E N C E S

1 M. Berger and M.P. Schneider, JAOCS, 69 N ~ 10 (1992) 683 2 P. Eigtved, T.T. Hansen and C.A. Miller, Proceedings of the World Conference on Biotechnology for the Fats and Oils Industry. T.H. Applewhite, AOCS, 1988. 3 J.A. Arcos and C. Otero, JAOCS, 73 N ~ 6 (1996) 673 4 C. Waldinger and M. Schneider, JAOCS, 73 N ~ 11 (1996) 1513 5 G. Boswinkel, J.T.P. Derksen, K. Van't Riet and F.P. Cuperus, JAOCS, 73 N ~ 6 (1996) 707. 6 T. Garcia, M. Martinez and J. Aracil, Enz. Microb. Technol. 15 (1993) 607 7 T. Garcia, M. Martinez and J. Aracil, Trans IChemE C 71 (1993) 47. 8 T. Garcia, A. Coter6n, M. Martinez and J. Aracil, Chem. Eng. Sci. 51 (1996) 2841. 9 T. Garcia, A. Coter6n, M. Martinez and J. Aracil, Trans IChemE C. 73 (1995) 140

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

Kinetic Modeling Elementary Steps

333

of Paraffins Hydrocracking based and the Single Event Concept.

G. Martens and G.F. Froment* Laboratorium voor Petrochemische Universiteit Gent, Belgium

upon

Techniek

Abstract

The kinetics of paraffins hydrocracking are developed in terms of elementary steps and single events,thus substantially reducing the number op parameters with respect to an adequate lumping or molecular approach. Plausible assumptions and thermodynamic constraints further reduce the number of independent rate coefficients to i0. These were determined from an experimental program with judiciously chosen Cs-paraffins.

i.

INTRODUCTION

So far the kinetic modeling of catalytic processes involving complex feedstocks and a large number of reactions between the various species has been based upon drastic simplifications yielding a small number of lumps and a limited number of reactions between them. Generally a lump consists of a large number of components belonging to a certain class of hydrocarbons. It is further defined by a physical property, like the boiling range or the density. Lumping is practiced in the modeling of the hydrocracking of vacuum gas oil or of wax produced by the Fischer-Tropsch synthesis. The reactions of the various lumps are described in terms of those of key components, thus inevitably leading to a biased product distribution and to rate coefficients which depend upon the composition of the lump, i.e., which are not invariant neither with respect to the feedstock composition nor to the operating conditions. Consequently a continuous experimental effort is required to support the kinetic model and its application. Froment and co-workers [1..3], on the other hand, excluded all lumping from the reaction network generation. Lumps are only introduced to account for the feedstock compositions, which is not very detailed with today's analytical tools and to set the partial pressures entering into the rate equations for the initial reaction steps. The lumps contain all possible components corresponding to their definition. In a judiciously selected lump the components should be at equilibrium. If the reaction network and the rate equation were based upon molecular structures the kinetic model would clearly comprise a gigantic number of rate coefficients since the rate coefficient for the h y d r o i s o m e r i z a t i o n of paraffins varies with the number of C-atoms in the molecule. To avoid this a more pronounced mechanistic description of the reaction scheme in terms of elementary steps is required. Hydrocracking catalysts contain a metal component (Pt, Pd,NiMo,CoMo,NiW) and an acid component, generally zeolite Y or alumina. The metal loading is sufficient to ensure that the rate determining step is located on the acid sites. The elementary steps on these sites are written in terms of carbenium ion chemistry. 2. O C T A N E

HYDROISOMERIZATIONAND

HYDROCRACKING

The various types of elementary carbenium ion steps o c c u r i n g in paraffins h y d r o i s o m e r i z a t i o n and hydrocracking are shown in Figure 1 for noctane. The elementary steps are protonation and d e p r o t o n a t i o n , h y d r i d e shifts,alkyl shifts,branching isomerization through a protonated c y c l o p r o p a n e intermediate and cracking through ~-scission. ,Present address: Department University, College Station, Texas

of Chemical 77843-3122,USA

Engeneering,Texas

A&M

334

+

§

i li +

!i-!

Craddng

§

11o 11_..

Figure i. Elementary and hydrocracking.

carbenium

3. R A T E

THE

EQUATIONS

FOR

ion

steps

ELEMENTARY

STEPS

in

paraffins

hydroisomerization

In hydrocracking at low temperatures and on a Pt-USY zeolite necessary to explicitly account for the physisorption of the feed nents [4]. This is preferably done through a Langmuir-isotherm:

it is compo-

Ku Pi

CPi - Csae I + ~ KLI Pl

(i)

1

Dehydrogenation

into olefins Pi(ads)

The olefins

are protonated

takes place on the metal:

= Oij + H 2

with

on Bronstedt Oij(ads)

KDH'iJ"

acid sites,

+ H § . RII §

Co~j .px~

(2)

Cp~

yielding

carbenium

ions (3)

335 which isomerize through on the a c i d sites:

hydride-

, methyl-shifts

and

PCP-branchings,

still

Ril ~ - Rno § Cracking

also

occurs

on these

sites,through

(4) ~-scission:

(5)

Ril § - R;x/ + Or~(ads)

The steps on the a c i d sites are rate d e t e r m i n i n g so that the net rate of f o r m a t i o n of the p a r a f f i n s is the s u m m a t i o n of the net rates of f o r m a t i o n of the o l e f i n s :

Rp i - S The o l e f i n s that

are

involved

Roi~

(6)

in p r o t o n a t i o n s , d e p r o t o n a t i o n s

Rol j = E k D e ( 1 ; O i j ) 1

CRil.- E k p r ( 1 ) O o i j C H § j

and

+ ~kcz(Y;a,Oij) y

~-scissions,so

CR ,

(7)

The c a l c u l a t i o n of the c o n c e n t r a t i o n of the c a r b e n i u m ions a p p e a r i n g r e q u i r e s in the first p l a c e an e q u a t i o n for t h e i r rate of f o r m a t i o n :

3

in

(7)

1

i

o

(8)

+ ~ kcr (U; 2 ,o~) c,w- ~ kcr (2; q, o~)c,. u

In the p s e u d o

steady

q

state RaiI. : 0

Eqns acid

(8) and sites

(9),

together

with

a balance

Ct = CH" + $ $

form

a set

of

linear

equations

in

(9) on the

concentration

of B r o n s t e d t

CRn"

CH., CRi I. .Its s o l u t i o n

(i0) yields

to be s u b s t i t u t e d into the net rate of f o r m a t i o n of olefins, (7). S v o b o d a et al. [i] h o w e v e r d e r i v e d from t h e i r k i n e t i c s t u d y of c r a c k i n g that (I0) r e d u c e s to Ct=CH+. 4.

REDUCTION

OF THE

NUMBER

OF

the v a l u e s Cs-hydro-

PARAMETERS

The network of C8 h y d r o i s o m e r i z a t i o n and hydrocracking contains 22 paraffins (17 o c t a n e s and 5 p r o d u c t s of cracking), 75 o l e f i n s (66 o c t e n e s and 9 p r o d u c t s of c r a c k i n g ) , and 57 c a r b e n i u m ions. To r e d u c e the n u m b e r of rate p a r a m e t e r s a n u m b e r of a s s u m p t i o n s has to be i n t r o d u c e d . 1 - O n l y s e c o n d a r y a n d t e r t i a r y c a r b e n i u m ions are c o n s i d e r e d , since m e t h y l and p r i m a r y ions are far less stable. 2-For a g i v e n type of R § (s or t) the rate c o e f f i c i e n t s are i n d e p e n d e n t of the c h a i n length. 3-The v a r i o u s k'isom o n l y d e p e n d on the type of r e a c t a n t ion (s or t) and p r o d u c t ion. 4-The v a r i o u s k'cr~k o n l y d e p e n d on the type of r e a c t a n t and p r o d u c t ion, not

336 on the p r o d u c e d olefin. 5-k'pr o n l y d e p e n d s on the type of R +. 6 - k ' m d e p e n d s on the type of ion and on the p r o d u c e d olefin. 5.

SINGLE

EVENT

RATE

COEFFICIENTS

In the above a s s u m p t i o n s the effect of the s t r u c t u r e of the c a r b e n i u m ions of a g i v e n type is not considered. To a c c o u n t for this effect B a l t a n a s et al. [2] i n t r o d u c e d the "single event" concept. A c c o r d i n g to the t r a n s i t i o n state t h e o r y the rate c o e f f i c i e n t s of the t r a n s f o r m a t i o n of a r e a c t a n t into the t r a n s i t i o n state c o m p l e x can be written:

k8

T

A S o*

k I = ~exp(

The

rotational

contribution

for the effect of following relation

the

R

to

AS ~

structure

on

)exp(-

was

A H 0*

RT

)

considered

k' .Baltanas

et

(ii)

to be al.

representative

[2]

derived

k / - ne.k

with

ne -

a~'z agl, #

the

number

of

single

the (12)

events

in

an

elementary

the rate coeffient of the single event. The global symmetry number

step

and

k

a~, r a n d o~,,

of the r e a c t a n t and the t r a n s i t i o n state also a c c o u n t for chirality. The c a l c u l a t i o n of ~ r e q u i r e s the c o n f i g u r a t i o n of the reactant, but also of the t r a n s i t i o n state. Recent q u a n t u m c h e m i c a l software, in p a r t i c u lar the ab initio versions, a l l o w the c o n f i g u r a t i o n to be d e t e r m i n e d in a r e l i a b l e way. 6.

THERMODYNAMIC

CONSTRAINTS

ON

THE

RATE

COEFFICIENTS

W i t h the a p p r o a c h b r i e f l y p r e s e n t e d here the n u m b e r of rate c o e f f i c i e n t s for p a r a f f i n s is r e l a t i v e l y small: 2 for o l e f i n s p r o t o n a t i o n ( one for sec R +- and one for tert R + - f o r m a t i o n ) , 4 for h y d r i d e shifts (kHs(S;s),kHs(S;t),kHs(t;s),kHs(t;t)),4 for m e t h y l - s h i f t s , 4 for P C P - i s o m e r i z a t i o n and 4 for ~scission. G i v e n the large n u m b e r of o l e f i n s the n u m b e r of i n d e p e n d e n t deprotonation coefficients is v e r y large, but this can be d r a s t i c a l l y r e d u c e d by i n t r o d u c i n g t h e r m o d y n a m i c constraints. For d e p r o t o n a t i o n the f o l l o w i n g r e l a t i o n was d e r i v e d by V y n c k i e r and Froment [3] :

kDe (m; Oij) i kD e (m, O r) fisom (Or"Oij)

(13)

where m stands for sec or tert and Or for a r e f e r e n c e o l e f i n isomer with the double b o u n d in such a p o s i t i o n that b o t h a s- or t-R + can be f o r m e d by protonation. Consequently, for a g i v e n c a r b o n n u m b e r there are o n l y two independent deprotonation rate coefficients: k ~ ( s ; O r) and km(t;Or). For octane h y d r o c r a c k i n g this r e d u c e s the k m from 85 to 7 ( 2 for C 8, 2 for C 5, 2 for C 4 and 1 for C 3 ). A l o n g s i m i l a r lines the f o l l o w i n g c o n s t r a i n t s were d e r i v e d for the k's of the three types of i s o m e r i z a t i o n :

k.s(~;s)

k.s(t;s)

k~cp(t;s) kp~(s)k~e(t;O)

k~zs(S; t)

kMs(S; t)

kpcp(S; t)

kpr( t) kDe(S;O)

(14)

337 This r e l a t i o n r e d u c e s the n u m b e r of i n d e p e n d e n t rate c o e f f i c i e n t s for HS-, MS- and P C P - i s o m e r i z a t i o n from 4 to 3. From the s e c o n d a s s u m p t i o n m e n t i o n e d above it follows that in ( 1 4 ) the ratio k ~ ( t ; O ) / k ~ ( s ; O ) h~s to be i n d e p e n d e n t of the olefin, e v e n w i t h a d i f f e r e n t n u m b e r of C-atoms. In o c t a n e h y d r o c r a c k i n g this f u r t h e r reduces the n u m b e r of k ~ from 7 to 5 (2 for C8, 1 for C 5, 1 for C 4 and 1 for C 3 ions). The c o n c e p t p r e s e n t e d here p e r m i t s the single event rate c o e f f i c i e n t s to be d e t e r m i n e d from h y d r o c r a c k i n g of short paraffins, w h e r e b y the c o m p l e t e p r o d u c t d i s t r i b u t i o n can be m e a s u r e d in a c o n v e n i e n t w a y e.g. by gas c h r o m a t o g r a p h y . To access all rate c o e f f i c i e n t s in an u n a m b i g u o u s way the c o m p o n e n t s to be c r a c k e d have to be j u d i c i o u s l y chosen. 7.

SELECTION

OF

REPRESENTATIVE

FEED

COMPONENTS

The f o l l o w i n g c o n s i d e r a t i o n s led to the final s e l e c t i o n of c o m p o n e n t s to be c r a c k e d : Previous w o r k did not s u c c e e d in d e t e r m i n i n g kMs(t;t) and ~r(t;t) [i]. In these e x p e r i m e n t s u s i n g n-octane, 2-Me-heptane and 2 , 5 - d i M e - h e x a n e the amount of t r i b r a n c h e d p a r a f f i n s was too low to be d e t e c t e d b e c a u s e of the fast (t;t) cracking. To d e t e r m i n e the above m e n t i o n e d rate c o e f f i c i e n t s 2,3,4 t r i M e - p e n t a n e was chosen. A m b i g u i t i e s were e n c o u n t e r e d for k~p(t;t). This single event rate coefficient is i n v o l v e d in the v a r i o u s e l e m e n t a r y steps t r a n s f o r m i n g di- into t r i b r a n c h e d R + and mono- into d i b r a n c h e d R +. Each of the t r a n s f o r m a t i o n s mentioned above contains next to PCP(t;t) also parallel PCP(t;s) or PCP(s;t) steps, so that the a s s o c i a t e d rate c o e f f i c i e n t s c o u l d not be i n d e p e n d e n t l y d e t e r m i n e d . There is no such a p r o b l e m w i t h the f o l l o w i n g four (t;t) P C P - s t e p s w h i c h t r a n s f o r m 2,4 d i M E - h e x a n e into 2,2,3- and 2,3,3triME-pentane :

W i t h a Cs-feed k~p(t;t) can o n l y be o b t a i n e d from the b e h a v i o u r of 2,4diMehexane, 2,2,3- and 2 , 3 , 3 - t r i M e p e n t a n e . F i n a l l y 2 , 3 , 4 - t r i M e p e n t a n e was s e l e c t e d since this c o m p o n e n t is not r a p i d l y t r a n s f o m e d b y (t;t) c r a c k i n g and thus also y i e l d s i n f o r m a t i o n on (t;t) a l k y l s h i f t s w i t h i n t r i b r a n c h e d paraffins. 8.

EXPERIMENTAL

PROGRAM

AND

RESULTS

The c a t a l y s t u s e d was a U S - Y zeolite c o n t a i n i n g 0.5 wt% Pt. The r e a c t o r was a B e r t y r e a c t o r for gas p h a s e operation. The t e m p e r a t u r e r a n g e d from 200 to 260 ~ the p r e s s u r e s from I0 to 50 bar, the r a t i o ~/HC from 30 to I00 and the space time from i0 to 420 gcat.s/mol. The e f f l u e n t c o m p o s i t i o n was a n a l y z e d by m e a n s of an o n - l i n e GC. The e x p e r i m e n t a l p r o g r a m c o n s i s t e d of : 201 e x p e r i m e n t s w i t h p u r e n-octane, 35 e x p e r i m e n t s w i t h a m i x t u r e of 90 mol% n - o c t a n e and i0 mol% 2 - m e t h y l p e n t a n e , 33 e x p e r i m e n t s w i t h a m i x t u r e of 90 mol% n - o c t a n e and i0 mol% 2,5 d i m e t h y l h e x a n e and 57 e x p e r i m e n t s w i t h a m i x t u r e of 90 mol% n - o c t a n e and i0 mol% 2,3,4 t r i m e t h y l p e n t a n e . The product distributions at various conversions confirm previous o b s e r v a t i o n s : e q u i l i b r i u m is r e a c h e d b e t w e e n the v a r i o u s i s o m e r s w i t h i n the mono- and d i b r a n c h e d families. This is not the case w i t h i n the t r i b r a n c h e d family because of the rapid (t;t) ~-scission [5]. An example of a c o n v e r s i o n v e r s u s space time curve is g i v e n in Figure 2 for n-octane.

338

K

T=S33

T=513

5o

c:

40

"-

3o

K

o 20

Io

o

I so

I

I

10o

15o

Space

Figure

2.

Conversion

9. P A R A M E T E R

of n - o c t a n e

I

I

I

200

~5o

Soo

t l me ( O c a t .

versus

space

850

sl tool )

time.

ESTIMATION

T h e n u m b e r of p a r a f f i n r e s p o n s e s p e r e x p e r i m e n t was 18 (4 less t h a n the n u m b e r of p a r a f f i n s in the r e a c t i o n n e t w o r k b e c a u s e of the o v e r l a p p i n g of 3 - M e - C 7 w i t h 3 - E t - C 6 a n d 3-Et-3-Me-Cs; 4 - M e - C 7 w i t h 3,4-diMe-C6; 2,3-diMe-C 6 with 2-Me-3-Et-Cs). Since replicate experiments were performed a weighted least squares objective f u n c t i o n was used. Its m i n i m i z a t i o n i n v o l v e d two techniques i) R o s e n b r o c k in the e a r l y s t a g e s 2) M a r q u a r d t in the final stage. Where well established rules of carbenium ion chemistry were available these were introduced as c o n s t r a i n t s on r a t i o s of r a t e c o e f f i c i e n t s of s o m e e l e m e n t a r y steps. The s t a t i s t i c a l t e s t s on the fit a n d on the p a r a m e t e r s [6] w e r e s a t i s f i e d . P a r i t y p l o t s for some of the r e s p o n e s are s h o w n in F i g u r e s 3 a n d 4. T h e p a r a m e t e r values, a l s o that of the physisorption c o n s t a n t of n - C 8 i s o m e r s are g i v e n in T a b l e I. Table 1 Composite

Arrhenius

and van

't H o f f

parameters

for C 8 h y d r o c r a c k i n g

Parameter

k'0 (mo i / (goath) )

E*0 (kJ/mol)

k'ms(S; S)

0.18

48.7

i0 I0

k*Ms (S ; t ) = k*Ms (t ; S )

0.33

i0 I~

48.1

k*MS(t ;t)

0.43

1011

47.9

k*~:p (s ; s )

0.61

i0 v

48.5

k*~zp (s ; t ) = k*pcp (t ; s )

0.86

108

48.0

k*~p (t ; t )

0.21

108

47.7

k*c~k (s ; s )

0.12

i0 I0

70.9

k*cr~k (S ; t )

0.45

106

17 .i

k'cr~k (t ; S )

0.88

1012

77.3

k*er~k (t ;t )

0.15

109

Parameter

K 0 (bar I)

AH~s k J / m o l

KL,8

8.3

74.8

10 .8

9.4

339 The rate coefficients k + in this table are products of the single event rate coefficient and the equilibrium constant for p r o t o n a t i o n / d e p r o t o n a t i o n [6]. The k+0 are really products of Csm,Ct,A*pr=A'~/A'mp and A'imm or A'cr+k. The prime indicates that the frequency factor of the elementary step does not include the entropy contribution due to the change in global symmetry since this is already accounted for by means of the number of single events, ~. E"0 in Table 1 is the algebraic sum of the p r o t o n a t i O n enthalpy and the activation enthalpy of the elementary step. The k'ms(S;S) and k'Ms(S;t) may contain a minor contribution of (s;s) and (s;t) ethyl shifts because of the overlapping of the GC peaks of ethyl- and methyl paraffins.

0.0of,

,.**i

o

o.,,,, o.,,,,

o.o,,, ,.,,,l

o

o

r

,.i,4

,.,,.

,. ,,,a ,.ooo,

,.1,i

,,,1,+

,.1,,i

::::::

o

o 0.***i

,.,**+

,.1,**

1.,**,

,.,,,

,.,,i

,.,,1+

1.1,+,

,.,,,,

Figure 3. Observed and calculated responses for 2,4-dimethylhexane

i0.

~176

i.iii

Figure 4. Observed and responses for propane.

calculated

CONCLUSION

A complete set of single event rate coefficients for paraffins hydroisomerization and hydrocracking has been obtained from an experimental program involving a number of judiciously chosen C 8 hydrocarbons. The set of parameters is statistically significant, satisfies the rules of carbenium ion chemistry and leads to an excellent fit of the experimental data. The fundamental modeling adopted in this work ensures that the set of parameters is valid also for the hydroisomerization and hydrocracking of higher paraffins. ii.

REFERENCES

Svoboda G.D.,Vynckier E.,Debrabandere B. and Froment G.F., SingleEvent Rate Parameters for Paraffin Hydrocracking on a Pt/US-Y Zeolite,Ind. Eng. Chem. Res. 1995,34,3793 Baltanas M.A.,Van Raemdonck K.K.,Froment G.F. and Mohedas S.R., Fundamental Kinetic Modeling of H y d r o i s o m e r i z a t i o n and hydrocracking on N o b l e - M e t a l - L o a d e d Faujasites. i. Rate Parameters for Hydroisomerization, Ind. Eng. Chem. Res. 1989, 28, 899 Vynckier E. and froment G.F., Modeling of the kinetics of Complex Processes Based upon Elementary Steps. Kinetics and Thermodynamic Lumping of Multicomponent Mixtures, Astarita G., Sandler S.I., Elseviers Science Publishers, Amsterdam, 1991,p 131. Steijns M. and Froment G. F., H y d r o i s o m e r i z a t i o n and Hydrocracking. 3. Kinetic Analysis of rate data for n-Decane and n-Dodecane. Ind. Eng. Chem. Prod. Res. Dev. 1981,20,660. Martens J.A. ; Jacobs P.A. Conceptual Background for the conversion of hydrocarbons in Heterogeneous Acid Catalysis, Theoretical Aspects of Heterogeneous Catalysis, Moffat J.B. Eds. , Van N o s t r u n d Reinhold,

340 New York ,1998 Froment G.F., Bischoff E d , J o h n W i l e y & Sons,

12 N O M E N C L A T U R E A' = f r e q u e n c y factor contribution, hl

C H.

= concentration

Co~~

= surface

Cp~

= surface

of

of

elementary

free

concentration concentration

CR.

= surface

Csat

= concentration

C t = concentration

concentration of of

K.B., C h e m i c a l NY, 1 9 9 0

total total

active of of

step

reactor

not

acid

Analysis

including

sites,

and

Design,

symmetry

mol/gc~

O U, mol/gca t Pi,mol/gca t

of

Rik+ , mol/gc~

physisorption active

acid

sites, sites,

mol/gca t

mol/gc~

De = deprotonation DH = dehydrogenation ES = e t h y l s h i f t HD = hydrogenation k = single-event rate coefficient, hI k' = r a t e c o e f f i e c i e n t of e l e m e n t a r y step, h I k* = c o m p o s i t e single event rate coefficient, mol/(gc~.h) kcr(ml;m2, O U) = s i n g l e e v e n t r a t e c o e f f i c i e n t for cracking of a c a r b e n i u m i o n of t y p e m I to f o r m a c a r b e n i u m i o n of t y p e m 2 a n d a n o l e f i n Ou, h -I k ~ ( m , O U) = s i n g l e e v e n t r a t e c o e f f i c i e n t for deprotonation of a c a r b e n i u m i o n of t y p e m t o f o r m a n o l e f i n Ou, h -I kisom(ml,m 2) = s i n g l e e v e n t r a t e c o e f f i c i e n t for isomerization of a c a r b e n i u m i o n of t y p e m I to f o r m a c a r b e n i u m i o n of t y p e ~ , h ~ kpr(m) = s i n g l e e v e n t r a t e c o e f f i c i e n t for protonation of a n o l e f i n w i t h formation of a c a r b e n i u m i o n of t y p e m, h I KDH.U = d e h y d r o g e n a t i o n equilibrium c o n s t a n t of Pi to f o r m Oij, b a r KL.i = L a n g m u i r p h y s i s o r p t i o n c o n s t a n t of Pi, bar1 MS = methyl shift = n u m b e r of s i n g l e e v e n t s O U = olefin with index j formed from paraffin with index i Or = r e f e r e n c e olefin Pi = g a s - p h a s e partial pressure of Pi, b a r Pi = p a r a f f i n with index i Pr = protonation PCP = branching via protonated cyclopropane R§ = carbenium ion Rik§ = c a r b e n i u m ion with index k formed from paraffin with index i Rou = n e t r a t e of f o r m a t i o n of O~, mol/(gc~.h) Rpi = n e t f o r m a t i o n of Pi, mol/(gc~-h) RRik+ = n e t r a t e of f o r m a t i o n f o r Rik+, mol/(gc~.h) Acknowledqment The authors are the experimental

greatful data.

to

A.

Collier

and

B.

Debrabandere

for

providing

2

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science Ltd. All rights reserved.

341

A comparative kinetic study of CH 4 oxidation by NiO/A1203, Pt/A1203 and NiO-Pt/A1203 catalysts T.-N. Angelidis and V. Tzitzios Aristotle University, Faculty of Chemistry (Box 114), 54006 Thessaloniki, Greece

Abstract

A NiO-Pt/fi-A1203 oxidation-reduction catalyst (3%Ni and 0.1%Pt, C53-1, Catalyst and Chemicals Europe S.A.) is applied in industry (EKO refinery, Thessaloniki, Greece) for nitrogen gas production, through LPG combustion by atmospheric air. The aim of the present research work is to compare the kinetics of total oxidation of methane by the separate constituents of the above catalyst, as well as by the mixed catalyst. The ~-alumina based NiO (3%Ni), Pt (0.1%Pt) and NiO-Pt (3%Ni, 0.1%Pt) catalysts were prepared by the impregnation method and characterized by BET, XRD and SEM-EDX. The activity tests show that activity follows the order, NiO-Pt>Pt>NiO. The reaction kinetics were measured at low partial pressures of oxygen and methane near the stoichiometric point. The kinetic results were interpreted with the Langmuir-Hinshelwood (L-H) and Elye-Riedel (E-R) mechanisms. The reaction follows an E-R mechanism on the NiO catalyst. On the Pt catalyst, reaction follows a mixed L-H/E-R mechanism, but at low methane partial pressures and high oxygen partial pressures an E-R mechanism prevails. The reaction rate kinetic and adsorption constants and activation energies were calculated according to the above mechanisms. The validation of the proposed mechanisms was checked statistically and through thermodynamic consideration. The reaction rate over the Pt catalyst is considerably higher (ten times) than over the NiO catalyst at the stoichiometric point. The rate of the reaction over the mixed catalyst is the sum of the rates over the separate catalysts and a synergistic effect is not seems possible.

1. INTRODUCTION An oxidation/reduction catalyst (3% Ni and 0.1% Pt on a-alumina, C53-1, Catalyst and Chemicals Europe S.A.) is applied in industry for Nz(g ) and CO2(g) production. The catalyst is described in detail elsewhere [1 ]. The flowsheet of the production unit is shown in Figure 1. A desulfurization ZnO catalyst (XC-117) is applied before the main catalyst to prevent poisoning by sulfur. The above catalyst may be of interest in environmental applications for removal of NOx, CO and hydrocarbons from combustion exhaust gases, due to his oxidation/reduction properties. The presence of nickel will limit his application only to stationery sources, where preheating is easier, to prevent nickel carbonyl formation (nickel carbonyls are stable up to 200~ [2]). The kinetics of the deNOx action of this catalyst has been studied in the earlier stages of his development [3].

342

HEATER ~ CATALYTIC ELECTRICAL

A: CATALYSTXC-117 B: CATALYSTC-53-1 GASES

- MEA

;;%

NITROGEN

315"Cl ~lR LPG~:~ REBOILER~ I I

AIR

ilI

Ico, I

I

ABSORBER . . . . . J'~

[ 'I

MEA

]

g_ _ ~ Figure 1. The industrial unit flowsheet (EKO refinery and petrochemical plant, Thessaloniki, Greece). The main effort of the present research was to examine the kinetics of the oxidation action of the separate constituents of the above catalyst and of the mixed one during methane oxidation. Methane was selected, since is the most refractory hydrocarbon and thus, is often used as model hydrocarbon compound for activity tests. In addition, methane itself is a potent greenhouse gas and the emission control of unburned methane from exhaust gases (specially from natural gas combustion) is essential. High load nickel based catalysts are applied in SYNGAS reforming by partial oxidation of methane and extensive research work is published concerning the improvement of the process conditions [4-5]. The application of nickel based catalysts in CO2 reforming by methane over nickel based catalyst is also the subject of recent publications [6-9]. Pt based catalysts are extensively applied in oxidation of CO and hydrocarbons from combustion exhaust gases [ 10]. The generally proposed mechanism for methane oxidation over noble metals (Rh, Pd, Pt) as well as transition metal oxides [ 1l-13] can be summarized as follows: CH4(g )

HCHO(g)

CO(g)

;T CH4(a ) - tt ~

C t t 3 . (a) + O

~

ttCItO(a)

H2(g )

T; ~

CO(a) + 2H(a) + O

----*CO2(g) + H 2 0 ( g )

In both cases the formation of adsorbed methyl radicals is the rate controlling step. The main differences between noble metals and transition metal oxides are: 9 methane is adsorbed on the noble metals surface at reasonable partial pressures, while is not adsorbed at all on metal oxides in the same partial pressures region [ 13] 9 oxygen is readily adsorbed and dissociates on noble metals catalysts, while is not easily adsorbed on metal oxides. In the last case oxidation may be under mixed control (oxygen adsorption and adsorbed methyl radicals formation) [13] 9 in the case of nickel oxide oxygen is readily adsorbed and dissociates due to the p-type semiconductor' s characteristics of the oxide [ 12-13]. According to the Langmuir-Hinshelwood (L-H) and the Eley-Riedel (E-R) approach, the acceptance of the above statements leads to the conclusion that an E-R mechanism prevails (methane from the gas phase reacts with adsorbed oxygen at substantial partial pressures of

343 oxygen). This is what most researchers observed during the kinetic study of methane oxidation over noble metals and active metal oxides catalysts [ 14-15].

2. E X P E R I M E N T A L 2.1. Catalyst Preparation and Characterization The supported catalysts' samples with composition similar to that of the industrial catalyst (3% Ni in NiO/A1203, 0.1% Pt in Pt/A1203 and 3%Ni-0.1%Pt in NiO-Pt/A1203) were prepared by conventional wet impregnation. Hydrogen hexachloroplatinate (IV) hydrate (Riedel-deHaen) and nickel nitrate hydrate (Riedel-de Haen) were applied to provide the active metals and fi-Al203 (HARSHAW A1-3971P) was used as the support material. The required quantities of metals and support were mixed with water. Water was removed by heating under agitation. The resulting solids were dried at 105~ for 24h, crushed and sieved, the 75-212im fraction being retained. They were then heated to 600~ in a current of hydrogen (2h) and in a current of helium (lh). The catalyst characterization was performed by AAS (PERKIN ELMER 2380) for bulk chemical analysis, by nitrogen adsorption/desorption for BET surface and pore volume measurements, by XRD (SIEMENS D-500 CuKfi) for crystalline identification and by SEM-EDX (JEOL 120CX, LINK AN 10S, ZAF4) for the texture and composition dispersion of the catalyst particles. Bulk analyses show that the prepared catalysts were of the expected chemical composition. XRD studies show that alumina exists mainly in gamma crystalline form and nickel mainly as NiA104. SEM-EDX studies show that the NiO-Pt/AI203 particles have a more clear crystalline form that confirms their low porosity compared with the other catalysts and that there is no observable difference of chemical composition between particles. 2.2. Activity and kinetics measurements The reactor was a 1cm i.d. x 35 cm long quartz tube heated by a temperature controlled tubular furnace. The reaction temperature was monitored by a thermocouple placed near the packed catalyst bed. Oxygen and methane certified calibration gas mixtures balanced by helium were used as reacting gases and pure helium was used as diluent (all from AIR LIQUID). The gas streams were measured with mass flow controllers and mixed before the reactor inlet. The resulting gas mixture flowed through the packed bed. For kinetic measurements, the reactor was operated in a differential mode with the conversion not exceeding 10%, so that the temperature was nearly uniform in the packed catalyst bed. Separate experimental tests showed that bulk mass transfer and intraparticle mass transfer resistance could be eliminated by using a gas space velocity greater than 26500 h ~ and catalyst particles less than 212im in size. The reaction conversion was controlled by the catalyst loading. The partial pressure of the reacting gas species was varied over the range of 0.01-0.08 bar for the Ni catalyst and 0.01-0.05 bar for the Pt and mixed catalysts and the temperature was varied over the range 525-575~ and 475-550~ respectively. Before any kinetic measurement, the catalysts were always treated with a gas flow of hydrogen for 0.5h and helium for 0.5h at 600~ The chemical analysis of reactants and the products gas mixtures was performed by gas chromatography (SHIMADZU GC-14B) equipped with Poropac Q and Molecular Sieve 5A columns and a TCD (Thermal Conductivity Detector). The mass balance of the reaction was always checked by analysis of all the feed and the product gases (O2, CH4, CO2 and CO). Carbon monoxide production was only observed during the activity experiments at temperatures over 800~ So the production of CO: was used to calculate the reaction rate:

344 (1)

Rate(cH4) = Rate(c02) = N t X c o 2 / S

where N t is the total molar gas flowrate in mole sec-1, Xco2 is the molar fraction of C O 2 in the product gas stream and S is the surface of the considered catalyst (calculated from the specific surface and the weight of the applied catalyst sample).

3. RESULTS AND DISCUSSION 3.1 Activity tests Figure 1 shows the light-off curves of C H 4 oxidation over the three prepared catalysts and the industrial one. The experiments were carried out at the stoichiometry of the methane total oxidation reaction (PcHn/Po2 = 89 The calculated light-off temperatures were 780, 610, 560 and

CONVERSION

CH 4 , %

CONVERSION

CH 4, %

100

I N iO/A

80

I i2 0 3

m

m TOTA L T O CO

60

_ __

-t

40

_

TO

CO 2

NiO/AI2 03

-~ / ..~ / NiO-Pt/AIz 031

20

P t / A l : O~

INDUSTRIAL] .

_

O

600 TEMPERATURE,

800 oC

1000

Figure 2. Light-off of C H 4 oxidation over the tested catalysts (PCH4 = 0.01 bar and Po2 = 0.02bar)

400

u=

--

1

.

.

.

.

6~) . . . . . TEMPERATURE,

.

myd~ljm oC

1000

Figure 3. Light-off of C H 4 oxidation over the NiO/A1203 catalyst with CO formation (PCH4 = 0.01 bar and Po2 = 0.02 bar).

570~ for NiO, Pt, Pt-NiO and industrial catalyst respectively. The oxidation activity of the mixed and Pt catalysts is considerably higher than the respective activity of the NiO catalyst. So, the oxidation activity of the mixed catalyst is mainly attributed to the presence of Pt. The prepared in laboratory mixed catalyst and the industrial one seems to have comparable activities, and the results obtained by the first one are representative for the later. In Figure 2 the partial oxidation of methane to CO at temperatures over 800~ is shown. CO production at high temperatures disappears in the presence of Pt in the mixed and industrial catalyst. 3.2. Kinetic results

Figures 4 and 5 show the variation of C H 4 oxidation rate with the partial pressure of 02 and respectively. Various rate equations derived from different reaction mechanisms as well as

CH 4

345 rate, moleCH4/sec

m2

rate, moleCC4/sec

NiO/AI2 O3 TEMPERATURE 8E- l0

~

5o0" C

ll

52s,.c

O

540" C

9

550" C

[

m 2

3E-9 NiO/AI2 03 TEMPERATURE

2E-9

_

~

5oo, c

1

525" c

O

540" C

9

sso,, c

9

-

4E-10 1E-9

0 0.00

I 0.08

0.04

0.12

pO 2 , bar

Figure 4. Variation of the C H 4 oxidation rate over the N i O / A 1 2 0 3 catalyst with Po2 (PCH4 = 0.01 bar).

_

0 0.00

I 0.08

0.04

0.12

p U l l 4 , bar

Figure 5. Variation of the C H 4 oxidation rate over the NiO/A1203 catalyst with PCH4(1902 = = 0.02 bar).

the empirical power order equation were evaluated to regress the experimental data. It was found that the experimental data were best represented by the equation: RCH 4 --

k PfH4KoPo/(1+KoPo)

(2)

where R C H 4 is the methane oxidation rate, k and Ko can be considered as the surface reaction rate constant and the adsorption equilibrium constant of oxygen respectively. Eq.1 gives a linear relation between the methane oxidation rate and the partial pressure of methane for constant partial pressure of oxygen. The results of the statistical check of the experimental data were applied for the calculation of the constants k and Ko. The curves in Figures 4 and 5 represent the fitting of Eq. 1. Eq. 1 is representative of an E-R mechanism. According to the discussion of the possible mechanisms made in the introduction the formation of adsorbed methyl radicals is the rate controlling step and methane reacts from the gas phase with adsorbed oxygen. The activation energy of the surface reaction was found to be 36.5 kcal mole l from the Arrhenius equation. The heat of adsorption calculated from the Ko values at various temperatures was found to be 11.4 kcal mole -I and the respective entropy of adsorption (AS) -6.2cal mole-lK 1. The entropy value seems to be in agreement with the fundamental thermodynamic rules: A~S < 0 and [AS] < Sg where Sg is the entropy of oxygen gas (56,5-57.2 cal mole l K -1 at the experiments temperature range [16]). The entropy value does not agree with the Boudart rules: [A,S[ > 10 and [AS[ < 12.2-(0.0014Q) where Q is the heat of adsorption in cal mole 1. This means that the calculated heat of adsorption is low and the dependence of the respective adsorption constants on temperature is less than thermodynamically expected. This fact may be explained by crystalline changes of nickel compounds from NiA104 to NiO as temperature increases. The later is a more active oxygen adsorbent [4] and changes the adsorption behavior of the catalyst. Figures 6 and 7 show the variation of the methane oxidation as a function of the partial pressure of methane and the partial pressure of oxygen respectively. A mixed E-R/L-H mechanism described by Eq.3 was the basis for the discussion of the kinetic behavior:

346

RCH 4 --

(3)

klKoPoPcH4/(1 + KoPo) + k2KoPoKcH4PcH4/(1 + KoPo)2 Eley- Riedel Langmuir-Hinshelwood

Where k~ and k 2 are reaction rate constants and Ko and KCH 4 a r e the respective adsorption constants of oxygen and methane. For constant partial pressure of oxygen Eq.3 becomes a linear function of the partial pressure of methane, while at low partial pressures of methane only oxygen is adsorbed and the E-R mechanism prevails. This consideration is confirmed by the experimental data of Figure 6, since they can be analyzed by two separate linear functions with different slopes. The same happens with the experimental results of Figure 7, where at low partial pressures of oxygen, methane is partially adsorbed and the mixed mechanism prevails and at higher partial pressure of oxygen the E-R mechanism prevails and the rate becomes constant. At the stoichiometric point both figures show that the reaction follows an E-R mechanism that is described by the linear relation: RCH 4 --

(4)

klPcH 4

since KoPo2 >> 1. The values of the reaction rate constant were calculated at various temperatures from the first linear part of Figure 6 and the activation energy was found to be 33 kcal mole -1 (this value is in accordance with the activation energies found by other researchers [17-19] for methane oxidation over Pt catalysts, which were between 24 and 44.7 kcal mole1). The lack of experimental data, specially in the experimentally difficult range of low partial pressure of oxygen, does not permit a complete kinetic study. Figures 8 and 9 show the kinetic behavior of the two separate catalysts compared with the mixed one. The oxidation rate on the mixed catalyst seems to be the sum of the rates of the separate catalysts. A definite conclusion is difficult to be obtained since the oxidation rate of methane on the nickel catalyst is about ten times lower than the oxidation rate on the platinum catalyst. rate, m o l e C H 4 / s e c

rate, m o l e C H 4 / s e c

m 2

1.6E-8

6E

I

t Pt/AI 20j i TEMPERA TURE

~/

1.2E-8

ll

475; C

O

500" C

9

m 2

Pt/AI2 03 TEMPERATURE 475 ~ C

----l---

525"C

4E

500" C 525"C

9

550" C

8E-9

2E 4E-9

0 0.00

I 0.04

0.02 pCH 4 ,

,, 0.06

bar

Figure 6. Variation of the C H 4 oxidation rate over the Pt/A1203 catalyst with PCH4 (Po2 = 0.02 bar).

0 0.0 0

"

I 0.0 2

I 0.0 4 pO 2 ,

0.0 6

bar

Figure 7. Variation of the C H 4 oxidation rate over the Pt/A1203 catalyst with Po2 (PCH4 = = 0.01 bar).

347 moleCH

rate,

1E-8

4 / m 2 sec '

I

moleCH

rate, '

'

--

6E-9

4 / m 2 sec I

I 550" C

_

8E-9

Pt-NiO/AI203 4E-9

6E-9

_

Pt/AI2 03

n

NiO/AI 2 03

_ 550" c

4E-9_

2E-9

0E+0

A~--

Pt-NiO/AI203

-~-

Pt/AI2 03

-1---

NiO/AIz 03

2E-9

_

_

0.00

9

9

m

I

T

I---IJ

I-

0.02

9

I-~

n

I

0.04

0E+0 0.06

p C H 4 , bar

Figure 8. Variation of the C H 4 oxidation rate over the laboratory catalysts with PCH4(Po2 = 0.02 bar). 4. CONCLUSION

0.00

I 0.08

0.04

B

0.12

p O 2 , bar

Figure 9. Variation of the C H 4 oxidation rate over the laboratory catalysts with Po2 (PCH4 = 0.01 bar).

NiO/A1203, Pt/AI203 and NiO-Pt/A1203 catalysts were examined in laboratory scale for total methane oxidation concerning their activity and kinetics. At the applied conditions (low partial pressures of methane and oxygen near the stoichiometric point and temperatures in the range of 475-575~ the oxidation reaction seems to follow an E-R mechanism for the three catalyst (methane from the bulk of the gas phase reacts with adsorbed oxygen). The oxidation activity is attributed mainly to Pt. The crystalline structure of nickel seems to influence the oxidation rate. No evidence of synergistic effect of nickel and platinum in the mixed catalyst was observed.

Ackwoledgments

The authors want to thanks the administration of the HELPE (ex-EKO) refinery (Thessaloniki, Greece) for the provision of catalyst samples.

5. REFERENCES 1. T. N. Angelidis, V. G. Papadakis, E. Pavlidou, Applied Catalysis B: Environmental No.4 (1994) 301. 2. F.A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 5th ed.J. Wiley&Sons, 1988. 3. R. R. Sakaida, R. G. Rinker, Y. L. Wang, and W. H. Corcoran, A.I.Ch.E. Journal No. 4 (1961)658. 4. D. Dissanayake, M.P. Rosynek, K.C.C. Kharas and J.H. Lunsford, Journal of Catalysis No. 132 (1991) 117. 5. F. van Looij and J.W. Geus, Journal of Catalysis No. 168 (1997) 154. 6. S.B. Wang and G,Q.M. Lu, Applied Catalysis B: Environmental No. 16(3) (1998) 269.

348 7. Z. Zhaolong and X. Verykios, J. Chem. Soc. Chem. Commun. (1995) 71. 8. J. R. Rostrup-Nielsen and J-H. Bak Hansen, Joumal of Catalysis No. 144 (1993) 38. 9. E. Ruckenstein and Y.H. Hu, Joumal of Catalysis No. 162 (1996) 230. 10.F.R. Hartley (ed.), Chemistry of the Platinum Group Metals-Recent developments, Studies in Inorganic Chemistry No. 11, Elsevier, 1991. 11 .S.H. Oh, P.J. Mitchell and R.M. Siewert, Journal of Catalysis No. 132 (1991) 287. 12.H. Borchert and M. Baems, Journal of Catalysis No. 168 (1997) 315. 13.J.J. Spivey, Ind.Eng.Chem.Res. No. 26 (1987) 2165. 14.W.L. Liu and M. Flytzani-Stephanopoulos, Journal of Catalysis No. 153 (1995) 317. 15.J.J. Carberry, Chemical Engineering Progress No.2 (1988) 51. 16.A. Roin, HSC CHEMISTRY ver. 2.03 Outokumpu Research Oy, Pori, Finland 17.R.F. Hicks, H. Qi, M.L. Young and R.G. Lee, Journal of Catalysis No. 122 (1990) 280. 18.C.F. Cullis and B.M. Willatt, Journal of Catalysis, No. 83 (1983)267. 19.Y.F.Y. Yao, Ind. Eng. Prod. Res. Dev. No. 19 (1980) 293.

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Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

351

Kinetic and Mass-transfer Effects in the H y d r o g e n a t i o n o f X y l o s e to Xylitol

Jyri-Pekka Mikkola* "~,Tapio

Salmi a

and Rainer Sj6holmb

a) Laboratory of Industrial Chemistry, b)Laboratory of Organic Chemistry, Abo Akademi FIN-20500 ,~bo, Finland AbstractThe paper discusses the modelling of xylose hydrogenation kinetics over Raney nickel in aqueous solutions, the determination of the hydrogen solubility in the reaction mixture as well as evaluation of mass-transfer effects in the reaction system. The hydrogenation experiments were carried out batchwise in an automatic laboratory scale reactor operating at 40 - 70 bar and at between 80 and 140 ~ The catalyst-to-xylose ratio was approximately 5 wt-% of the xylose weight. The reactor contents were analysed off-line with a high performance liquid chromatograph. The solubility of hydrogen in the reaction medium was determined with a gaschromatographic system. The solubility was found to remain fairly constant during the hydrogenation. Only a slight increase in the hydrogen solubility was detected as xylose was hydrogenated to xylitol. The main hydrogenation product was xylitol, but small amounts of xylulose and arabinitol were detected as by-products. A semi-competitive kinetic model, based on hydrogen and xylose adsorption was developed. The parameters of the kinetic model were determined with non-linear regression analysis. It turned out that the kinetic model is able to describe the formation of both xylitol and the by-products. The mass-transfer effects in the batch hydrogenation were evaluated by using measured viscosities and estimated diffusion and mass-transfer coefficients. A process simulator, utilizing the kinetic and mass transfer effects, was developed to predict the behaviour of industrial reactors.

1. INTRODUCTION Xylose is a hydrolysis product of xylane, which is the main part of hemicellulose present in all flora, particularly in birch and reed. Xylitol has become a popular sweetening agent, particularly because of its anti-caries properties. Xylitol is obtained by catalytic hydrogenation of xylose. Hydrogenation of xylose is industrially carried out in three-phase (semi)batch reactors over suspended Raney nickel catalyst. The traditional approach is based on the analysis of the main components only, although it is known that xylulose, arabinitol, furfural and xylonic acid can appear as by-products. In order to optimize the reaction conditions of an industrial hydrogenation process it is, however, necessary to have a kinetic model for the formation of both the main and by-products. The present article is a continuation of our previous work (Mikkola et al. 1998), devoted to xylose hydrogenation. In the current work, a novel semi-competitive adsorption model for organic molecules and hydrogen is proposed, and experimental data for hydrogenation, hydrogen solubility and sugar equilibria are presented. Parameters of the mathematical models are determined by regression analysis. Finally, the mass transfer effects in large-scale reactors are demonstrated with a process simulator.

352 2. E X P E R I M E N T A L An automatic laboratory scale three-phase reactor system (600 ml, the liquid volume 300 ml) was utilized in the kinetic experiments. The stirrer (Rushton turbine) speed was adjusted to 1750 rpm to ensure the absence of mass-transfer limitations. The catalyst and the reactor were preheated under hydrogen atmosphere to the desired temperature. The xylose-solvent mixture was simultaneously heated under hydrogen flow in a preheater. The xylose-solvent mixture was fed into the reactor, after which the pressure was adjusted, and the stirrer was switched on. The reactor was operated at a pressure range of 4070 bar and at 80-140~ The concentration of xylose in deionized water was between 40 and 60 wt-%. A commercial Raney Ni catalyst was used as received (approx. 90% Ni). The amount of catalyst was at 5 wt-% of the xylose weight. The medium particle size was 22 ktm. The reactor contents were analyzed off-line with a high performance liquid chromatograph (HPLC) (RI- detector, Biorad Aminex HPX-87C carbohydrate column), column temperature was adjusted to 60~ and the flow rate to 0.4 ml/min. 1H NMR spectra were recorded by a JEOL JNM-A500 FT NMR spectrometer, operating at 500 MHz. The sample was prepared by dissolving 500 mg of xylose in 1 ml of D20. The sample was held for 1 hour at each temperature before the measurement to allow for equilibration. The signals of H-2 in the ct-anomer and H-3 in the 13- anomer were used for the integration of the signal intensities. The kinematic viscosities of the xylose-xylitol solutions were determined at 80-100~ by using an Ostwald viscosimeter. The solubility of hydrogen was studied with a pressurized circulation system comprising pipelines connected to reactor, a circulation pump, sampling loops and a GC equipped with a packed column (Chromosorb 103) and a TC detector. Nitrogen was used as a carrier gas. The column temperature used was 110~ Mixtures of water, xylose and xylitol were prepared to simulate the composition of the reaction medium at various stages of the hydrogenation. Hydrogen pressure was applied, temperature adjusted to the desired level and a HPLC pump was used to circulate the reactor contents between the sampling valve of the GC and the reactor.

3. H Y D R O G E N SOLUBILITY M E A S U R E M E N T S As expected, the hydrogen solubility was strongly dependent on temperature and pressure: the higher the temperature and/or pressure, the better the solubility. The solubility was found to be at a considerably lower level for the reaction mixture than for pure water. The solubility of hydrogen increased slightly during the course of the reaction, as xylose was hydrogenated to xylitol (Fig. 1.). A correlation was developed for the hydrogen solubility:

XH =

Wxytose W xylose+ W xylitol W~tose+ W~yl,,ot

9[In

[In

9

(0.9991)- 0.1144

(0.9993)

(r/ r)

+O.O004228.1n (pH/ bar)l + (Eq.1)

~0.11603 + O.O0041126.1n (pM/ bar)l (T I K)

where wi represents the relative amount of xylose and xylitol in the reactive mixture. The fitted hydrogen concentrations are presented as a function of the pressure in the system in Fig.2.

353

Fig.I. Change of the hydrogensolubility during the reaction.

Fig.2. Hydrogensolubilityas a function of the hydrogenpressure.

4. SUGAR MUTAROTAION AND PRODUCT IDENTIFICATION The NMR measurements of the sugar mutarotation equilibria, showed the temperature dependence of the anomer equilibria in water and aqueous ethanol, ct-xylopyranose is the dominant species in the crystalline state, whereas the equilibria are shifted towards the J-form in the dissolved state. NMR analysis did not show any trace of the aldehyde form of xylose. In the literature the relative amount of the acyclic carbonyl form, in the equilibria, is extremely low (0.02%, Monosaccharides 1995). However, it is usually assumed that the hydrogenation of sugars takes place from the acyclic form. The complete reaction scheme, which was based on peak identification from the HPLC analysis, is depicted in Figs.3 and 4. The anomer equilibria is discussed in more detail in our previous publication (Mikkola et al. 1998).

Fig.3. The sugar equilibria.

Fig.4. The reaction scheme.

5. REACTION MECHANISM AND RATE EQUATIONS The kinetic modelling was based on the stoichiometric scheme and reaction mechanism shown in Figs.4 and 5. Competitive and semi-competitive pseudo-homogeneous kinetic models were applied in the fitting of rate and adsorption parameters. The simplest approach

354 was to assume completely competing adsorption of the organic molecules and hydrogen, but since the sizes of xylose, respective hydrogen molecules are very different, a semi-competitive adsorption model was developed (Mikkola et a/.1998). The semi-competitive model was K based on the hypothesis that ~.) S + , - v-- ~ S* organic molecules can occupy several primary sites, i.e. sites for 2, H 2+ 2* - ~ 2(H*) the adsorption of a hydrogen atom. Also, a plausible 3.) K isom S ,, X Y assumption that the organic K species are not able to completely 4.) xy X Y+* -~~- - X *Y cover the catalyst surface, was K made. Thus, there always remain ~ S* + 2(H*) ~R~S'X' X* + 2* interstitial sites for hydrogen S = xylose K RDS,X2 X y = xylulose 6.) adsorption. Hydrogen was X = xylitol Xy* + 2 ( H * ) ,K,, X * + 2 * A = D-arabinitol assumed to undergo dissociative H = hydrogen x * = active site adsorption and the surface 7) X* ~ - - X + * reactions were presumed to be K RDS,A the rate limiting steps. As can be ~ Xy* + 2 ( H * ) ~ A * + 2* understood, the competitive Kb model is a special case of the 9.) A* ~--- A + * semi-competitive model. K RDS,F The mechanistic steps and rate,o.~ S* ---'K F* + 3 He O F equations are listed here. ,,.~ F* ~ F + * Fig 5. The surface reaction steps in xylose hydrogenation.

ke KsK c.

Competitive model. (Eq.2)

E =(Ks + K..-k .... )cs+ K.~c~+ K.~c.,.+ K~c~

kRDS,x1KsK2I~J'I~Z [1+(1-oc) RI

m

I

\llj

f csc /

j

= .fraction of sites

cov

ered with organic

= 2 (in case of dissociative hydrogen ads.)

Semi-competitive model. (Eq.3)

13

6. R E A C T O R M O D E L

The reactor model is based on the following principles: hydrogen is continuously bubbled through the liquid phase; the gas-liquid-solid mass transfer resistance is included in the model. Because of very small catalyst particle sizes (medium particle size 22 ~tm), the intraparticular diffusion resistance was ignored. Consequently, the mass balances for the organic compounds become,

dn, dt

-

m

o,r,

(Eq.4) where ni is the amount of species Ai.

After assuming a constant liquid volume and introducing the catalyst bulk density, (pB=mcatVL),the mass balance equation becomes,

dCl dt -p r,

355 (Eq.5) For dissolved hydrogen, the mass balance can be written as follows,

N~. A~ =N~s. A.s +d n./dt

(Eq.6)

where NCLH and NLSH denote the fluxes at the g-1 and 1-s interfaces, respectively. Pseudosteady state hypothesis for dissolved hydrogen implies that dnH/dt -- O. For the catalyst particle and the liquid film surrounding the particle, the hydrogen mass balance becomes

NL~.AL~+r.m~o.=O

(Eq.7)

where the consumption rate of hydrogen is calculated from the reaction rates,

-

rH =

II~,A-,IV, I/xT' lV " ~k)CH 2,j

(Eq.8)

where ZNk denotes the series of nominators given in the rate equations, Eq.2 or Eq.3, and Di the denominator of Eq.2 or Eq.3. A combination of Eqs.6 and 7 gives, consequently, the relationship (Eq.9)

N~,. A~ +r. m~o,= 0

The flux at the gas-liquid interface is described with the law of Fick, i.e.

N~.=k~.(c*.-c.)

(Eq.10)

where Cn is the saturation concentration of hydrogen, (CH = XH Ctot). After inserting the flux expression, recalling that mcat = pBVL and a = AG#'Vc, we obtain (Eq. 1 l)

k~.a ( c . - c . ) + r . p . - O

As can be easily seen, the hydrogen concentration in the bulk phase coincides with the saturation concentration, provided that kLH a is high.

7. DIFFUSION COEFFICIENT OF H Y D R O G E N The diffusion coefficient of hydrogen in the reaction mixture was estimated by using the Wilke-Chang equation and the experimentally determined viscosity data. The Wilke-Chang equation is written as follows (Reid, et al., 1988; Wilke and Chang, 1955), where the values of the parameters were chosen accordingly: the association factor (~) = 2.6 (-- water), the average molecular mass (M) ,,~Xiorg Miorg + Xw Mw , VH2 14.3 cm3/g (Sherwood et al., 1975). =

-"

356

7.4.10_12.

D.. /

"/m s -l =

(I) M

g / mol

(T/K) (Eq.12)

cPJ~cm ,s) The dynamic viscosity (l.t) was calculated from the kinematic viscosity and the mixture density,

11 = V Pm,x

(Eq. 13)

in which

pm =Pw+Wx'P'x

(Eq.14)

In the above equation, Pw is the density of water (Kfister-Thiel 1990), Wx is the fraction of organic substance (xylose/xylitol) in the mixture (wt-%) and Px an experimentally determined scaling parameter. The density was estimated to vary between 1.18 and 1.13 kg/l at 80-140~ The viscosity was estimated to vary between 5.5 and 3.3 cP at 80~ and from 3.1 to 2.8 at 100~ The calculations indicate that the diffusion coefficient of hydrogen varies between 8.8-10 -l~ m2/s and 1.5.10 -8 m2/s at the temperature range from 80 to 140~ The role of intraparticular mass transfer effects was investigated by using the concept of a pseudo-first order reaction with respect to hydrogen. The apparent hydrogenation rate is given by the expression, ,

R =?~ek

,

,

r

(Eq. 15)

where Vie and k' denote the first-order effectiviness factor and rate constant for spherical particles. In this case, the effectiveness factor is related to the Thiele modulus (@) according to

E1 1]

3 tanh

(Eq.16)

where ~-(k'Pr,/Dem) v2 R p . Rp and pp denote the particle radius and density. The effective diffusion coefficient (Den2) was obtained from the molecular diffusion coefficient (Din), the particle porosity (ep) and tortuosity (xp). An estimate of the Thiele modulus was obtained from the experimentally observed rate (R') and the physical parameters. The diffusion effect increases as the reaction rate increases, thus the observed initial rate at the highest pressure and temperature, 80 bar and 140~ was used in calculations. The results are summarized in Table 1. The calculations indicate that the Thiele modulus can maximally attain the value ~ 1.04, which gives the effectiveness factor tie --0.93. Thus we conclude that the effect of intraparticular diffusion is in practice negligible. Table 1. Evaluation of the r61e of intraparticular diffusion. T= 140~ P= 70 bar Rp = 23 l.tm xp = 1 pp = 1500 kg/m 3 CTOT= 35.7 mol/dm 3 Den2 = 0.75.10 -8 m2/s Xn2 = 6.19.10 .4 ~) = 1.04 tie = 0.93

R'= l~p = Dn2 =

0.01279 mol/(g.min) 0.5 1.5-10 .8 m2/s

357 8. C O N L C U S I O N S The kinetic parameters were obtained from the model with non-linear regression analysis. The obtained parameters for the competitive, respective semi-competitive models are listed in Table 2. As can be seen, the amount of parameters to be fitted in the models is fairly large. Consequently, it was rather cumbersome to obtain reliable estimates for the adsorption coefficients. The obtained fit and standard error of estimate were of the same order of magnitude, for both the competitive and semi-competitive model. An interesting observation was that if allowing floating of the parameter (x, i.e. the competitiveness factor, its value was always determined with rather low error. Although the obtained fit for the semi-competitive model was of the same order of magnitude than the traditional, competitive model, we suggest that our novel approach introduces a more realistic view of the real state of the surface stage. Table2. Estimated parameter values for xylose hydrogenation over Raney Ni. Water - competitive model: Water S o l v e n t - semi-competitive model: Explained (%): 97.19 Explained (%): 96.90 Std. Error of estimate: 0.2500E+01 Std. Error of estimate: 0.2626E+01 Estimated Est. Relative Parameter/ Estimated Parameter symbol and units: Parameters Std Error (%) Std. Error Parameters .509E+00 (x 0.zi61E+05 2].1 4.7 0.100E+01 kxy mol/s'gcat 0.951E+00 7.4 13.5 0.200E+01 kRos,Xl mol/s'gcat 0.782E+04 160.7 0.6 0.301E+01 kRDS,X2 mol/s, gcat 0.621E+02 171.4 0.6 0.894E+02 kRDS,A moVs'gcat 0.100E-04 85.9 1.2 0.648E-02 kRDS,F mol/s 0.135E+05 50.1 2.0 0.134E+01 EKxy J/mol 0.530E+05 5.3 18.8 0.501E+05 EROS,X1 J/mol 0.326E+05 292.0 0.3 0.122E+06 ERDS,X2 J/mol 0.102E+06 93.5 1.1 0.715E+03 ERDS,A J/mol 0.689E+05 ]8.6 5.4 0.397E+05 ERDS,F J/mol 0.281E-1 0.123E-1 0.151E-2 0.424E-1 0.208E-1

Ks Kx KL KF KH

l/mol 1/mol l/mol 1/mol l/mol

Sample model fits are represented in Figs. 6 and 7.

~~~'~ .

data set 2

data set 2 50

. . . . .

0 '0 '0

~-

3~ \

4~ 40 35

30 25

~ /

o

25

20

/\

o

20

/o /

~

/

~

15 10

s~4 0 '~ 0-::

0

~ 9 .:

20

, ~..

40

9 ,~.

60

~ 9

80

. . . . . . . 100 120 140 160

Fig 6, The fit of the competitive model: 100~ 60 bar, 50 wt-% xylose in water.

y ~'~

~o ). -

20

,~

40

.J O

60

)

80

100

120

140

160

Fig 7. The fit of the semi-competitive model: 100~ 60 bar, 50 wt-% xylose in water.

358 To account for the coupled kinetic and mass transfer effects in large scale, a simulator was compiled. Fig. 9 demonstrates the effect of the mass-transfer limitations in the hydrogen supply. In the smaller (inserted) picture, the mass-transfer effects (kLHa) are minimized by setting a high value to kLH a, whereas in the larger one, severe hydrogen mass-transfer limitations prevail. The simulator prognoses for the concentration profiles of the various species is satisfactory.

Fig 8. XyroSIMV1.0, a study of the mass-transferlimitations.

REFERENCES:

Collins P., Ferrier R. (1995), Monosaccharides, Wiley, Chichester K/ister-Thiel, Tabelle per le Analisi Chimiche e Chimico-Fisiche, U. Hoepli Editore spa (1990), Milano Italy, 1990 (Titoli originali: Logaritmische Rechentafeln f/Jr Chemiker) Mikkola J.-P., Salmi T., Sjtiholm R. and M~iki-Arvela P. (1998), Catalysis in Multi-Phase Reactors, March 1998, Toulouse, Catalysis Today (in press). Mikkola J-P, Salmi T. and Sj6holm R. (1998), CHISA'98, August 1998, Prague Reid, R.C., Prausnitz, J.M., Poling, B.E., (1988), The Properties of Gases and Liquids, 4 th Edition, McGraw Hill, New York Seidell A. (1941), Solubilities of Organic Compounds, D.Van Nostrand, NY, 9 (2), 210-212 Sherwood, T.K., Pigford, R.L., Wilke, C.R., (1975), McGraw-Hill Inc., San Fransisco Wilke, Chang, Am. Inst. Chem. Engrs. J., (1955), 1,264, Wisniak, J., Hershkowitz M., Leibowitz R. and Stein S. (1974), Ind.Eng. Chem. Prod. Res. Develop., 13 (1), 75-79 Wisniak, J., Hershkowitz M., Leibowitz R. and Stein S. (1974), Ind.Eng. Chem. Prod. Res. Develop., 13 (4), 232-23

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

359

Methylation of Biphenyl over Zeolite H-ZSM-5 in Gas Phase with Methanol in Presence of Water: Effect of the Catalyst Impregnation by Tetraethyl Orthosilicate S. Dubuis, R. Doepper* and A. Renken Institute of Chemical Engineering, Swiss Federal Institute of Technology, 1015 Lausanne, Switzerland

Abstract The gas phase methylation of biphenyl with methanol over H-ZSM-5 and H-ZSM-5 modified by Si(OCzHs) 4 was investigated in presence of water and under atmospheric pressure. The modification of the outer surface of the catalyst pellets by sililation permits to increase the selectivity for the position 4 and 4' of the biphenyl. The ratio of the selectivity of the 4,4'-dimethylbiphenyl over the sum of the selectivities of the 12 dimethylbiphenyl isomers is increased from 0.55 up to 0.86. The 4-methylbiphenyl represents 45% of the methylbiphenyl isomers over H-ZSM-5 and 75% over modified H-ZSM-5. 1. INTRODUCTION Polynuclear aromatic hydrocarbons are envisaged as raw materials for advanced products such as liquid crystal polymers and heat-resistant polymers. The biphenyl-4,4'-dicarboxylic acid is one of the monomers used for the synthesis of such linear long chain polymers through polycondensation. The biphenyl-4,4'-dicarboxylic acid is produced by the oxidation of a 4,4'-dialkylbiphenyl. The 4,4'-dimethylbiphenyl is more convenient to use than the 4,4'-diethyl- or the 4,4'-dipropylbiphenyl because it requires softer oxidation conditions. The synthesis of the 4,4'-dialkylbiphenyl is usually carried out with conventional protic Friedel-Crafts catalysts and involves many problems from the chemical as well as from the economical-technical point of view. The propylation of biphenyl over shape selective catalysts was studied for some years and good results were obtained over Mordenite and Y-type zeolite [1, 2]. Selective C2- or C 1-alkylations, which are interesting for the industry, are not possible with simple alkylating agents over these catalysts because no steric hindrance is induced for ethyl- and methylsubstituants. This disadvantage was, till today, resolved by the use of complex alkylating agents like tetraethylbenzene and pentamethylbenzene respectively [3-5]. In such cases, the selectivity for the 4,4'-isomer is due to the conditions of transalkylation, which takes only place at the intersections of the catalyst pores. The methylation of biphenyl over H-ZSM-5 catalyst in gas phase with methanol could be an alternative route to produce 4,4'-dimethylbiphenyl because this zeolite induces steric hindrance on methyl-substituants. The modification of the outer surface of the catalyst pellets by impregnation with a solution of tetraethyl orthosilicate to enhance the catalyst selectivity [6] was performed. The results of the methylation of biphenyl over H-ZSM-5 and modified H-ZSM-5 are presented herein. The thermodynamic equilibrium for the three methylbiphenyl isomers favours the Corresponding author

360 3-methylbiphenyl, then the 4-methylbiphenyl and finally the 2-methylbiphenyl. The catalyst process has to avoid the formation of the 3-methylbiphenyl to enforce the 4,4'-dimethylbiphenyl selectivity among the twelve possible dimethylbiphenyl isomers.

+

x MeOH

=

~

m

= 12 isomers

y~

+ x MeOH /

0

N

Figure 1 Methylation of biphenyl with methanol and parasite reactions over acid catalysts.

2. EXPERIMENTAL

2.1 Catalysts and Reagents The H-ZSM-5 (Zeocat PZ-2/50) as 1/16 in. extrudates was supplied by CU Chemie Uetikon AG (Switzerland). The methanol (99.9%) was supplied by Romil (U.K.) and the biphenyl by Aldrich Chemie AG (Switzerland) with a purity of 99%. The tetraethyl orthosilicate used to modify the H-ZSM-5 was supplied by Aldrich with a purity of 99.9%. Two batches (M1 and M2) of 60 g of H-ZSM-5 were soaked for a night at room temperature in a solution of 13.3 g of Si(OC2H5)4 dissolved in 150 ml of cyclohexane supplied by Fluka Chemie AG (Switzerland). These two batches were prepared to test the reproducibility of the impregnation procedure. One third batch (MM) of 60 g of H-ZSM-5, was soaked under the same conditions with a solution of 26.7 g of Si(OC2H5)4 dissolved in 150 ml of cyclohexane to study the influence of the silicate concentration on the impregnation. The cyclohexane was evaporated from the three batches. The catalyst pellets were then dried at 120~ under vacuum for 3 hours and calcinated during 5 hours at 540~

361 2.2 Experimental set-up and procedure A tubular stainless steel reactor with an internal diameter of 32 mm heated by an electrical oven is used for the methylation of biphenyl under atmospheric pressure. The catalyst bed length is equal to 11.5 cm for the charges of 53 g of catalyst used in all experiments. The fixed bed can be considered as an ideal plug flow reactor. Biphenyl, water and methanol are dosed by adjusting nitrogen flows through three different bubble-columns regulated at specified temperatures and pressures to obtain the desired partial pressure of the different feed components. Nitrogen is also used as diluent and oxygen is used for catalyst regeneration. Mass flow controllers (Bronkhorst High-Tech B.V.) regulate the gas flows. The temperature of the catalyst-bed is measured by two K-type thermocouples (Philips AG). The thermal gradient within the fixed bed does not exceed 1.5 K. The composition of the gas flow at the outlet of the reactor is analysed by gas chromatography (Shimadzu 14A, capillary column HP5, 30m x 0.32 mm I.D x 0.25 ~tm film thickness, FID) and then cooled to condense the aromatic products. The condensed fraction is analysed with a GC/MS system (HewlettPackard, G1800A, GCD System, capillary column HP-5, 30m x 0.25 mm I.D x 0.25 gm film thickness, EID) to permit the identification of the alkylated aromatic products. The methylation reactions were carried out during three to four hours and then the catalyst bed was regenerated. The regeneration is carried out after 10 minutes of nitrogen purge with a 10 percent oxygen flow in nitrogen and an increase of the temperature from the reaction temperature up to 800 K. The regeneration is stopped when the temperature of the thermocouple at the end of the catalyst bed is stable after having passed through a maximum. The catalyst activity is totally recovered by the regeneration procedure.

3. RESULTS AND DISCUSSION 3.1 Methylation over H-ZSM-5 The methylation of biphenyl (BP) over H-ZSM-5 zeolite, under the conditions used, produces mainly 3- and 4-methylbiphenyl (4-MBP), 4,4'-dimethylbiphenyl (4,4'-DMBP) and cracking products (benzene, toluene, xylene, trimethylbenzene .... ) (Figure 1). The production of 2-methylbiphenyl isomer is negligible. The methanol reacts also independently to form dimethylether and gasoline (MTG reaction). This consumption of methanol by undesired reactions limits the biphenyl conversion to about 0.25. The reaction system is subject to a deactivation due to the catalyst coking but it can be regenerated by coke burning. The higher the methanol inlet molar fraction is, the faster is the catalyst deactivation. The biphenyl conversion falls from 0.19 to 0.13 within 3 hours (Figure 4). The reaction was led in presence of water to reduce the biphenyl cracking and the production of dimethylether and gasoline. The transient behaviour of the catalyst is also strongly influenced by the presence of water. Whereas, in the absence of water the 4,4'-dimethylbiphenyl selectivity continuously increases over a time period of 3 hours, the start-up time of the reactor is reduced to about 1 hour in the presence of water in concentrations higher than 10%. But due to coking, the deactivation of the catalyst can not be avoided (Figure 4). Coking and cracking become predominant at temperatures higher than 560 K. All the selectivities mentioned hereafter correspond to the selectivities observed after start-up time and they are considered as initial selectivities.

362 0.6

Z 0

.~_

0.5

V'~_~

O.4

[3

---"~~

V

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

~ 0.3 "o E lg

i~_ 0.2

A

S4,4, DMBP

O

S4_MB P

[~

9

~

Xv

S3_MB P

S4,4,_DMBJ~DMBP

0.1 __ "1''

'

I

0.01

0.02

._

'

/

'

O.03

___ 9 I

0.04

Methanol inlet molar fraction, Y~OH,O[-]

Figure 2 Selectivities and ratio of selectivities of the products as function of the methanol inlet molar fraction. Y = 5 4 0 K, nlca t = 53 g H-ZSM-5, YsP,0= 0.01, YH20,0 = 0.2, x' = 6660 kgcatS/molBp 0.60-

0.55 " " O.50 O "~ 0.45 ._~ 0.40 " "~ 0.35 ".

~

~ o.3o -~ E 0.25

4,4'-DMBP

.

4-MBP 3-MBP

> 0.20 " .m

S4,4,_DMBP / ~ SDMBP

"~ 0.15

S4.MB P / ,~, S M B P

o,qJ 0.10 -~ 0.05 40OO

A

l~lb'

!

6000

Ak ,, . . / . . . . .' . .

Modified

8000

I

10000

i

!

12000

,

I

14000

Residence Time, 1,' [kg,=t*s/mOIBp ]

Figure 3 Selectivities and ratio of selectivities of the products as function of the modified residence time. T - 540 K, mcat = 53 g H-ZSM-5, YBP,0--- 0.01, YMeOH,0= 0.01, yn2o,o = 0.2 Because the 4,4'-dimethylbiphenyl is produced by a consecutive methylation reaction, a high selectivity for the 4-methylbiphenyl is necessary to obtain the 4,4'-dimethylbiphenyl with a good selectivity in the consecutive second methylation. The 3-methylbiphenyl is the main product over H-ZSM-5 with a selectivity between 0.4 and 0.5 (Figure 2). The

4-methylbiphenyl is the second most important product. The 4,4'-dimethylbiphenyl is, with a selectivity in the order of 0.07, the third product. An increase of the methanol inlet molar fraction induces a decrease of the reaction selectivity for the position 4 and 4' of the biphenyl. This can particularly be seen by the decrease of the selectivity ratios of the desired isomer among all the possible isomers. Similarly, an increase of the modified residence time (Figure 3) results in the same effect as an increase of the methanol inlet molar fraction. The influence of these two parameters can be explained by the reactions of methylation and isomerisation that occur at the outer surface of the catalyst pellets. If the methanol concentration is higher, the methylation rate is increased mainly at the outer surface of the catalyst pellets where no diffusion limitation and steric hindrance occur. With an increase of the residence time. the products are staying for a longer time in contact with the catalyst and the isomerisation reactions are favoured. The 3-methylbiphenyl, which is more stable thermodynamically, becomes also more important. A value equal to 0.55 for the ratio of the 4,4'-dimethylbiphenyl selectivity over the sum of the selectivities of the twelve dimethylbiphenyl isomers is reached at short residence time and low methanol inlet molar fraction for an inlet molar fraction of biphenyl of 0.01 that corresponds to the most favourable conditions. 3.2 Methylation over modified H-ZSMd I

OM,

Time on stream [ h : ~ n ]

Figure 4 Conversion of biphenyl as function of time on stream over four different catalyst charges. T=540 K, m,,=53 g H-ZSM-5, y,,,=O.Ol, yM,o,,~o=O.Ol, y, 2 00=0.2, ~ ' = 6 6 6 0kg,,, slmolBp . The impregnation of the outer surface of the catalyst pellets with Si(OC2H5)4,which is too big to penetrate into the pores of the zeolite [6], induces a reduction of the activity of the catalyst (Figure 4). For identic reaction conditions, the biphenyl conversion decreases from 0.19 to 0.14. The modification procedure is reproducible as the results over catalyst charge M1 and M2 show. An increase of the tetraethyl orthosilicate concentration in the impregnation solution does not permit to ameliorate the catalyst selectivity. The catalyst

364 activity is similar for all the three modified catalyst charges. The passivation of the outer surface of the catalyst by sililation is efficient to favour the methylation of the position 4 and 4' of the biphenyl (Figure 5). The selectivity of 4-methylbiphenyl is higher than the 3-methylbiphenyl selectivity. The sum of the selectivities of 4-methylbiphenyl and 4,4'-dimethylbiphenyl can reach up to 0.77. The ratio of the 4-methylbiphenyl selectivity over the sum of the selectivities of the three methylbiphenyl isomers is increased from 0.4 to 0.7 by the catalyst modification at a modified residence time equal to 6660 kgcatS/mOlBp. The ratio of the 4,4'-dimethylbiphenyl selectivity over the sum of the selectivities of all the twelve dimethylbiphenyl isomers is enhanced from 0.49 to 0.76. This ratio reaches 0.86 at the shortest residence time tested (Figure 7).

Figure 5 Selectivities and ratio of selectivities of the products over H-ZSM-5 and modified H-ZSM-5 at two biphenyl inlet molar fractions. T - 540 K, meat=53 g, YMeon,0=0.02, Yu20,0=0.2, Z'=6660 kgcatS/mO1Bp An increase of the inlet molar fraction of methanol acts similarly over unmodified and modified H-ZSM-5, but the effect is reduced for the latter (Figure 6). This shows that the passivation by the tetraethyl orthosilicate is not sufficient to avoid completely the reactions at the outer surface of the catalyst. The enhancement of the selectivity for the position 4 and 4' depicts in Figure 7 at lower residence time confirms this fact. The selectivities of the methylbiphenyl and dimethylbiphenyl isomers are remarkably constant in the temperature range studied (Figure 8).

365

0.8

.'T:. 0.7 O 0.6 > 0.5 (D O.4 "o E: 0.3

._Z',

V O

~

9

S4,4,.DMB P

0

S4_~p

O

S3_MB P

v

S.~ZSMBp

._> .,~ 0.2 ~

0.1

0.0

I

'

0.010

I

0.015

'

I

0.020

'

I

0.025

'

I

0.030

'

I

0.035

'

I

0.040

Methanol inlet molar fraction, Y~OH,O[-] Figure 6 Selectivities and ratio of selectivities of the products as function of the methanol inlet molar fraction. T = 5 4 0 K, mca t = 53 g modified H-ZSM-5 (M1 and M2), YBP,0 = 0.01, YH20,0 = 0.2, "r' = 6660 kgca t s/mo1Bp 0.9

//

'--' 0.8 I i1._1 o "~ 0.7

v

:>_- 0.6

0

"d

S4,4,_DMBP

9

_.e

0.5

"~ r-

0.4-

~ 0.3 .>__ ~ 0.2

O

$4_M8P

r"l

S3_Mgp

9

V

S4,4,.DMBW~'SDMBP

S4_MBW~SMBP [] []

0.1

0.0

4000

.

,

5000

.

,

6000

.

70'00

.

//

,

13000

.

14()00

Modified Residence Time, x' [kgcat*S/mOIBp ]

Figure 7 Selectivities and ratio of selectivities of the products as function of the modified residence time. T=540 K, mcat=53 g modified H-ZSM-5 (M1 and M2), yBP,0=0.01, YMeOn,0=0.01, YH20,0=0.2

366 1.0

z

o

O.90.8-

>, 0.7 >

12; 0.6 (I.) (1) 03 0.5 "o C 0.4

V

-------

0

0

>,

"~

~d

(1) 00

9

S4,4,_DMB P

rl

S3_MBP

9

S4,4'_DMBg~-SDMBP

v

0.3

o O

S~MBP

S.JZS~

0.2 0.1

0.0

i

520

&

_A

&'

I

525

'

I

530

'

~

5

'

I

540

9 '

I

545

'

i

550

Temperature [K]

Figure 8 Selectivities and ratio of selectivities of the products as function of the temperature. ITlcat = 53 g modified H-ZSM-5 (M1 and M2), YBP,0= 0.01, YMeOH,0= 0.01, YH20,0 = 0.2, Z' = 6660 kgcatS/mO1Bp 4. CONCLUSION The modification of the outer surface of the H-ZSM-5 zeolite by impregnation with a solution of tetraethyl orthosilcate permits to increase the selectivity of the 4,4'-dimethylbiphenyl produced by the methylation of biphenyl in an heterogeneous catalytic gas phase process. The total selectivity for the position 4 and 4' reaches 0.77. The 4,4'-dimethylbiphenyl corresponds to 86% of the twelve dimetyhlbiphenyl isomers over the modified zeolite. In an industrial point of view, the increase of the conversion of the biphenyl in avoiding a decrease of the 4,4'-dimethylbiphenyl is possible by a multistage methanol injection in the catalyst bed. AKNOLEDGMENT The financial support provided for this work by the Swiss National Science Foundation is gratefully acknowledged. REFERENCES [1]

[2] [3] [4]

[5] [6]

Lee, G. S., Maj, J. J., Rocke, S. C. and Garces, J. M., Catal. Lett., 2 (1989), 243-248. Lee, G. S., Garces, J. M. and Maj, J. J., Catal. Sci. Technol., 1 (1.990), 385-387. Brechtelsbauer, C. and Emig, G., Appl. Catal., A, 161 (1997) 1-2, 79-92. Brechtelsbauer, C. and Emig, G., Chem. Eng. Technol., 20 (1997), 582-588. Takeuchi, B., Shimioura, Y. and Hara, T., Appl. Catal., A, 137 (1996), 87-91. Wang, X. S., Wang, G. R., Guo, H. C. and Wang, X. Q., Stud. Surf. Sci. Catal., 105 (1997) PA-C, 1357-1364.

Reaction Kinetics and the D e v e l o p m e n t of Catalytic Processes G.F. F r o m e n t and K.C. W a u g h (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

Kinetics and m e c h a n i s m studies of the catalytic d e h y d r o g e n a t i o n of isobutane on P l a t i n u m - I n d i u m catalyst D. Casanavel, K. Fiaty" 2, J.A. Dalmon 1, M. Forissier 3 ~institut de Recherches sur la Catalyse (IRC) UPR CNRS 2 av. Albert Einstein 69626 Villeurbanne Cedex, France -~Laboratoire d'Automatique et de G6nie des proc6d6s, UPRES-A CNRS Q5007 Universit6 Claude Bernard Lyon-I et CPE Lyon, B~t 308G, 43 Bd du 11 Novembre 1918 3 Laboratoire du G6nie des Proc6d6s Catalytiques (LGPC) CPE Lyon, B~t 308G, 43 Bd du 11 Novembre 1918 Abstract The catalytic dehydrogenation of isobutane into isobutene was studied in a differential reactor using platinum-indium supported on silicalite as catalyst. The rate controlling mechanism was found to be the surface reaction of the adsorbed isobutane. The adsorption and kinetic coefficients were determined by the Newton method applied to the reaction rate equation. Statistical tests were performed on these coefficients in order to evaluate their significance on the multiple regression. I. INTRODUCTION The dehydrogenation of isobutane is an important reaction from the industrial view point. Isobutene is indeed a raw material for the production of MTBE, synthetic rubber and many others organic products. Two types of catalyst have been developed and patented for this reaction: chrome based catalysts [1-3] and more recently platinum supported catalysts [4-7]. Both of them suffer from deactivation due to coke deposition and often a feed of hydrogen is used in order to improve the catalyst stability. The first studies of the kinetics and mechanism of the reaction were performed on chromia-alumina catalysts, for the dehydrogenation of n-butane [8-11] and isobutane [12]. All the kinetic equations have been developed on the Langmuir-Hinshelwood scheme, with a dual site mechanism, in agreement with the experimental data. For the nbutane, the rate determining step was found to be the surface reaction, which involves two adjacent active centers, whereas for isobutane the reaction is both adsorption of butane control and surface reaction control. As far as the platinum based catalysts are concerned, kinetics and mechanism of the dehydrogenation of iso- and nbutane have been studied by Lyu Cam Loc and Co-workers [13-19], with platinum on y- alumina catalysts, promoted or not with indium or tin. The mechanistic study, made by means of the response method in transient isotopic exchange, have allowed to propose the following stage wise scheme: @CaHIo + 2 * <=:>C4H9" + H* O C4H9" + * r + H* O C4H8" ,::::>C4Hs + * O2H* c=>H2 + 2 * where stage O appears to be slow and * represents an active surface site. The rate determining step is the adsorption of butane. The experimental rate laws for dehydrogenation of isobutane (iC4) and n-butane (nC4) are respectively: r=k I

Pic 4 *Yi Pic4_ + k2 P05 H2

(1)

PnC4 * ~' i 3 ~ (2) ! + k4PnC4= where k~, k2, k3, k4 are constants and Yi is a correction factor introduced for taking into account the back reaction, iC4_-, nC4__represent respectively isobutene and butene. On chromia based catalyst, Masson et al [12] proposed the following scheme for the dehydrogenation of isobutane in three steps: O iC4Hio + * --)' iC4Hio* O iC4Hio* --~iC4Hs* + H2 r= k

9Corresponding author e-mail: [email protected] 1.fr

367

368

OiC4Hs*

~iC4H s

+ *

The three above steps are respectively: non-dissociative adsorption of isobutane at the surface of the catalyst, surface reaction of the adsorbed isobutane given the adsorbed isobutene, desorption of adsorbed isobutene. In this study the kinetics of isobutane dehydrogenation on Platinum-Indium (Pt-In) catalysts is investigated, in order to determine the rate law of the reaction, for further use in membrane reactor modeling. This law has to describe the rate of isobutene production in the equilibrium range, where the reverse reaction could not be neglected. For this purpose, kinetic studies have been made in a differential reactor, where perfect-mixed flow is assumed. The influence of the partial pressures of reactants and products and the influence of the temperature on the reaction rate were studied for both the dehydrogenation of isobutane and for the hydrogenation of isobutene. 2. EXPERIMENTAL

2. i. Materials in all experiments, isobutane and isobutene with high purity from Air Liquide were used. The gases were further purified using filters and adsorbents. The catalyst is a platinum-indium catalyst supported on silicalite, obtained from "lnstitut de Recherches sur la Catalyse" (IRC) [20]. It contains 0.5% of platinum. 2.2. Experimental set-up Experiments were performed using a differential quartz microreactor shown in figure 1. The reactor was supplied with different ratios of the following gas mixture: isobutane, hydrogen and nitrogen for the dehydrogenation of isobutane, and isobutene, hydrogen and nitrogen for the hydrogenation of isobutene. The mixture flow rate was controlled by mass flow controller. Furnace equipped with PID controller was used to heat the reactor. Reactants and products analysis were performed by gas chromatography with two detectors: FID for hydrocarbon analysis and TCD for hydrogen one.. 2.3. Procedure In order to have isobutene yields lower than 5%, all runs were conducted with very short contact time (1-5 ms) because of the high catalyst activity. About 6 to 10 mg of catalyst powder were introduced in the reactor between two quartz wool plugs, forming a 2 mm high catalyst bed. The catalyst was first activated with an hydrogen flow of 50 ml/min, at 723K for two hours, after a slow temperature raise (l~ Then the reactor was swept with nitrogen, while heating or cooling it at the chosen reaction temperature. After all, it was fed by the gas mixture at the desired ratios, with a flow rate of 100-200 ml/min. Before a series of runs begun, the catalyst deactivation was measured until there was hardly no deactivation, which usually occurred after about I0 hours on stream. Two kind of experiments were conducted to study in one hand the effect of reactants partial pressure at 723K, in other hand the effect of reaction temperature ranging 673K-723K. Gas chromatography analysis were performed online at different time intervals. Only are detected isobutane, isobutene and hydrogen. Therefore we assumed that side reactions did not occurred. From the gas analysis and the flow rate value Qs, the production rates of isobutene (dehydrogenation) and isobutane (hydrogenation) were then calculated according to the following relationships: Qs 1 ric4= = Pic4= RT m

(3)

Qs 1 ric4 = Pic4 RT m

(4)

3. PRELIMINARY STUDIES: DIFFUSION This study was done in order to find the operating conditions where the differential reactor operated in chemical regime. It means that the measurement have to be realized without any disturbance from diffusion limiting effects by reactants and products. Two kind of diffusion can occurred: extragranular diffusion which concern mass transfer from bulk fluid to the external surface of the pellet in the catalyst bed, and intragranular diffusion for mass transfer into and through the pores within the pellet with reaction taking place on the catalytic surface of the pores. Quick estimates of the external diffusion can be done by applying Mears' criterion which uses the measured rate of reaction to learn if mass transfer from the bulk gas phase to the catalyst surface can be neglected. This is the case when: r~ < 0.15 (5) kDCbulk where n is the reaction order, R the catalyst particle radius, p the bulk density of catalyst bed and Cb,~k the bulk concentration. The mass transfer coefficient kD can be calculated from appropriate correlation. We can also based our analysis on the estimation of the pressure gradient between the bulk and external surface of the catalyst from the

369 measurement of the reaction rate which in steady-state is the same as the reactant diffusion rate. Therefore the following expression can be used: AP = R___.~TrObS (6) kDa where a is the specific area. In our study it was found for the highest reaction rate that the isobutane pressure gradient is about 710 -5 atm which is negligible comparing to its bulk pressure 0.2 atm.

3

--i

iC'i~':H~-~ Reaction mixture

N2~ H~

t~urnace

__~

Oz/Ar Activation-Regeneration

I:

p

H2

Chromatograph

9

r

Chromatograph

(!)

Figure 1: Experimental set-up In the case of internal diffusion, Weisz-Prater Criterion is commonly used. There are not diffusion limitations and consequently no concentration gradient within the pellet when: rObSoR2 rldp2 = ~ >> 1 (7) DeCs where Cs is concentration of the reactant at the external surface of the catalyst pellet, De the effective diffusivity, 1"1the effectiveness factor and dp the Thiele modulus. In our study it was found that a particle size of 0.5 mm is necessary to virtually eliminate diffusion control (rl =0.9). The kinetic of the dehydrogenation of isobutane will be studied with catalyst having particle size lower than 0.1 mm.

4. RATE LAW Before establishing the kinetic law, experiments were performed in order to determine the effects of isobutane and hydrogen partial pressures on the reaction rate at 723K. For the dehydrogenation reaction, first isobutane partial pressure was varied from 0.05 atm to 0.3 atm when feeding the reactor with nitrogen/isobutane gas mixtures at different compositions. Then isobutane partial pressure was set at 0.2 atm while varying hydrogen partial pressure in the range 00.17 atm. It was found, when plotting the logarithm of the measured reaction rate as function of logarithm of partial

370

pressure in each case, that the kinetic order were 0.63 and -0.5 respectively of isobutane and hydrogen. When studying hydrogenation reaction, firstly we set the partial pressure of hydrogen at 0.2 atm while isobutene partial pressure varied between 0.05 atm and 0.25 atm. Secondly, we fixed the partial pressure of isobutene at 0.2 atm while hydrogen partial pressure was varying in the range 0.05-0.25 atm. The kinetic orders were found to be 0.8 and 0.5 respectively for isobutene and hydrogen. The value of isobutane lower than 1 suggests a equilibrium adsorption step of isobutane in the mechanism while the hydrogen order is in agreement with the value obtained in literature. Therefore, among the large number of possible reaction mechanisms (Casanave [21]), we retained the five following kinetics models which can described the dehydrogenation of isobutane on Pt-ln supported on silicalite: 4.1. Kinetic law 1 and kinetic law 2

These kinetics law derived from the one developed by Lyu Cam Loc for the dehydrogenation of n-butane and isobutane on platinum catalysts supported on y-alumina [13-19]. It originated from the following reaction schemes: - Dissociative adsorption of isobutane at the surface of the catalyst kl iC4Hlo + 2 * r iC4Hg* + H* k_l - Surface reaction of the adsorbed isobutane k2 iC4Hg* + * 'r iC4H8* + H* k_ 2 - Desorption of adsorbed isobutene iC4Hs*

r iC4H8 + * Kic4= - Desorption of the adsorbed hydrogen 2H* ~ H2 + 2 * KH Depending upon whether adsorption, surface reaction or desorption are rate controlling and upon the way hydrogen is desorbed the possible rate equations derived were: a-Adsorption of isobutane controlling:" kinetic law 1 k 1Pic4

r=

-

k-IKic4 KH K2 ic4: PH,_

-

-

P

(8)

C1 + KiC4= Pc4 " + ~/KHPH2 + KiC':K2K ~ H Pc4 P~/~2) 2 with K 2 = k--L-2 k_2 where r is the dehydrogenation reaction of isobutane. b-Adsorption of isobutane and surface reaction are rate controlling steps: kinetic law 2

(9)

1 , _- PH2 ) k2 +k_I~]KHPH 2 (k,k2Pic 4 - k - l k - 2 K i c 4 = KH Pc4 r=

2 (I+Kic4-P'c4 _

(10)

+ KHPH2J +klPic4+k-2Kic4 ) ~/KHPH/P'c, = , =

, :

k 2 + k_ I~]KHPH2

4.2. Kinetic law 3 and kinetic law 4

These models are similar to the model developed by Masson [12] on chromia based catalyst. But in this case the adsorption of hydrogen at the catalyst surface was taken into account. The reaction scheme is in four steps: Non-dissociative adsorption of isobutane at the surface of the catalyst kl iC4HI0 + * r iC4Hio* k_l Surface reaction of the adsorbed isobutane k2 iC4Hto* + 2 * ~ iC4Hs* + 2H* k_2 Desorption of adsorbed isobutene -

-

-

371

iC.~Hs*

r iC4H8 + * Kic4=

- Desorption of the adsorbed hydrogen 2H*

r H2 + 2 * KH

Kinetic law 3: This law was obtained by assuming surface reaction to be the rate controlling step. Consequently, the rate equation for the reaction has the form: k2K,c 4 Pic4 - k-2Kic4:KHPic4 - PH2

r=

(l+Kic4P,c 4 +Kic4:Pc4

3

(ll)

+~KHPH2)

kl

with Kit_4-

(12)

k_ 1

Kinetic law 4

In this case the rate of isobutane adsorption at the catalyst surface is assume to be slow. Therefore the rate equation was: klPic4

-

k-i Kic4= KH K2

P

ic~= PH~_

r=

(13) I+Kic4_Pc4 ' -

+~/KHPH2 + KiC4-KH p ic4~ Ptl 2 K2

4.3. Kinetic law 5

The reaction scheme of this kinetic model is a variant of the one above mentioned. It was develop in five steps where the surface reaction was assumed to be hold in two consecutive steps. The mechanism proposed is: Non-dissociative adsorption of isobutane -

iC4Hto + -

*

r iC4Hlo* KiC4 First step of surface reaction of the adsorbed isobutane kt * 'r [iC4Hlo---i C4H8]* + H* k_ ! Second step of surface reaction of the adsorbed intermediate specie

iC4Hio* + -

k2 <::=> iC4Hg* k_ 2 Desorption of adsorbed isobutene

[iC4Hjo---i C4H8]* + -

*

+ H*

iC4Hs* -

<::> iC4H8 + * KiC4= Desorption of the adsorbed hydrogen

2H*

r I-l, + 2 " KH where is [iC4Hlo ---iC4Hs] an intermediate specie. We assumed the first surface reaction to be the rate determining step and the four others were considered to be equilibrium steps. The rate equation is then: klKic4 Pie4 r=

k-lKic4__ KH K2

PIE4=PH2

(14)

. . . .

The five relations (kinetic laws) are supposed to describe the behavior of the dehydrogenation of isobutane on PtIn/silicalite and contains a number of parameters of which the values are not known beforehand and need to be estimated. For hydrogenation of isobutene, the rate laws are the opposite of those obtained for the dehydrogenation of isobutane.

372

5. PARAMETERS ESTIMATION AND STATISTICAL TEST There exist several techniques which can be used to determine the unknown parameters such as the method of maximum likelihood and the method of moment. An overview of these techniques was reviewed by Froment [22]. The most widely used method is the least squares technique which consists of minimizing the objective function F defined as the sum of squares of residuals and given by: F = ~-~/r cal r~ 2 ~.J-.i / (15) J where n is the total number of observations. The residual is the difference between the predicted data by one of the three expressions of the reaction rate derived from the mechanism ( r cat ) and the experimental data ( r ~ NewtonRaphson method is used for minimization. Other numerical techniques are presented by Van der Linde et al [23]. Results of the minimization are shown in table 1. Table I: Results of the multiple regression Kinetic law 1 k I (j,tmol.s-~.g -I) 284 k. I (~.unol.s-l.g l ) 4.28 105 k2 (~mol.s-l.g -I) K2 = 1090 k. 2 (~trnoi.s-l.g-i) _ Kit4= (atm "l) 2.80 KH (atm l ) 14.3 F 177

on the different reaction rate equations. Kinetic law 2 Kinetic law 3 545 Kic4 = 1 1.52 104 271 306 4400 1500 1 2.91 0.105 4.67 42 906

Kinetic law 4 384 5880 K2 = 0.009

Kinetic law 5 308 1.82 107 K 2 = 5.85 104 Kic4 = 1.39 3.18 19.6

6.9 320 266

The best fitting result was obtained with kinetic law 5 where surface reaction was assumed to be the rate determining step. In figure 2, the calculated and observed reaction rates were plotted as function of partial pressure of isobutane (PH2 = 0 ) or partial pressure of hydrogen (P ic4 = 0.2 atm) for dehydrogenation of isobutane. The calculated and experimental reaction rates for hydrogenation of isobutene versus the partial pressure of isobutene (Pro = 0.2 atm ) or the partial pressure of hydrogen (P ic4--= 0.2 atm) were shown in figure 3. A good agreement between the reaction rates calculated from the retained model and the experimental data are obtained for dehydrogenation as well as hydrogenation reactions. r (~moi/s.~ 6

o

;o~ ~o~, o ~

o~:o~;~;~

0.!

P(ntm)

Figure 2: Dehydrogenation rate calculated from kinetic law 5

0.2 P(atm)

Figure 3: Hydrogenation rate calculated from kinetic law 5

Significance of the parameters were tested by means of statistical criteria. Student's t-test is used to establish whether the estimates are significantly different from zero. The confidence intervals were determined from the covariance matrix.In the case of least squares method, the covariance matrix V is approximated by the following expression: V = s2(jTj) -!

(16)

where s 2 the variance is an estimate of the population variance o 2 defined by: s2 =

F (17) n-p where n is the number of observations, p is the number of parameters. The Jacobian matrix J and the criteria F are evaluated at the estimated parameters.

373 The confidence interval was then constructed for every element 0~ of the parameter vector | level i-et and the Student's t-distribution as follows: 0i = Oi + t O l - P ) ~ i i

at a given probability

(18)

where 0i is the estimated value of the i-th element of parameter vector, t ~ -p) is or/2 percentage point of Student's tdistribution with n-p degrees of freedom, and Vii the i-th element of the primary diagonal of the covariance matrix V. The result of the test is given in table 2 for the kinetic law 5. Table 2: Parameter estimated and 95% confidence limits, number of experiments n = 26 Kinetic law5 kl (ktmoi.s".g") 308 + 73 k.i 0-tmol.s~.g "1) 1.82 107+ 3 10 tt K2 5.85 104+9 l0 s Kic4(atmz) 1.39 + 0.5 Kit4= (atm "l) 3.18 + 1.8 KH (atm ~) 19.6 + 8 From this statistical analysis it can be seen that the estimated values K2 and k_= are not contained within the 95% confidence intervals and are not significantly different from zero. More over the analysis of the correlation matrix N defined as: V0 N=[N~.i] with Nij = ~ ~

(19)

showed the interaction between these two coefficients. Therefore the can be combined in one coefficient k'.~ defined by the following relation' k-i _

k'_~ - ~

(20)

K2

Furthermore the term KiC4--~/KH Pic4= P ~ 2 represented the surface rate covered by the intermediate adsorbed specie K2 and can be neglected when comparing it with one. The kinetic law 5 can then be simplified, given the following rate equation: r=

k iKic4 Pic4 - k'-I Kic4_ KHPic4_ PH2 -

(21)

By applying multiple regression techniques and statistical tests at 95% confidence limits to the simplified reaction rate equation, the following coefficients were obtained: kl = 348 + 99 lamol.si.g l k'.l = 386 5:150 ~tmol.s'l.g"l Kic4= = 2.6 + 1.1 atm "= K , = 16.5 + 5.8 atm "~

Kic4 = 1.1:!: 0.5 arm "!

It was checked that this above mentioned rate law we derived, and which was consistent with experimentally determined rate, gives the correct value of the equilibrium constant at the operating temperature. 6. CONCLUSION The present work deals with the kinetic law study of the dehydogennation of isobutane using Pt-ln/silicalite as catalyst. Five possible mechanisms were tested and whereby the coefficients of the corresponding rate equations are calculated by the Newton method and the selection of the proper mechanism is based on the condition the experimentally-observed dehydrogenation as well as hydrogenation reactions rates must be well described by the reaction rate equation retained. It was shown that the surface reaction of the adsorbed isobutane given a reactive intermediate specie is the rate controlling step. From the statistical test, it was shown that the reaction rate expression obtained can be simplified. Therefore the dehydrogenation rate of isobutane on Pt-ln/silicalite at the vicinity of the thermodynamic equilibrium is of the form: [=

k ! Kic4 Pic 4 - k'-! Kic4= KHPic4= PH~

This simplified equation implies a mechanism in five steps: adsorption of isobutane, first surface reaction, second surface reaction, desorption of the adsorbed isobutene and desorption of the adsorbed hydrogen. This mechanism which

374

is a probable one is in agreement with the kinetic data. This agreement does not prove that the reaction mechanism is the correct one. Therefore additional chemical information such as results of tracer studies, identification of reaction intermediates by spectroscopic or other means (Katzer et al. [24]) is needed.

ACKNOWLEDGMENTS

The authors acknowledge the financial support for this study provided by Total Raffinage Distribution, and the research group of P. Meriaudeau at IRC for kindly providing the catalyst. NOMENCLATURE dp: pore diameter F: objective function J" jacobian matrix K ~c4 : adsorption equilibrium constant of isobutane (atm"t) K ic4-- : adsorption equilibrium constant of isobutene (atm "1) K u : adsorption equilibrium constant of hydrogen (atm~) m: mass of catalyst (g) N: correlation matrix PH2: partial pressure of hydrogen (atm) P ic4: partial pressure of isobutane (atm) P ic4--:partial pressure of isobutene (atm) Qs: stream volumetric rate r: rate of reaction (~tmol.st . g t ) r ic4="rate of isobutene production (lamol.st.g~) r ~c4:rate of isobutane production (~tmol.st.g~) s: standard deviation T: reaction temperature (K) V: variance-covariance matrix REFERENCES

[i] E. Pop, N. Goidean, D. Goidean and G. Serban, DE 2401955, 17 Jul 1975 [2] B. Schramm, J. Kern, H. Schwahn, A.W. Preuss, K. Gottlieb and H. Bruderreck, DE 3739002 AI, 24 May 1989 [3] J.F. Kirner, GB 2162082 AI, 29 Jan 1986 [4] E. Box, US 3692701, 19 Sep 1972 [5] F.C. Wilhelm, US 3998900, 21 Dec 1976 [6] S.A.I. Barri, and R. Tahir, EP 351066 AI, 17 Jan 1990 [7] P.R. Cottreil and M.E. Fettis, US 5087792, I I Feb 1992 [8] S. Carra, L. Forni and C. Vintani, J. Catal.,9(2)(1967)154-165 [9] S. Noda, R.R. Hudgins and P.L. Silveston, Can J. Chem. Eng., 45(5)(1967)294-299 [I 0] J. Happel, M.A. Hnatow and R. Mezaki, Adv. Chem. Ser.,(1970)92-109 [11] S. Carra and L. Forni, Catalysis Review, (1971)159-198 [12] J. Masson, J.M. Bonnier and B. Delmon, J. Chimie Phys., 76(5)(1979)458-464 [13] Lyu Cam Loc, N. A.Gaidai, B.S. Gudkov, S.L. Kiperman and S.B. Kogan, Kinet. Katal, 27(6)(1987)1184-1189 [ 14] Lyu Cam Loc, N. A.Gaidai, B.S. Gudkov, M.M. Kostyukovskii, S.L. Kiperman, N.M. Podkletnova, S.B. Kogan and N.R. Bursian, Kinet. Katal, 27(6)(1987)1190-1196 [I 5]Lyu Cam Loc, N. A.Gaidai, B.S. Gudkov and S.L. Kiperman, Proc. 9th Int Congr. Catal., 3(1988)1261-1267 [16] Lyu Cam Loc, N. A.Gaidai, S.L. Kiperman, N.M. Podkletnova and S.B. Kogan, Kinet. Katal, 29(5)(1988)989-996 [17] Lyu Cam Loc, B.S. Gudkov, N. A.Gaidai, S.L. Kiperman, Ho Sci Thoang and S.B. Kogan, Proc. Int Conf. (1990) 385-390 [! 8] Lyu Cam Loc, N. A.Gaidai, S.L. Kiperman and S.B. Kogan, Kinet. Katal, 31(2)(1990)421-424 [19] Lyu Cam Loc, B.S. Gudkov, N. A.Gaidai and S.L. Kiperman, Kinet. Katal, 31(3)(1990)541-545 [20] P. Meriaudeau, C. Naccache, A. Thangaraj, C.L. Bianchi, R. Carli and S. Narayanan, J. Catal., 152(1995)313 121] D. Casanave, (1996), PhD thesis, universite Claude Bernard Lyon- 1, Lyon, France [22] G. Froment, AICHE J., 21 (6)(1975) 1041-1056 [23] S.C. van der Linde, T.A. Nijhuis, F.H.M. Dekker, F. Kapteijn and J.A. Moulijn, App. Catal. A:General, 151(1997)27-57 [24] J.R. Katzer, in Chemistry and chemical engineering of catalytic processes, Eds R. Prins and G.C.A. Schuit, Sijthoff and Noordhoff, 1980, serie E: Applied Sciences 39, 3-48.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science Ltd. All rights reserved.

375

Synthesis of Ethylbenzene from Diethylbenzenes In the presence of Benzene using Triflic Acid as Catalyst Mohammed C. AI-Kinany and S0iiman H. Ai-Khowaiter Petroleum & Petrochemicals Research Institute, King Abdulaziz City for Science and Technology, P. O. Box 6086, Riyadh- 11442, Saudi Arabia. ABSTRACT In order to produce higher yield of ethylbenzene from diethylbenzenes which are resulted from the processes of ethylation of benzene using zeolite, Bronsted and Lewis acids as catalysts, a series of isomerization and transalkylation reactions of o-, p-, and m-diethylbenzenes in benzene with different molar ratios 1"1 and 1:6 with each isomer respectively using trifluoromethanesulphonic acid (triflic acid) as catalyst were caried out in liquid phase under atmospheric pressure and room temperature. In the case of 1:1"1 molar ratio of p-diethylbenzene to benzene to catalyst higher yield of ethylebnzene (ca. 51 mol%) were obtained, based on benzene and p-isomer after 6 h of the reaction time. With o-isomer (43 mol%) of ethylbenzene was produced. But the reaction between m-diethylbenzene with benzene under the same conditions gave lower yield of ethylbenzene (ca. 29 tool%). It was also found that when the molar ratio of isomer to benzene decreases the yield of ethylbenzene decreases. In the case of 1:6 molar ratio of p-isomer to benzene the yield of ethylebenzene decreased to (22.2 mol%). The same trend was found with the other two o- and m-isomers. The conclusion which can be drawn from these series of reactions is that the rate of transalkylation of diethylbenzenes in benzene with all molar ratios is faster than the rate of isomerization, and the rate of transalkylation isomers leading to higher yield of ethylbenzene with 1"1 molar ratio in the following order:

p-diethylbenzene

> o-diethylbenzene > m-diethylebnzene

Therefore, triflic acid as a catalyst to produce more yield of ethylbenzene from the products of ethylation reaction is more superior than the other catalysts such as solid super acids, zeolite, Bronsted acids, and aluminum chloride which is currently used in industries and leads to the formation of complexes with ethylated products in the catalyst layer. INTRODUCTION Alkylation of benzene with ethylene to produce ethylbenzene, an intermediate in the production of styrene, is well established process (15). The major byproducts formed during the reaction are polyethylated benzenes such as o-, p- and m-diethylbenzenes, triethylbenzene isomers and tetraethylbenzene isomers. Higher ethylated products, including pentaethylbenzene and hexaethylbenzene are also produced. Considerable attention has been given to this chemistry because large amounts of ethylbenzene were required to meet the increasing demand for styrene in this highly competitive market. The importance of ethylbenzene as a precursor to styrene and other petrochemical derivatives is shown in scheme (1).

376

~1 .J Styrene- Im~ Ethylbenzene

Poiystyren

I

i Styrene-ButadieneRubber(sBR:). I .

.

.

.

.

.

.

.

,~ Styrene-Butadiene- Acryionitrile , , J Stvrene_Acrvlonitrile__]

Scheme (1) The following catalysts have been used. 1. Metal halide catalysts (Lewis acids) : such as anhydrous aluminum chloride, boron trifluoride and titanium tetrachloride. 2. Protonic acid catalysts (Bronsted acids) : such as sulphuric acid, hydrofluoric acid, and super acids. 3. Solid acid catalysts" such as zeolites (6), and solid phosphoric acid. Transfer of an alkyl group from one to another molecules is the basis of the transalkylation reaction. This reaction is industrially of some interest as some of the low valued products like polyethylbenzenes can be converted to their monosubstituted homologues, which are in higher demand and of higher value. The transalkylation reaction, first demonstrated in 1894 (8), forms the basis of an important commercial synthetic process for the synthesis of ethylbenzene from fresh benzene and recycled higher ethylbenzenes in the presence of a metal halide (i.e. A1C13, BF3), protonic acid, or solid acids. Transalkylation is usually carried out on an industrial scale employing aluminium chloride (1) or trifluoroboron(5) at 100?C, and zeolite (2'3) catalysts at temperatures higher than (9) 200?C.Supported phosphoric acid catalysts are ineffective for this reaction. The diethylbenzenes, formed using this kind of catalyst, are not recycled for transalkylation, and to minimize their formation higher ratios of benzene to ethylene are often used. The transalkylation of diethylbenzene with benzene has also been studied over solid superacidic catalyst, such as Nafion-H(]~ proposed reaction mechanism, for transalkylation of diethylebenzene with benzene in the presence of Lewis acid is presented in scheme (2). Et

~Et + 0 _ . . , , _ 2 (

E

Scheme(2)

The other reaction involving disproportionation of diethylbenzene to triethylbenzenes and ethylbenzene is shown in scheme (3).

377 Et

Et

Et

Ethylbenzene Triethylben

Scheme (3)

Transalkylation of diethylbenzenes to obtain higher yield of ethylbenzene using the above and a wide range of catalysts has been intensively studied during the past decade. The major reactions include isomerization and transalkylation with their relative contributions controlled by the nature of catalyst and the reaction conditions. Considerable efforts have been aimed at the isomerization and disproportionation of diethylbenzene isomers, in the absence of benzene, to produce ethylbenzene by using trifluoromethanesulphonic acid catalyst (7). The present work was carried out to investigate the efficiency of trifluoromethanesulphonic acid as a catalyst for isomerization and transalkylation of diethylbenzene isomers in the presence of benzene at room temperature in order to synthesise higher yield of ethylbenzene. EXPERIMENTAL

Materials : Benzene (Sigma Chemicals Ltd., 99.9%) was purified according to the standard method, and dried prior to use over sodium wire. It's purity was checked by Gas Liquid Chromatography (G.L.C.) analysis and found 99.99%. Trifluoromethanesulphonic acid was a commercial sample (Fluka Chemie 98% and d=1.696) purified by double distillation under dry nitrogen at reduced pressure immediately before use. Hexane (Fluka Chemie) which was used as solvent for reaction product extraction was high grade and spectroscopically pure. Its purity was checked by G.L.C. to avoid any confusion in retention time between impurities in this solvent and the mixture of the isomerization and transalkylation products. o-Diethylbenzene , p-diethylbenzene and m-diethylbenzene ( all Fluka 97%) were used directly without further purification. General Procedures : Isomerization and transalkylation reactions were carried out in a closed system with continuous stirring under dry nitrogen, atmospheric pressure, and room temperature. Sampling: G.L.C. analysis of the reaction mixtures was carried out using a Varian 3400 instrument fitted with a 60m x 0.32 mm ID capillary column (phase DB 1, film thickness 1.0 micron) and flame ionization detector (FID). The G.L.C. instrument was calibrated by extemal standardization method using standard mixtures of benzene, ethylbenzene, o-, p-, and m-diethylbenzenes, 1,3,5-triethylbenzene and 1,2,4,5-tetraethylbenzene.

RESULTS AND DISCUSSION One of the characteristic features of the ethylation of benzene with ethylene is that an ethyl group is transferred readily either from one position to another in the same ring, or from one benzene nucleus to another. It results in the formation of higher ethylated products under ethylation conditions.

378 Generally speaking, the extent of isomerization and transalkylation which occurs during the ethylation reaction is determined mainly by the severity of the reaction conditions, such as, the type of catalyst and co-catalyst used, the ratio of catalyst to benzene substrate, the solubility of ethylated products in the catalyst layer, the reaction temperatures and pressures, and the reaction time. In our previous investigations (7), anhydrous trifluoromethanesulphonic acid (triflic acid) was found to be a very effective catalyst at room temperature and atmospheric pressure for isomerization and disproportionatof o-, p-, and m- diethylbenzenes in the absence of benzene.The higher yield of ethylbenzene (ca.50 mol%)was obtained with o-isomer after 15 min.of the reaction time. It is much more efficient than other catalysts such as, aluminium chloride promoted with water (li) and Y type zeolite (12) It was of some interest to investigate the efficiency of trifluoromethanesulphonic acid as a catalyst at room temperature and atmospheric pressure for isomerization and transalkylation of o-, p-, and m-diethylbenzenes with benzene in order to produce higher yield of more valuable ethylbenzene. The following set of reactions were carried out at room temperature with a mole ratio of each isomer to benzene to catalyst, 1:1:1, and 1:6:1.

Isomerization and Transal~lation of Diethylbenzenes with Benzene Using a closed system at room temperature and nitrogen atmospheric pressure, a set of experiments was carried out in which o-, p- and m-diethylbenzene to benzene to triflic acid molar ratio was 1:1:1 and 1:6:1 respectively. In all cases the product distribution of isomerization and transalkylation reactions was followed by analysing samples of the organic layer at 0.5 h intervals.

With o-diethylbenzene and benzene Reaction between o-diethylbenzene, benzene and trifluoromethanesulphonic acid with a molar ratio 1"1"1 respectively at room temperature caused rapid transalkylation to ethylbenzene and triethylbenzene as well as isomerization to m-, and p- diethylbenzenes as shown in figure (1). It can be seen that after 15 min. (14.63%) of ethylbenzene, (13.08%) of m-isomer, (2.88%) of p-isomer and (0.61 mol%) of triethylbenzene were formed.When the mole percentage values of diethylbenzene isomers were normalized, it can be seen from 60 m-Diethylbz o-Diethylbz

50

o~cl~n~o

20

10

p-Diethylbz | ,2,3-Triethylbz

~~

~x~.--~----___~~x-~

~

~~•215

0 0

1

2

3

4

5

6

7

TIME (HR) Fig. (1) p r o d u c t d i s t r i b u t i o n f r o m i s o m e r i z a t i o n a n d t r a n s a l k y l a t i o n o f o - d i e t h y l b e n z e n e w i t h b e n z e n e (1:1 m o l a r ratio)in the p r e s e n c e o f a n h y d r o u s triflic acid c a t a l y s t at r o o m t e m p e r a t u r e .

figure (2) that isomerization of o-isomer gives predominantly the m-isomer during the early stages of the reaction, and the m-isomer isomerizes to the p-isomer as expected of the 1,2-shift mechanism (Scheme 4).

379 Et

Et

'~176 ~

~o ~k~o-isomer

4

Et ~

.+, ,r

~L

Et ~1~

Et

";-~~~r,-+-,1 _"+ ]~,~.

I

II ~

~o4~ -~/ \.~ ~ ~. ~ . + ~ 60

-~/

+/+

+

1

2

Et

-

I , + ~1 H

+

* o-diethylbz

Et

"

- .

~_~+~

+~+---_+

4

5

0 0

3

6

7

TIME (HR) Fig. (2) N o r m a l i z e d d i e t h y l b e n z e n e isomers distribution fr om i s o m e r i z a t i o n a n d t r a n s a l k y l a t i o n of o - d i e t h y l b e n z e n e with b e n z e n e (1:1 m o l a r ratio)in t h e p r e s e n c e of a n h y d r o u s triflic acid ca ta lyst at r o o m t e m p e r a t u r e .

With p-diethylbenzne and benzene At 22~ with a p-isomer to benzene to triflic acid molar ratio 1"1"1, the mole percentage of ethylbenzene was (21.56%) after 15 min. and rose slowly to (50.91%) over further 6 h in the reaction mixture, as shown in figure (3). Under these conditions the rate

y, /,(

"

1

:,____ ........................

2

3

4

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

5

6

TIME (HR) Fig. (3) p r o d u c t d i s t r i b u t i o n f r o m i s o m e r i z a t i o n a n d t r a n s a l k y l a t i o n o f p - d i e t h y l b e n z e n e w i t h b e n z e n e ( 1 : 1 m o l a r r a t i o ) i n the p r e s e n c e o f a n h y d r o u s triflic a c i d c a t a l y s t at r o o m t e m p e r a t u r e .

of transalkylation to ethylbenzene and triethylbenzene is faster than that of isomerization to o-, and m-diethylbenzene. So, for example, when the mole percentage values of diethylbenzenes normalized, as shown in figure (4) the composition of the isomers changed

380 120 9m-Diethylbz + p-Diethylbz

100

k+~+,

* o-Diethylbz

.

40

.

.

.

.

_

.

.

.

.

.

.

~F~_F,.~..__+ ~ + _ . _ . . _ + ~ + ~ +

.

...__._+

/ S ~ , ,*-/-v-,~ * _ _

0

,T-~-*v-:-, . ~ - - * ~ - , ,t-7--*v-:-, ~ * : - ~ - . t~ . . . .

1

2

3

4

5

6

TIME (HR) Fig. (4) Normalized diethylbenzene isomers distribution from isomerization and transalkylation of p-diethylbenzene with benzene (1:1 molar ratio)in the presence of anhydrous triflic acid catalyst at room temperature.

from o-(1.23%), p- (28.75%) and m- (68.77%) after 1 h. It can also be seen from figure (4) that (a) the rate of isomerization to the thermodynamically more stable m-isomer is faster than that to o-isomer; (b) the isomerization of p-isomer gives predominantly m-isomer during the early stages and over further 6 h of the reaction and; (c) the m-isomer then ultimately isomerizes to o-isomer as expected from the following 1,2- shift mechanism scheme (5). Et

Et

fast 1.2-Shlft

Et

Et Et

Et

/L 1.2-Shill

p-isomer Et

H

o-isomer Scheme (5)

With m-diethFlbenzene and benzene Under the same conditions, the reactions of m-isomer, benzene and triflic acid with a molar ratio 1"1"1 respectively caused very slow transalkylation to ethylbenzene and triethylebenzene as well as isomerization to o-, and p- isomers as shown in figure (5). 9B e n z e n e o m-Diethylbz ~ o-Diethylbz

" ~ . _ ~

* Ethylbz x p-Diethylbz ~ 1,2,3-Triethylbz

40 _1 o

=E

30

/ 0

1

2

3

4

5

6

TIME (HR) Fig. (5) p r o d u c t d i s t r i b u t i o n f r o m i s i m e r i z a t i o n a n d t r a n s a l k y l a t i o n o f m - d i e t h y l b e n z e n e (1 "1 m o l a r ratio)in t h e p r e s e n c e o f a n h y d r o u s t r a f i c acid catalyst.

381 Both reactions were much slower than those of o-, and p-isomers. The reaction mixture after 15 min. contains benzene (50.28%), ethylbenzene (4.43%), m-, p-, and o-isomers (44.38%), (0.54%) and (0.18%) respectively and 1,2,3-triethylbenzene (0.14 mol%). The mole percentage of ethylbenzene in the reaction mixture was (4.43%) after 15 min. and this rose slowly to only (29.0%) after 6 h. As with o-, and p-isomer the rate of transalkylation to ethylbenzene and triethylbenzene was faster than the rate of isomerization to o-, and pisomers to some extent. The rate of isomerization of m-isomer to o-isomer is slower than that to pisomer.This is seen more clearly in figure (6), where the normalized mole percentages of the three diethylbenzene isomers are presented. 120 100

= - " ' - - - - ~ . _____~...______

80 m-Diethylbz + p-Diethylbz o-Diethylbz

60 40 20 + _____ + _ _ _ _ _ + _____ + _____ + - - - - - - + - - - - -

0

1

2

3

+

4

TIME

+ 5

+

.+ 6

(HR)

Fig. (6) Normalized diethylbenzene isomers distribution from isomerization and transalkylation of m-diethylbenzene with benzene (1:1 molar ratio)in the p r e s e n c e of anhydrous triflic acid catalyst at room temperature,

It might have been predicted on steric grounds and intermediates that the p-isomer would be the product of isomerization by a 1,2- shift scheme (6). Et +d__ Et m-isomer

--1'2-Shift ~ E t H Et

~J II

- -

fast 1,2-Shift

Et

0-is0mer

Et

-a+-0

Et

p-isomer

Scheme (6)

In various results obtained by Olah (l~ in his study of the isomerization and transalkylation of m-diethylbenzene with tenfold excess of benzene in the presence of superacidic perfluorinated resin sulphonic acid (Nation-H) catalyst. The mixture when passed in the gas-phase and atmospheric pressure over the catalyst at 193~ was found to isomerize and transalkylate. Mono-, di- (o-, m-, and p-), and triethylbenzenes with composition of (30.6%), (1.5%), (63.2%), (3.9%) and (0.8%) respectively were the only products obtained under the experimental conditions. Benzene could not be determined in the reaction mixture. The only yield of ethylbenzene obtained was more consistent with ours after 6 h.

382 under the best operating conditions i.e. at room temperture. This indicates that the activity of triflic acid significantly exceeds those other solid superacidic catalyst Lewis acids and zeolite (12) at room temperature. From the above experiments, figure (7) shows the mole percentages conversion of o-, p-, and m-diethylbenzene isomers to ethylated products by isomerization and transalkylation reactions. It can be seen that the rate of conversion decreases in the following order: o-diethylbenzene 100

I

> p-diethylbenzene > m-diethylbenzene

.--------'~'~-+-----+------+--+--+-----+'----+--+

/

/I.+I--+

§

.

6o 1 s

- o-Diethylbz

4O

2O

0

0

1

2

3

4

5

6

7

TIME (HR) Fig.(7) conversion o f diethylbenzenes from isomerization and transalkylation o f diethylbenzene isomer with benzene ( 1 : 1 molar ratio ) in the presence o f anhydrous triflic acid atroom temperature.

Figure (8) shows the mole percentages of ethylbenzene in the reaction mixtures of isomerization and transalkylation of the three isomers of diethylbenzene. It can be seen that higher yield of ethylbenzene was obtained with p-isomer (ca. 51%) and it decreases in the following order. p-diethylbenzene > o-diethylbenzene > m-diethylbenzene

6O 5O ,o

t -

___...-+I+ "/"

/ i

______._.t-

.

20

10

0

0

1

2

3

4

l

- Ethylbz from o-Diethylbz I Ethylbz from p-Diethylbz Ethylbz from m-giethylbz 5

6

7

TIME (HR) Fig.(8) yield o f e t h y l b e n z e n e from the reaction mixtures o f i s o m e r i z a t i o n and transalkylation o f d i e t h y l b e n z e n e isomer with b e n z e n e (1 : I molar ratio ) in the presence o f anhydrous triflic acid at r o o m temperature.

Effect of Molar Ratio on the Yield of EthFlbenzene In order to investigate the effect of molar ratio of isomer to benzene on the conversion of isomer and the yield of ethylbenzene, the mole ratio of isomer to benzene was varied from 1:1 to 1:6 at room temperature. Figures (9 & 10) show the effect of variation of the molar ratio of diethylbenzene isomer to benzene on the conversion of isomers and the yield of ethylbenzene respectively. There is a marked decrease in the conversion of isomers

383 to ethylated product with a decrease in the molar ratio. With o-isomer, the maximum convrsion of (98.63%) and maximum yield of ethylbenzene (43%) were observed at a 1:1 molar ratio after 6 h of reaction, while at 1:6 molar ratio the maximum conversion of (94%) and maximum yield of ethylbenzene (5.3%) were observed after the same reaction time. Same trends were obtained with p- and m-isomers as shown in figures 9 and 10. It can be seen that the yield of ethylbenzene based on the conversion of isomers decreases in the following order:

,2o[

p-diethylbenzene > o-diethylbenzene > m-diethylbenzene 100

+ ~ + ~ + ~

~___+__.__.+___~ ~+~.___.t---+

~o /111

6o

I-o-O,e..~,~z

IlU

I + p-Diethylbz

4O 2O 0 0

1

2

3

4

5

6

7

TIME (HR) F i g . ( 9 ) c o n v e r s i o n o f d i e t h y l b e n z e n e s f r o m i s o m e r i z a t i o n and t r a n s a l k y l a t i o n o f d i e t h y l b e n z e n e i s o m e r w i t h b e n z e n e (1:6 m o l a r ratio ) in the p r e s e n c e o f a n h y d r o u s triflic a c i d at r o o m temperature.

- Ethylbz from o-Diethylbz

Ethylbz from p-Diethylbz Ethylbz from m-Diethylbz

4O _1 O

30

_..._~+~+~+~+

20 10

j~_-

o .Z~ 0

~ - - : ~ - ~ 1

2

,.=-~*-~*__-7-;-,*-:-7,*.7", 3

4

5

6

, , ,

7

TIME ( H R ) Fig.(10) y i e l d o f e t h y l b e n z e n e f r o m the r e a c t i o n m i x t u r e s o f i s o m e r i z a t i o n a n d t r a n s a i k y l a t i o n o f d i e t h y l b e n z e n e i s o m e r w i t h b e n z e n e (1:6 m o l a r ratio ) in t h e p r e s e n c e o f a n h y d r o u s triflic a c i d at r o o m temperature.

Reusability of the Cata!Fst Reusability of the catalyst is commercially important. Considering this, we checked the recycling and reusability of triflic acid catalyst. In this case the catalyst was recovered as described in our previous work (13), and reused by repeating one of the above experiments (with o-, 1"1 molar ratio). The Results were similar to the results of a fresh catalyst.

CONCLUSIONS Trifluoromethanesulphonic acid was found to be a promising superacid catalyst for isomerization and transalkylation of diethylbenzene isomer with benzene at room temperature to produce higher yield of ethylbenzene. Its catalytic activity is comparable

384 with the other catalysts such as solid superacid, zeolite, Lewis acids and Bronsted acids. ACKNOWLEDGEMENT

The authors would like to thank Mr. Fahad H. AI-Malki for his assistance in the experimental works.

REFERENCES

1. 2. 3. 4. 5. 6. 7.

A.C. McFarlane (Monsanto Co.), Oil and Gas J., 99, 1976. F.G. Dwyer, and P. J. Lewis: Chem. Eng., 83(1), 90, 1976. F.G. Dwyer, P. J. Lewis, and F. H. Schneider: Oil and Gas J., 75, 55, 1977. E.K. Jones, Oil and Gas J., 58(9), 80, 1960. H.W. Grote (UOP C.), Oil and Gas J., 56, 73, 1958. J. Butler, J. Waguespack and K. Hall, EP 726242 A1 19960814. M. C. A1-Kinany and S. H. A1-Khowaiter, Proceedings of 2nd Middle East Refining and Petrochemicals Conference and Exhibition, PETROTECH'98, Bahrain, September 1998, vol. 1. P.523-42. 8. C. Radziewanoki, Ber, 27, 3235, 1894. 9. E.K. Jones, Oil and Gas J.. 58(9), 80, 1960. 10. G. A. Olah, and J. Kaspi, Nouveau J. DE Chieme, 2(6), 585-591, 1978. 11. G. A. Olah, W. Max, M. W. Meyer and N. A. Overchuck, J. Org. Chem., 29, 2313, 1964. 12. J. Rane, Shashikala, C. V. V. Satyanarayana, D. K. Chakrabarty, Appl. Catal., 69(2), 177-86, 1991. 13. B. L. Booth, M.C. A1-Kinany and (in part) Khosrow Laali, J. of Chem. Soc., Perkin. Trans,., 1, 2049, 1987.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

385

Elementary steps of reaction pathway in the pilot plant phot0mineralisation of s-triazines on to photocatalytic membranes im~obilising titanium dioxide and promoting photocatalysts

Alberto Moroni a, Ignazio Renato Bellobono a,** and Bernd M. Gawlik b aDepartment of Physical Chemistry and Electrochemistry, University of Milan; via C. Golgi, 19; 1-20133 Milano (Italy) bjoint Research Centre of the Commission of the European Union; E n v i r o n m e n t Institute; 1-21020 Ispra (VA) (Italy)

Abstract

The T i O 2 - m e d i a t e d p h o t o o x i d a t i o n of atrazine and prometryn in 5.0x10-6-2.0x10 -4 M aqueous solutions on to PHOTOPERM | CPP/313 membranes immobilising 30• wt.% Ti02, with or without 6 wt.% of a synergic mixture of tri-(t-butyl)and tri-(ipropyl)vanadate(V), as p h o t o p r o m o t e r c o - i m m o b i l i s e d in to the membrane, and by employing hydrogen peroxide or ozone as oxygen donors, was studied kinetically. Experiments were carried out at 2 9 6 • in a P H O T O P E R M | WP pilot plant, using monochromatic irradiation (254 nm) in the absorption range of semiconductor, with an absorbed power of 11• W. Pseudo-first order kinetics was observed in the first step of transformation of substrates into intermediates, following a L a n g m u i r - H i n s h e l w o o d behaviour of initial rates. The final product was the relatively photostable 2,4,6-tri-hydroxy-s-triazine. By using s u b s t o i c h i o m e t r i c hydrogen peroxide, however, or much better ozone, as oxygen donor, this latter intermediate was completely mineralised. Integral p h o t o m i n e r a l i s a t i o n was even more fast, by a factor of about 10, when using the photopromoted p h o t o c a t a l y t i c membranes. Furthermore, based on HPLC analysis, a m e c h a n i s m of p h o t o m i n e r a l i s a t i o n is proposed, and the role of oxidising radicals discussed.

1. INTRODUCTION s-Triazine derivatives are among the most widely used herbicides, to control broadleaf and grassy weeds, as well as other crops. Their environmental presence both in surface and groundwaters has been ascertained [1 ]. If an assessment * Part 69 of the series "Photosynthetic Membranes" ** Author to whom c o r r e s p o n d e n c e should be addressed.

386

approach similar to that adopted by the World Health Organization will prevail [2], the limit values for s-triazine herbicides in human drinking water should be of a few ppb. This will require advanced water purification systems, among which photocatalytic technologies, particularly those based on the use of titanium dioxide semiconductor, have received great attention [3]. The photocatalysed degradation of atrazine, and other s-triazines, however, when using TiO 2 suspensions, has been reported to present no destruction of the heterocyclic ring, the final product observed being 2,4,6-tri-hydroxy-striazine (cyanuric acid), photochemically stable in the oxidative conditions tested [4]. In previous papers of this series [5-7], on the contrary, in which immobilisation of massive amounts of semiconductor has been carried out onto photocatalytic, PHOTOPERM membranes (Chimia Prodotti e Processi, Cinisello Balsamo; Milan) prepared by photografting, to which appropriate photocatalytic promoters could be added in the composite structure, practically quantitative photomineralisation of atrazine and other s-triazines has been achieved during pilot-plant experiments. By this way, a technological and economic solution for these photocatalytic processes has been found, able to exploit at the same time all known advantages of membrane technology (modularity, optimal photoreactor modeling, continuous processing, and so on). At the same time, also a very efficient method resulted, to tailor the photocatalytic process, by co-immobilising other photocatalysts or other photocatalytic promoters or other catalytic species, as desired, thus adjoining the pertinent performances. As to the unsuccess in getting complete mineralisation when using suspended semiconductor, two main reasons have been preliminarly explored and clarified [7-8]: on one side the scarce efficiency of TiO 2 in aqueous suspensions, as documented by the very limited quantum yields reported in the literature (the same holds for unsufficiently thin and compact layers of semiconductor immobilised on surfaces), and on the other side the mechanism. When based exclusively, or almost exclusively, on reaction with -OH radicals, such as it occurs if low efficiency semiconductor is employed or if hydrogen peroxide is used as oxygen donor, cyanuric acid, as the final degradation intermediate, results. In the present paper, attention was focused on the kinetic role of intermediates during photodegradation onto the photocatalytic membranes, with the purpose of identifying the relevant pathways, and their possible elementary steps, as well as of investigating systematically the reasons of the very satisfactory performance of the membrane process, as regards integral photomineralisation. 2. E X P E R I M E N T A L 2.1.

Materials.

Standard herbicide and its degradation intermediates, as well as other chemicals, were the same used in previous studies

387

[6-8]. No buffer system was added; 5.8 during irradiation experiments. 2.2.

Photocatalytic

membranes

pH

ranged

and i r r a d i a t i o n

from

about

5 to

experiments

Photocatalytic membranes (PHOTOPERM | CPP/313) were supplied by Chimia Prodotti e Processi (Cinisello Balsamo, Milan, Italy). The standard membranes immobilised 30• wt.% TiO2; the photopromoted membranes co-immobilised 6 wt.% of a synergic mixture of tri-(t-butyl)and tri-(ipropyl)vanadate(V), as photopromoter. Irradiation experiments were performed in a PHOTOPERM | WP pilot plant, already described [9], using monochromatic irradiation (254 nm) in the absorption range of semiconductor, with an absorbed power of ii• W, as measured actinometrically. Mean temperature, during the runs, was regulated at 296• K. A solution volume of 20 L was used in all experiments, containing 5.0x10-6-2.0x10 -4 M atrazine [CAS: 1912-24-9] or prometryn [CAS: 7287-19-6] in deionised water, to which a stoichiometric amount of hydrogen peroxide was generally added, as oxygen donor. In some sets of kinetic runs, however, less or greater than stoichiometric hydrogen peroxide was used; and in some others ozone was exclusively employed, as described [5-7]. Analytical methods, concerning total organic carbon (TOC) determination, and HPLC quantitation of intermediates, were the same already described [6-8]. Organic nitrogen analysis was also carried out spectrophotometrically in the present paper by the Kjeldahl method and the Nessler reagent. Absorbable organic halogens (AOX) analysis was carried out by active carbon absorption, fol• by successive combustion of the sorbent at 950~ in oxygen atmosphere, and microcoulommetric titration of resulting halogenide ions, in an AOX 70.10 analyser (Dani, Monza, Milan, Italy). 3. RESULTS

AND D I S C U S S I O N

As had been observed in previous work, on related striazines, but employing oxygen or ozone as oxygen donors [57], in a range of low concentrations such as that examined in the present paper, the observed kinetics of substrate disappearance were of apparent first order. Furthermore, in the present work, by comparison of HPLC and AOX analysis the percentage of reaction leading to dechlorination of substrate in the first reaction step could be calculated as 30• in the case of atrazine, leading to intermediate I3 (see route b in Scheme i), while the main route, covering 70• of the whole first reaction step (a in Scheme i), through the detected intermediates I4 and I5, formed from I2 and I3, was able to yield the detected intermediate I7, either directly (by routes h and g of Scheme i), or indirectly, by passing through the detected intermediate I6, which was quantitatively converted into I7 (route d in Scheme I). As, in the presence of ozone and by photopromoted membranes, the successive intermediate, for

388 Scheme

1

Cl

N~ N

ATRAZINE

b1:3o%

=70%

N~/~NI1 C1

OH

OH

Cl

12 OH

N"~N

'

; ...................................................................................... ! .

el 1 N~N

1[4

.

.

.

.

.

.

.

.

.

cl I

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

g h

T

I N~N

NH2~~~ OH

I

N ~ N

HO ~ ~ ~

~o~

.

r

C1

N~N

.

15

N~N

c~

HO

.

OH

16 NH 2

~_ 1 0 0 %

ei N~-.N

o~

h+

2

OH

............................................. ~.

HO

:-I--~

T

02-"

N~

OH

complete mineralization

17 NH 2

OH

I9

Scheme +

I N~N

d

.41.............................................

HO

][8

389 which detection was possible, resulted in isolation of cyanuric acid (19). Its formation may be well rationalised, by route e of Scheme I, and the possible occurrence of intermediate 18. The further degradation of 19 (see Scheme 2) may be hypothetised, on the basis of the oxidising ability of holes, as proposed in the literature [i0], even if for substrates different than cyanuric acid. Contrarily to runs carried out with oxygen or ozone, however, when operating in the presence of hydrogen peroxide, even after prolonged irradiation (> I000 min), the final presence of cyanuric acid, which resisted further oxidation under these experimental conditions, could be detected for both herbicides tested. This had also been remarked in some of the previous experiments on to photocatalytic membranes with simazine [7] in the presence of hydrogen peroxide. When the latter was used, mineralization stopped when a TOC decrease correspondent to 4/7 with respect to the initial concentration was attained [7]. This is also in agreement with the results of other photodegradation experiments, using mainly oxygen and titanium dioxide suspensions, performed with various striazines [4, 11-12]], with the difference, indeed, that on to photocatalytic membranes, in the presence of only dissolved oxygen [5, 7], mineralization was still complete. A first explanation to this behaviour has been given [7], and is confirmed by the present data. It is well known [12] that the prevailing radical species acting in the mechanism of photocatalytic degradation are .OH radicals, when hydrogen peroxide is present, so that the lack of any reaction between 2,4,6-tri-hydroxy-s-triazine and these radicals is inferred. This is very plausible also on the basis of the chemical structure of cyanuric acid. With oxygen or ozone, on the contrary, on to the very active photocatalyst immobilized within the membrane pores, rate of scavenging of electrons in the conduction band is strongly enhanced. By this way, not only the probability of their recombination with holes is decreased or radically suppressed (the latter situation is observable when titanium dioxide on to the membrane is activated by promoting photocatalysts [7]), but the production of .O2H radical species (in acid-base equilibrium with 02-- ) is also possible, by reaction with dissolved oxygen or ozone, thus leading to further degradation of cyanuric acid. Apparently, this does not occur with aqueous suspensions of titanium dioxide, while in the presence of hydrogen peroxide, even on to the more reactive photocatalytic membranes, the formation of reactive .O2H radicals is not very favoured. Further evidence along this line of interpretations stems from the apparent first order dependence of observed kinetic contants for prometryn [8] and atrazine in the present work, as stated above. Degradation experiments for both herbicides were carried out at different initial concentrations C0, in order to evaluate the concentration dependence of initial rates of degradation. As has been done in previous work [5-9, 13-14], the Langmuir-Hinshelwood model and the corresponding kinetic rate law, combining apparent adsorption equilibrium and apparent zero-order surface reaction, was used to interpret

390 initial rate data form of eqn. (I):

of

experimental

results,

in

the

linearised

(1)

I/r 0 = (l/k) + (i/k KC0)

where r 0 represents the rate of decrease of the pertinent species as a function of time, C O the initial concentration of the organic compound, which is being photodegraded, k the reaction rate constant, e x p r e s s e d in mol/min, in order to make it independent on the solution volume, and K (M -I) the apparent adsorption constant of the substrate on to the p h o t o c a t a l y t i c membrane. The validity of this model, as has been proposed, in their criticism of the formal LangmuirHinshelwood equation, by Turchi and Ollis [15], and amply d e m o n s t r a t e d experimentally, also by our own work [5-9, 13-14, 16], should not be stressed in the strict sense above. While k is a kinetic constant d e p e n d e n t on the properties of the catalyst and on the reaction conditions, K has also, as a matter of fact, partly at least, a kinetic meaning, being d e p e n d e n t on the m e c h a n i s m of reaction, as well as on operative conditions of the photoreactor, such as flow rate, c o n c e n t r a t i o n of oxygen suppliers, irradiation wavelength [16]. By calculating, from curves, such as those of Fig. i, the apparent first order kinetic constants kobs, and by plotting them as a function of initial concentrations, the curve of Fig. 1 was o b t a i n e d for the experimental data of atrazine.

0,042 0,040 ,--,r

0,038

"~ 0,036

,,, 0,034 r

o

-~ 0,032

/

0,0300,028 0,026 . 0,00 0,02 i

,

i

,

i

,

i

0,04 0,06 0,08 C o / m g L1

,

o, io ' o, 2

Fig. i. D e p e n d e n c e of apparent first order rate kobs, on the initial c o n c e n t r a t i o n of atrazine.

constants,

A possible e x p l a n a t i o n of this dependence may be the fact that k o b s i n c l u d e a term r e f e r r i n g to the concentration of *OH. The latter, in the experimental conditions of this work, were

391

produced both by reaction of the photogenerated holes on the semiconductor with water and by direct photolysis. As the concentration of hydrogen peroxide increased with increasing Co, on one side the concentration of -OH radicals, produced photolytically, increased with increasing C O . On the other side, however, due to the reaction of hydrogen peroxide with holes generating molecular oxygen, the production of semiconductor driven -OH decreased, thus generating the profile of Fig. i, above a critical concentration of substrate. In the range of concentrations, for which the behaviour of atrazine was regular (increase of kob s with increasing CO) , by eqn. (i), the k and K parameters have been calculated by regression analysis, for the first degradation step, in the presence of hydrogen peroxide as oxygen donor, and with the standard photocatalytic membranes. These values are reported in Table I, where they are compared with the corresponding values for prometryn [8], and with the values, for these same herbicides, measured in the presence of ozone as oxygen donor, and with photopromoted membranes containing the trialkyl vanadates additives, in the present work. TABLE 1 Parameters k and K, according to eqn. (i), (uncertainties expressed as standard deviations) for the first step of photodegradation of prometryn and atrazine in aqueous solution (volume of irradiated solutions 20 L), in the presence of stoichiometric hydrogen peroxide, by standard photocatalytic membranes PHOTOPERM | CPP/313, and in the presence of ozone, by photopromoted membranes, in pilot-plant experiments using monochromatic irradiation (Ii• W of the emitted energy at 254 nm was effectively absorbed by the membranes in the absorption range of semiconductor). Standard membranes and H202 Prometryn Atrazine

[8]

Photopromoted membranes and 03 Prometryn Atrazine

k (mol/min)

K ( M -I)

(5.6+0.5)X10 -4 (5.0+0.4)X10 -4

(1.8+0.4)xi03 (1.3+0.3)xlO 3

(2.0+0.2)x10 -2 (2.3+0.2)x10 -2

294+53 195+65

By considering that the absorbed power, in the experimental conditions of this work, corresponded to 1.4x10 -3 Einstein/min, quantum yields for the photodegradation of the two substrates (first step of their transformation into intermediates) may be easily evaluated at "infinite" concentration, by the ratio between the k values of Table 1 and photonic flux above. In these conditions, availability of substrate is such as to inhibit recombination of -OH radicals

392

to form hydrogen peroxide, a process which obviously depresses photonic efficiency for the attack of substrate molecules by these radicals. The quantum yield values above are quite satisfactory, two order of magnitudes greater than the literature values. This has been amply discussed in preceding papers, relatively to s-triazines [5-8], in which the positive influence brought about by trialkyl vanadates in the membrane, has also been underlined. By the latter, indeed, overall quantum yields greater than the maximum theoretical ones may be reached, owing to the fact that, by using ozone as oxygen supplier, the vanadate additives exhibit a dark catalytic effect, which contributes very efficiently to integral mineralisation, even of the most sluggish intermediates [6-8]. 4. R E F E R E N C E S

1 2 3 4 5 6 7 8 9

i0 ii 12 13 14 15 16

A. Di Corcia, R. Samperi, A. Marcomini and S. Stelluto, Anal. Chem.,65 (1993)907. World Health Organisation, Guidelines for Drinking Water Quality, vol.l: Recommendations, WHO, Geneva, 1993. D.F. Ollis and H. Ai-Ekabi (eds.), Photocatalytic Purification and Treatment of Water and Air, Elsevier, Amsterdam, 1993. E. Pelizzetti, V. Maurino, C. Minero, V. Carlin, E. Pramauro, O. Zerbinati, M.L. Tosato, Environ. Sci. Technol., 24, 1559 (1990). I.R. Bellobono, B. Barni and F. Gianturco, J. Membrane Sci., 102 (1995) 139. F. Gianturco, C.M. Chiodaroli, I.R. Bellobono, M.L. Raimondi, A. Moroni, B. Gawlik, Fresenius Environ. Bull., 6, 461 (1997). I.R. Bellobono, P.L. Pinacci, G. Riva, C. Lagrasta, Fresenius Environ. Bull., 7, 277 (1998). O. Borio, B.M. Gawlik, I.R. Bellobono and H. Muntau, Chemosphere, in the press. B. Barni, A. Cavicchioli, E. Riva, L. Zanoni, F. Bignoli, I.R. Bellobono, F. Gianturco, A. De Giorgi, H. Muntau, L. Montanarella, S. Facchetti and L. Castellano, Chemosphere, 30 (1995) 1847. N. Serpone, E. Pelizzetti and H. Hidaka, in Photocatalytic Purification and Treatment of Water and Air, (D.F. Ollis and H. Ai-Ekabi, eds.), Elsevier, Amsterdam, 1993, 225. M.C. Gonzales, A.M. Braun, A. Bianco Prevot and E. Pelizzetti, Chemosphere, 28 (1994) 2121. D.F. Ollis, E. Pelizzetti and N. Serpone, Environ. Sci. Technol., 25 (1991) 1523. I.R. Bellobono, A. Carrara, B. Barni and A. Gazzotti, J. Photochem. Photobiol., A:Chem., 84 (1994) 83. C. Lagrasta, I.R. Bellobono and M. Bonardi, J. Photochem. Photobiol., A:Chem., 119 (1997) 201. C.S. Turchi and D.F. Ollis, J. Catal., 122 (1990) 178. B. Barni, A. Cavicchioli, E. Riva, L. Zanoni, F. Bignoli, I.R. Bellobono, F. Gianturco, A. De Giorgi, H. Muntau, L. Montanarella, S. Facchetti and L. Castellano, Chemosphere, 30 (1995) 1861.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

Methane

oxidation

over

supported

nickel

393

catalysts

A.M. Diskin, R.H. Cunningham and R.M. Ormerod* B irchall Centre for Inorganic Chemistry and Materials Science, Department of Chemistry, Keele University, Staffordshire, ST5 5BG, United Kingdom *to whom correspondence should be addressed Abstract The catalytic oxidation of methane over supported nickel catalysts has been studied using conventional catalytic reactor measurements, temperature programmed reaction spectroscopy and gas pulsing experiments. The influence of support material, catalyst pre-treatment and operating temperature have been studied. The nature of the support material has a large influence on the subsequent activity and CO selectivity of the nickel catalysts. Temperature programmed measurements have been used to study methane activation, the surface reaction pathways and to evaluate the nature and level of any carbon species deposited during reaction. Temperature programmed oxidation reveals several types of carbon are formed on the catalyst during catalytic methane oxidation. Gas pulsing experiments have been carried out at different reaction temperatures to determine how the initial methane conversion, product selectivity and surface carbon yield vary as a function of contact time, and show that metallic nickel provides the active site for methane partial oxidation.

1. I N T R O D U C T I O N The conversion of natural ~,,as., containing mostly methane, to value-added products, such as easily transportable fuels, is driven by the tremendous abundance of natural gas in remote areas, as well as other economic factors. Thus in recent years the direct and indirect conversion of methane to value-added products has attracted a great deal of attention [1-8]. The direct conversion routes involve partial oxidation of methane to methanol, formaldehyde or alkenes. However, the high reactivity of the products relative to the reactant methane makes this approach difficult. The indirect routes involve oxidation of methane to syngas, either by steam reforming or partial oxidation or by a combination of the two, followed by conversion of syngas to upgraded products by either the methanol synthesis route or the Fischer-Tropsch process. Although steam reforming is predominantly used to convert methane to syngas [91, there are drawbacks to the process relating to the energy and capital costs of maintaining the reaction conditions of superheated steam, high temperatures and pressure. In addition, the water gas shift reaction produces significant quantities of carbon dioxide in the product gas, and the hydrogen to carbon monoxide ratio is higher than the optimum required for both methanol synthesis and the Fischer-Tropsch process !101. Deactivation of the nickel based catalysts through coking and sintering of the nickel particles is also a problem I111. Partial oxidation of methane to syngas represents a potential alternative to steam reforming. Compared to methane steam reforming, it is more energy efficient, potentially more selective and yields a lower H2:CO ratio which is more favourable for methanol synthesis and FischerTropsch processes. The strongly e~dothcrmic nature of steam reforming means that partial oxidation is of particular itnportance in fuel processing applications ,is the basis for start-up from cold. In solid oxide fuel cells self-sustained internal steam reforming is precluded during

394 start-up conditions and for operation at low power, as electrochemical losses are insufficient to meet both heat loss from the stack and the endothermic requirements of steam reforming, whereas partial oxidation offers the potential for both start-up and self-sustaining low power operation, as well as for remote and small-scale applications. Mixed partial oxidation and steam reforming (autothermal reforming) therefore provides the basis for operation from zero to full power. There have been numerous studies published on the partial oxidation of methane over supported metal catalysts [ 1-8,12-14]. It has been reported that high activities and selectivities to syngas can be obtained over supported catalysts, including nickel catalysts [5,14]. The actual mechanism of catalytic partial oxidation of methane and the nature of the active sites have been the subject of considerable debate [ 1,3,4,7,8,13,15]. In this short paper we describe a study of the catalytic oxidation of methane over nickel catalysts dispersed on various oxide supports. The influence of catalyst pre-treatment, the nature of the support material, operating temperature and methane/oxygen ratio have been studied. The CH4/O2 ratio of 4:1 used here is higher than in most studies of methane partial oxidation, where a 2:1 ratio is generally used, since low oxygen partial pressures are desirable for internal reforming in SOFCs. In addition to carrying out conventional catalytic reactor measurements, temperature programmed reaction spectroscopy (TPRS) has been used to study the methane activation process, the surface reaction pathways, and the nature and level of any carbon species deposited during reaction. Gas pulsing experiments measuring methane conversion and product selectivity as a function of contact time with the catalyst are also described. 2. E X P E R I M E N T A L 2.1

Catalyst p r e p a r a t i o n Nickel catalysts were prepared by impregnation of the oxide support to incipient wetness with an aqueous solution of nickel (II) nitrate (Fluka puriss.). All the support materials were dried at 403 K before use. Nickel catalysts dispersed on alumina, silica, ceria, titania and zirconia were prepared by incipient wetness with nickel loadings of 5 wt% and 10 wt%. Following preparation the catalysts were dried in air at 403 K overnight before being calcined in situ in a 10% O 2 ~ e stream at 873 K for one hour. For .catalytic measurements on prereduced catalysts, all samples were reduced in a 10% H2/~e stream at 953 K for one hour, with the exception of the Ni/CeO2 catalyst which was reduced at 1073 K. The catalysts have been characterised using X-ray diffraction, EXAFS and TPR. These results will be described in more detail elsewhere [ 16].

2.2 Catalytic experiments

All the experiments described here were carried out using a custom-built continuous flow stainless steel apparatus, consisting of a tubular fixed-bed reactor heated by a furnace over the range 293-1273 K. The reactor exit was connected to a continuously sampling mass spectrometer (Leda-Mass Satellite) via a heated quartz capillary to avoid water condensation. Linear temperature control and the continuous sampling nature of the mass spectrometer allowed both temperature programmed reaction measurements and gas pulsing experiments to be carried out, and any transient phenomena to be identified. Temperature programmed measurements were carried out using a heating rate of 10 K min -l. Temperature programmed reduction (TPR) and oxidation (TPO) experiments were carried out in 10% H~JHe and 10% O2/He gas mixtures, respectively. Temperature progranlmed partial oxidation of methane used a 4:1 CH4/O2 mixture diluted in helium. Gas pulsing experiments were carried out using a purpose-built pulsing loop. There was a constant helium flow through the reactor and the reactant gas mixture was flushed in by the carrier ~as ,,. -. The time interval between pulses was 5 rains.

395 3. RESULTS 3.1

T e m p e r a t u r e P r o g r a m m e d Oxidation of M e t h a n e Temperature Programmed Reaction Spectroscopy (TPRS) was used to study methane activation on the nickel catalysts and the catalytic oxidation of methane over the temperature range 300-1160 K using a 4:1 CH4/O2 ratio. Figure 1 shows a typical TPRS spectrum, obtained in this case for a 5 wt% Ni/SiO2 catalyst. It can be seen how the methane conversion increases with reaction temperature. This was found to be true for all the catalysts studied. The CO selectivity also increases with temperature, with a sharp increase in selectivity observed at high temperature; this occurred for all the catalysts studied except Ni/CeO2, where the CO selectivity was observed to pass through a maximum at about 1080 K. Figure 2 shows the selectivity towards CO formation over pre-reduced 5 wt% Ni catalysts supported on alumina, silica, titania and ceria, determined from TPRS over the temperature range 760 K to 1160 K. Table 1 shows the maximum selectivity for CO for each catalyst in the TPRS experiment. Table 1 Maximum CO selectivity and onset temperature for methane conversion for temperature programmed reaction of a 4:1 CH4/O2 mixture

Catalyst

Maximum Sco

Onset of methane conversion / K

0.85 0.39 0.57 0.60 0.53 0.33 0.29

647 657 594 712 645 627 589

Ni/SiO2 (pre-reduced) Ni(O)/SiO2 (non-reduced) Ni/TiO2 Ni(O)/TiO2 Ni/CeO2 Ni/CeO2 (partially pre-reduced) Ni/AI203

methane 5 ~

ol

CO

r ~ ,..,,

oxygen o...

01

200

400

600

800

1000

1200

Temperature I K Figt, re I. Temperature progr~L,ilt~led o• ratio = 4).

of methane over a 5% Ni/SiO2 catalyst (CH4/O2

396 1

0.9

NilSiOz

0.7 0.8

N~~O

~o6 +

=

"~ 0.5 ,-~0.4 ffl

O.3

0.2 0.1 0

650 too 750 800 850 900 950 1000 1050 1100 1150 1200 Temperature I K

Figure 2. CO selectivity over supported nickel catalysts for temperature programmed oxidation of methane (CH4/O2 ratio = 4). The TP spectra can be used to determine the temperature at which methane activation starts to occur. This is manifested by a decrease in the methane signal but more particularly by the onset of evolution of gaseous products, namely CO and/or CO2, and H2 and/or H20. The temperatures at which methane activation commences for each catalyst are also shown in Table 1. There are significant differences between the catalysts. Pre-reduced Ni/AI203 and Ni/TiO2 were found to activate methane at the lowest temperatures, 589 K and 59.4 K, respectively, whereas the non-reduced Ni/TiO2 catalyst showed no activity below 712 K. In all cases the first carbon oxide detected was CO2. The effect of the initial oxidation state of the nickel was investigated by carrying TP partial oxidation of methane directly following calcination of the catalyst without carrying out the reduction treatment. In general the pre-reduced catalysts gave rise to a higher transient CO selectivity than the non-reduced catalysts, and the sharp rise in CO selectivity is not observed until higher temperatures for the non-reduced catalysts. For example, for the Ni/SiO2 catalyst the CO selectivity increases to a maximum of 86% at 1160 K on the pre-reduced catalyst compared to 39% on the non-reduced catalyst, with the sharp increase in CO selectivity occurring at -1030 K on the pre-reduced catalyst but not until 1140 K on the non-reduced catalyst. Catalytic Partial Oxidation of Methane The activity of the nickel catalysts and their selectivity towards syngas formation were studied at reaction temperatures of 973 K, 1073 K and 1173 K. Table 2 shows the steady state methane conversion and CO selectivity for the non-reduced catalysts at these reaction temperatures for a 4:1 CH4/O2 ratio. The continuous sampling of the mass spectrometer enabled the exit gases to be monitored continuously, enabling the initial and non-steady state methane conversions and product selectivities to be determined throughout the reaction. This behaviour will be described in detail elsewhere I! 7 I. 3.2

397 Table 2 CO selectivity and methane conversion for a 4-1 CH4/O2 mixture over non-pre-reduced supported nickel catalysts at different reaction temperatures. Catalyst

Methane conversion / %

Ni(O)/AI203 Ni(O)/SiO2 Ni(O)/ZrO2 Ni(O)/CeO2 Ni(O)/TiO2

CO selectivity

973 K

1073 K

1173 K

973 K

1073 K

1173 K

15 14 92 4 0

95 94 97 19 15

99 30 99 96 38

0.00 0.00 0.72 0.10 0.00

0.77 0.78 0.85 0.00 0.00

0.91 0.49 0.92 0.90 0.45

After each catalytic experiment a temperature programnled oxidation (TPO) experiment was carried out to study the extent of carbon deposition during catalytic methane oxidation, and the nature of the carbon deposited. Very significant differences in the extent of carbon deposition were observed between the different catalysts. For all reaction temperatures studied the ceria and titania supported nickel catalysts showed essentially no carbon deposition, whereas substantial carbon deposition occurred on the Ni/AI203 catalyst during methane oxidation at 1073 K and 1173 K. Figure 3 shows the removal of carbon species by oxygen from the surface of the non pre-reduced 10% Ni(O)/AI203 catalyst, after passing 4:1 and 6:1 CI-t4/O2 gas mixtures over the catalyst for 6 hours at 1173 K. Interestingly more carbon is deposited for the mixture with the lower CH4/O2 ratio. For the 4:1 mixture carbon is removed in a single desorption peak with the maximum rate of removal occurring at 1200 K; a small amount of carbon is still being removed at the highest reaction temperature of 1250 K. In contrast, at a 6:1 CH4/O2 ratio four desorption maxima are observed, at 900 K, 1050 K, 1145 K and 1230 K, with the lowest temperature peak being the largest.

H4/O2 = 4.0

=m

coN

CH4102 = 6.0

m

c.o~

200

400

!

600

!

800

1000

1200

I

Isothermal from 1223K

Temperature / K Figure 3. TPO of non prc-redttced Ni(O)/AI203 following exposure to 41 and 6:1 CH4102 mixtures at 1173 K for 6 hours.

398 3.3

Pulse reaction studies The initial reactivity of CH4/O2 mixtures over reduced and non-reduced nickel catalysts was studied using a pulsed technique. For each catalyst sample the variation in the methane conversion, product selectivity and surface carbon yield, with the number of C H j O 2 pulses was determined. Figure 4 shows how the methane conversion, CO and CO2 selectivity, and surface carbon yield vary with the number of CH4/O2 pulses for a 4:1 CH4/O2 mixture over a pre-reduced Ni/AI203 catalyst at 1173 K. 0.9

100 ,,e

"

=.,~...~'.4.-=

0.8-

~'

"~-. o

.

9

-L

="

".,

9

9 q~D" "

eI

._> 0.so _~ o.4(D

I

-0

.-~

.0.

,

L 98

e

" ' . CH4

- 96

. O

r,r~ ~%'J

q'~,

-

92

-

90

-

88

I

9

0.3-

6 9

o ~

9CO=

~

'

"=A -

0.1 & 0

~-

9

- 94

0.6-

0.2

9 t~.

",

0.7 ~

-Ik

I-

i

C2 -

:

0

9

2

9 t

9 9 [~

9

t

9

i

9 ....

9 "

~-

--e

"

9 --

&"

'-I.

"~-

-A~*

9

4

--

~

.

-, 'bid

86 .

.. "I1"

84

cO

,=..,1

L_

0 > to o o .g:

9 E

82

9 "dk- .

I

I

I

l

I

I

I

6

8

10

12

14

16

18

Pulse

80 20

number

Figure 4. Variation of methane conversion, and CO, CO2, C2 and surface carbon selectivity, with the number of CH4/O2 pulses over pre-reduced Ni/AI203 at 1173 K. 4. DISCUSSION

Two distinct reaction schemes for the partial oxidation of methane to syngas have been proposed. The first, favoured by many workers [1,13,18,19], proposes a two stage process, involving complete oxidation of methane to CO2 and H20, followed by steam and dry reforming of the remaining methane by H20 and CO2 produced in the first stage of the reaction. The second mechanism, proposed by Schmidt and co-workers [7,8,20], involves pyrolysis of methane, followed by oxidation of surface carbon to CO. CO is thus formed as a primary product from methane oxidation without the involvement of CO2. CO2 can then be formed by reaction of CO with adsorbed oxygen atoms, in all our temperature programmed measurements CO2 is the first carbon oxide detected. This could be taken to indicate that CO2 is the primary product, and that CO is formed as the secondary product, by dry and steam reforming of the remaining methane. However, the more likely explanation is that at these low temperatures the reactivity of methane is low and the concentration of surface oxygen atoms relative to that of surface carbon, and hence CO, formed from methane pyrolysis, is high, and at these temperatures further oxidation of adsorbed CO to CO2 is favoured over CO desorption. As the reaction tenlperature increases so does the relative concentration of surface carbon and CO desorption becomes ,norc favourable, until above a certain temperature a sudden increase in CO selectivity is observed, t:or the Ni/SiO2 and Ni/TiO2 catalysts this is particularly dramatic, whereas the increase is tllttc'h leSS tnarked witll the Ni/AI203 catalyst, whilst the CO selectivity lot the Ni/CeO2 sample actually goes through a maximunl at -- 1()8() K, after increasing sharply

399 over the temperature range 920 K to 1080 K. It should be noted, however, that such data result from a dynamic experiment in which the surface of the catalyst is altering over the experiment, and hence with reaction temperature, for example, as a result of carbon deposition. The temperature programmed experiments also show there is a significant difference in the onset temperature for methane conversion over the catalysts, ranging from 589 K for the prereduced Ni/AI203 catalyst to 712 K for the non-reduced NiO/TiO2 sample. The data for the non-reduced catalysts suggest that either unreduced nickel oxide is active towards syngas formation, or that the CH4/O2 mixture reduces the NiO during the temperature ramp. Lunsford and co-workers have proposed that unreduced NiO/AI203 is active for the complete combustion of methane, with the subsequent reforming reaction occurring on reduced Ni/AI203 [13]. The temperature programmed data support the CH4/O2 mixture reducing NiO to nickel, with metallic nickel being the active species for methane partial oxidation, in agreement with the proposals of Au [3-5]. The catalysis data (Table 2) show there are clear differences in both the activities and CO selectivities of the catalysts. In general both methane conversion and CO selectivity increase with increasing reaction temperature, consistent with the findings of Schmidt and co-workers for Pt/A1203 and Rh/A1203 catalysts [7,20]. The titania and ceria supported catalysts are the least active and selective, particularly at low reaction temperatures where very low conversion and zero CO selectivity are observed. These catalysts also showed essentially no carbon deposition, even at the highest reaction temperature. This may be related to the reducible nature of the supports, and is the subject of further study. The alumina and zirconia supported catalysts are the most active and show high CO selectivity above 1073 K; TPO also reveals significant carbon deposition particularly for the Ni/AI203 catalyst. The Ni/ZrO2 catalyst is remarkably active even at the lowest reaction temperature studied, 973 K. TPO shows that significant carbon deposition occurs on some of the catalysts during catalytic methane oxidation, and that this is strongly bound only being removed from the catalyst by oxygen at temperatures as high as 1200 K. At lower levels of carbon deposition, carbon is removed in several lower temperature processes, suggesting that one form of carbon can be converted into a more strongly bound form with increased carbon deposition. The pulsed reaction method can provide insight into the reaction mechanism of methane partial oxidation and the nature of the active site at specific reaction temperatures lessening the problems of thermal gradients in the catalyst bed, and overcoming some of the disadvantages of temperature programmed experiments. For both non-reduced NiO/AI203 and pre-reduced Ni/AI203 catalysts the highest CO2 selectivity is observed at the start of the pulsing experiment, with high initial selectivity towards CO2 for the non-reduced catalyst. The data suggests that very rapid reduction of the NiO to metallic nickel is occurring with metallic nickel being the active species for methane partial oxidation, whilst NiO shows transient selectivity towards complete combustion of methane. The transient higher level of CO2 formation on the pre-reduced catalyst can be rationalised by either of the two mechanisms proposed for partial oxidation of methane, rather than simply in terms of CO2 being the primary product and CO the secondary product. The pulsed studies also clearly show a transient carbon deposition process, which coincides with the drop in CO2 selectivity to very low levels. 5. S U M M A R Y In summary, our results indicate that metallic nickel is the active species for partial oxidation of methane, with methane conversion and CO selectivity being favoured at higher reaction temperatures. The nature of the catalyst support has a considerable influence on the activity and CO selectivity of the nickel catalyst, with Ni/CeO2 and Ni/TiO2 catalysts being significantly less active for methane conversion and exhibiting much lower CO selectivity, as well as giving less carbon deposition than Ni/AI203 and Ni/ZrO2 catalysts. Temperature programmed measurements show the onset of methane conversion varies appreciably between the catalysts, ranging from 589 K on Ni/AI203 to 712 K on unreduced NiO/TiO2. Temperature programmed

400 oxidation indicates that several types of carbon species are formed on the catalysts during methane oxidation, differing in their strength of binding. As the anaount of carbon deposited increases conversion to more strongly bound species occurs. 6. REFERENCES

1

10 11 12 13 14 15 16 17 18 19 20

A.T. Ashcroft, A.K. Cheetham, J.S. Foord, M.L.H. Green, C.P. Grey, A.J. Murrell and P.D.F. Vernon, Nature, 344 (1990) 319. A.T. Ashcroft, A.K. Cheetham, M.L.H. Green and P.D.F. Vernon, Nature, 352 (1991) 225. C.T. Au and H.Y. Wang, J. Catal., 167 (1997) 337. C.T. Au, H. He, S.Y. Lai and C.F. Ng, J. Catal., 163 (1996) 399. C.T. Au, H.Y. Wang and H.L. Wan, J. Catal., 158 (1996) 343. O.V. Buyevskaya, D. Wolf and M. Baerns, Catal. Letts., 29 (1994) 249. D.A. Hickman and L.D. Schmidt, Science, 259 (1993) 311. P.M. Torniainen, X. Chu and L.D. Schmidt, J. Catal., 146 (1994) 1. M.V. Twigg (ed.), Catalysis Handbook, Manson Publishing, London, 1996. S.C. Tsang, J.B. Claridge and M.L.H. Green, Catal. Today, 23 (1995) 4. J.R. Rostrup-Nielsen and L.J. Christiansen, Appl. Catal. A:, 126 (1995) 381. P.D.F. Vernon, M.L.H. Green, A.K. Cheetham and A.T. Ashcroft, Catal. Today, 13 (1992) 417. D. Dissanayake, M.P. Rosynek, K.C.C. Kharas and J.H. Lunsford, J. Catal., 132 (1991) 117. V.R. Choudhary, A.M. Rajput and V.H. Rane, J. Phys. Chem., 96 (1992) 8686. D. Dissanayake, M.P. Rosynek and J.H. Lunsford, J. Phys. Chem., 97 (1993) 3644 A.M. Diskin, R.H. Cunningham and R.M. Ornaerod, in preparation. A.M. Diskin, R.H. Cunningham and R.M. Ormerod, in preparation. Y.F. Chang and H. Heinemann, Catal. Letts., 21 (1993) 215. Y. Matsumara and J.B. Moffat, Catal. Letts., 24 (1994) 59. D.A. Hickman and L.D. Schmiidt, J. Catal., 138 (1992) 267.

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Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

403

N o v e l s e n s o r for s t u d y i n g t h e t r a n s i e n t b e h a v i o u r o f an iron a n t i m o n a t e p a r t i a l oxidation catalyst D. Barth a, M. Sahibzada b and I.S. Metcalfe b ~'Department of Chemical Engineering, Imperial College, London SW7 2BY, United Kingdom bDepartment of Chemical Engineering, University of Edinburgh, Edinburgh EH9 3JL, United Kingdom Abstract Oxygen-ion conducting solid electrolyte cells were prepared with Pt reference electrodes and FezO3-Sb20 4 working electrodes. The working electrode served as both electrode and catalyst during the partial oxidation of propene. The electrochemical cell allowed thermodynamic catalyst surface oxygen activities to be followed and acted as a sensor of catalyst state. It is clear that the catalyst was much more selective for acrolein formation at lower oxygen activities as indicated by the sensor signal. There is evidence to suggest that the oxygen activity measured by the sensor is as a result of the presence of adsorbed oxygen and not lattice oxygen. Based on the results it appears that potentiometric sensors could provide important additional information for catalytic partial oxidation process monitoring and control.

1. INTRODUCTION Oxygen partial pressure can be determined electrochemically by measuring the potential difference between the two electrodes of an electrochemical cell consisting of a solid electrolyte membrane onto which two porous electrodes have been deposited. The charge overall transfer reaction on both sides of the cell is, 02 + 4e- r

202-

(1)

where 0 2- represents lattice oxygen in the solid electrolyte and e- represents an electron associated with the electrode. Consequently, the e.m.f, of the cell relates the oxygen partial pressures on both sides of the membrane through the Nernst equation, RT P"o2 E = ~ln~ 4F P'02

(2)

This equation relies on the assumption that oxygen at the electrode surfaces is in equilibrium with gas phase oxygen. In the case where one of the electrodes is exposed to a reacting gas mixture equilibrium between gas phase oxygen and oxygen at the surface need no longer exist and the e.m.f, of the cell becomes a reflection of the thermodynamic activity of oxygen on the electrode in question (providing a mixed potential does not occur),

E

RT In a ~

4F

P(;2

(3)

404 where a o is the thermodynamic activity of oxygen on the catalyst. Potentiometric sensors have been used to investigate the properties of the prevailing gas phase [ 1]. In contrast, the object of this work was to investigate potentiometric sensors for studying the catalyst itself. 2. EXPERIMENTAL 2.1 The Apparatus The solid-state electrochemical cell used in this study, was the following: Air, Pt [ ZrO2, 8% Y203 [ Fe203-Sb204, C3H6, 02, C02, C3H40, N2, He The working electrode was exposed to reaction conditions and served simultaneously as the catalyst for the reaction under study. The reference platinum electrode was exposed to air. The electrode potential difference, i.e. the e.m.f, of the cell, was monitored by means of a DD 10M potentiostat/galvanostat (Sycopel Scientific). The cell itself was in the form of an yttriastabilised zirconia (YSZ) thimble, with an internal painted Pt electrode and an external ring electrode made of the FezO3-Sb204. The YSZ thimble, with the inner Pt electrode in place, was provided by Allied Signal, who manufacture such cells for use as automobile exhaust gas sensors. The whole assembly, thimble and disk, was then placed in the reactor housing, which was made of aluminium. Analysis of the inlet and exit gas streams was performed using a gas chromatograph.

2.2 Fe203-Sb204electrode preparation The Fe203-Sb204 catalyst powder was prepared by following the method of Allen et al. [2]. The ratio of iron to antimony employed in this work was Sb:Fe = 2.5 g of the oxide catalyst powder were mixed together with 8.5 mg of fish oil (dispersing agent) and 3 g of ethanol (solvent) to produce a slurry. Additionally 80 mg of PVB (polyvinylbutyral) were added as a binder with 38 mg of a 1:1 plasticiser mixture of polyethylene glycol 400 and dibutyl phthalate. The slurry was then homogenised in a ball mill for 58 hours. The oxide electrodes were prepared by first applying a fine gold paste layer over the YSZ electrolyte (this gold layer acted as the electrical contact). The gold layer was sintered at 700~ for two hours. The oxide slurry was then painted over the Au-YSZ and fired at 600~ for three hours. 3. RESULTS AND DISCUSSION The technique of solid electrolyte potentiometry (SEP) was employed to study the state of the oxide catalyst under reaction conditions (propene oxidation). The Nernstian behaviour of the oxide electrode was confirmed using oxygen-nitrogen gas mixtures of known compositions at 400 and 450~ The e.m.f, values obtained were within 5 mV of those predicted by the Nernst equation. There was no strong hysteresis in the observed steady state reaction rates for the production of total oxidation products or acrolein when oxygen inlet concentrations or propene inlet concentrations were cycled. Experiments were performed in which the sensor was exposed to changes in gas phase concentrations. In Figures 1 and 2 the catalyst is exposed to propene in helium for 24 hours prior to introduction of oxygen. The e.m.f, or open-circuit potential (OCP) is then monitored

405 with time. Initially the OCP is very negative (-~ -400 mV) due to the prereduction of the catalyst. On introduction of the oxygen the surface is quickly oxidised leading to an OCP of around -150 mV (i.e. more oxidised than under steady operating conditions). However, this oxygen is on the surface of the catalyst and is therefore not expected to be very selective for partial oxidation [2] and indeed a poor acrolein selectivity is observed. With time this surface oxygen begins to migrate into the oxide catalyst (there is a driving force for this because of the prereduction of the oxide) and as this migration occurs the surface of the catalyst appears to become more reduced and the OCP drops to -200 mV to -250 mV. With this formation of reactive lattice oxygen we would expect to see an increase in the selectivity for acrolein (as lattice oxygen is thought to be more selective for partial oxidation [2]) and indeed this appears to be the case. In Figure 2 we can clearly see that when the catalyst OCP goes through its minimum there is a maximum in the yield of acrolein. At the same time as oxygen migrates from the surface to enter the lattice the yield of carbon dioxide is decreased. At longer times, as the gradients in the oxygen activity of the catalyst disappear, the surface returns to an oxidation state corresponding to around -200 mV and consequently the acrolein selectivity drops.

0

12

,o

0 -50 I -100

-100

OCP

-150 ~'-200

s

8 6 -o

ILl -250 -300

03H4(

-350 ~, ~ ; ~ - " : ' - - " : ~ - -400 ~,,,,,,,.,,. . . . -450 0

~,,-

. . . . .

t 500

i 1000

time [min]

4

10

OCP

-150 -200

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

6

~ ' -250

C3H4(

-300

-400

0

Figure 1: Transient behaviour at 400~ and 5% 02 over prereduced electrode surface

8

"~-~---=-~--~-~ -

-350

e-O~' i 1500

I5% C3H6.

.~ ..~_,,,. . . . . . . . . . . . . . . . . . . . . . -450 t 0 500

i

_

t 1000

time [min]

,--,-

=

C02

4

E "O

~2 >-

2

........

I 1500

Figure 2: Transient behaviour at 450~ and 5% 02 over prereduced electrode surface

If the catalyst is preoxidised in oxygen in helium, the starting OCP is in the region o f 50 to 0 mV (see Figures 3 and 4). On introduction of the reacting gases the surface is rapidly reduced to around -400 mV. Although one might expect to see high selectivities on this highly reduced surface any such effect is impossible to observe because of the rapidity of the processes involved. The large amount of available surface oxygen results in very high initial rates of carbon dioxide production and although carbon dioxide production rapidly drops and acrolein production rapidly rises it is difficult to perform an accurate evaluation of selectivity because of the short timescales. As the surface then approaches its steady state level of oxidation (i.e. becomes more oxidised) the yield of acrolein decreases. However, we can see in Figure 4 that during the first 500 minutes of the transient the catalyst remains in a reduced state realative to its final steady state and during this period the yield and selectivity for acrolein formation is enhanced. The catalyst's behaviour can be rationalised by considering there to be at least two different forms of surface oxygen available. It appears that the e.m.f, of the sensor is most sensitive to the more reactive adsorbed oxygen and not lattice oxygen. It may be expected that

406 the sensor would be most sensitive to the most reactive forms of oxygen on the catalyst surface as these forms of oxygen are probably also the most reactive forms for the electrochemical reaction which determines the e.m.f. (this is an important principle which could lead to the wider application of this type of sensor). This reactive adsorbed oxygen is selective for the formation of carbon dioxide and therefore we would expect the catalyst to favour acrolein formation when activities of this adsorbed oxygen are low. This has clearly been demonstrated through the use of the sensor under transient conditions.

0 r

12

-50 15 % 03H6 -100

10

-150

8

uJ -250

OCP

03H40

-400-350i'-~,~,~~ ~ ~ I 500

I 1000 time [min]

C02

I 1500

4

15~

-100

OCP

-150

6..o

~"~ " ~

-300

-450 , 0

0 -50

.~>-

2 0

Figure 3: Transient behaviour at 400~ and 5% O2 over preoxidised electrode surface

~'

.E

-2oo

03H40

uJ -250 -300

,-..,

eL 4

>"

-350 -400 -450

0

I 500

t 1000 time [mini

C02

I 1500

Figure 4: Transient behaviour at 450~ and 5% 02 over preoxidised electrode surface

4. CONCLUSIONS Oxygen-ion conducting solid electrolyte cells were used to determine catalyst oxygen activities during the partial oxidation of propene over Fe203-Sb204. It appears that the e.m.f. of the sensor is most sensitive to the more reactive adsorbed oxygen and not lattice oxygen. This reactive adsorbed oxygen is selective for the formation of carbon dioxide and therefore the sensor signal has been found to correlate with catalyst selectivity under transient conditions. Based on these results it is clear that potentiometric sensors can provide important information for catalytic partial oxidation process monitoring and control. The possibility that such sensors are most sensitive to very reactive forms of surface oxygen could lead to a wide range of catalytic applications. 5. ACKNOWLEDGEMENTS DB wishes to acknowledge the financial support of EPSRC. We wish to thank Allied Signal for supplying the YSZ thimbles and Dr E. van Steen for preparing the catalyst. ,

[1] [2] [3] [4]

REFERENCES W.C. Maskell and B.C.H. Steele, J. Appl. Electrochem. 16, 475 (1986). M. Allen, R. Betteley, M. Bowker, G.J. Hutchings, Catal. Today 9 (1991) 97-104. P. Petrolekas and I.S. Metcalfe, J. Catal. 152 (1995) 147-163. P. Petrolekas and I.S. Metcalfe, J. Catal. 157 (1995) 545-549.

Reaction Kinetics and the Developmentof Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science Ltd. All rights reserved.

407

C h e m i c a l e q u i l i b r i a i n d i r e c t s y n t h e s i s of d i m e t h y l e t h e r

M. Grzesika, b and J. Skrzypek a aInstitute of Chemical Engineering., Polish Academy of Sciences, Gliwice, Poland bFaculty of Food Technology, Academy of Agriculture, Krak6w, Poland

Abstract

The effect of the initial feed composition, t e m p e r a t u r e and pressure on the equilibrium conversions were examined for direct synthesis of dimethyl ether from CO2 and H2 or from CO and H2. The three-reaction system was analysed for both considered cases. From a thermodynamic point of view, direct synthesis of dimethyl ether from CO2 and H2 or from CO and H2 seems to be quite promising for industrial applications.

1. I N T R O D U C T I O N Dimethyl ether (DME) is an important chemical and a chemical intermediate. It is widely applied as a fuel for diesel engines or a fuel additive and also as a refrigerant instead of the freons. DME can be synthesised directly either from CO2 and H2 [1] or from CO and H2 [2] if hybrid catalysts are used. The following main reactions may occur: -

synthesis of DME from CO2 CO2 +3H2 <==> 2CH3OH <==> CO2 + H2 <==>

- synthesis of DME from CO CO + 2H2 <==> 2CH3OH <==> CO + H20 <==>

and H2 CH3OH + H20, CH3OCH3 + H20, CO + H20,

and H2 CH3OH, CH3OCH3 + H20, CO2 + H2.

(1) (2) (3) (4) (5) (6)

All the reactions are reversible. The reactions (1) and (4) are exothermic and proceed with volume contraction. The purpose of present study is to determine the effect of process p a r a m e t e r s on equilibrium conversions in the reaction systems (1-3) and (4-5).

408 2. M A I N R E S U L T S The effect of t e m p e r a t u r e (400-700 K), p r e s s u r e (1-10 MPa) and initial molar fraction of C O 2 / C O / (0.2-0.5) on the e q u i l i b r i u m c o m p o s i t i o n s of the r e a c t a n t s has been evaluated. The e q u i l i b r i u m c o n v e r s i o n s are defined as follows:

~1" = An co2/co/(1 or 4)/Fon, ~2" = A n D M E / F o n ,

~3" = An co2/co/(3 or 6)/Fon, w h e r e Fon is a total initial n u m b e r of moles. S e l e c t e d r e s u l t s of n u m e r i c a l c o m p u t a t i o n s are i l l u s t r a t e d in Fig. 1- 5 (CO9 --> D M E ) and in Fig. 6-10 (CO --> DME). The a l g o r i t h m u s e d w a s s i m i l a r to t h a t p r e s e n t e d in [3]. 0.15 -

Chemical equilibrium Synthesis of dimethyl ether XOCO2=0.2, XoH2=0.8, p=l MPa

!"O

9~ 0.10

" ,

+

400

@

reaction (1)

~

reaction (2)

+L 13

.

+

,

+

***n

[]

I"!

§

I OO@O

000

O

§

+ §

eve,

or,

v

,.

. . . . . . . . . .

500 550 600 Temperature, K

aoa

aaaO

650

700

400

450

+

§

9

an aa

00~

reaction (31

+

§

[]

§247

a aa

450

,,a

[]

§ + § §

n n a a

[] [] a a

.... tion(S),~[] []

+

..

0.00

reaction (1) reaction (2)

++§247

+

.

4"

C h e m i c a l reactions:

§

.;~ 0.05

I I Chemical reactions-

Chemical equilibrium Synthesis o f d i m e t h y l ether X o C O = 0 . 2 , X o H 2 - 0 . 8 , p=5 M P a

+

4.+

a

500 550 Temperature,

600 K

650

700

Fig.1. Direct synthesis of DME from C02 and H~ Fig.6.Direct synthesis of DME from CO and He. 0.25-~

C i t e m i c a l e~luilibrium S y n t h e s i s o f d i m e t h y l ether XOCO2=0.2, XoH2--0.8, p=5 MPa i J

| ~

0.20-

.~

0

I~ 0.10

r. 9

9 9+

Chemical reactions:

'- 0.15- . o;~ =

_ =

Chemical equilibrium I Synthesis of dimethyl ether [ XoCO=0.3, XoH2=0.7, p=l MPa I

§ 9§

-

!§

§

§

~ .

i ***

0.05 0.00~

400

,***

. . . . . . . .

450

§

§

+

reaction (1)

~

reaction (2)

~ I"!

§

**** ~ nan

-~-

reaction (3)

§

§

a

~

*.**

a

a

9 +'§

***

a OQO

.

.

.

.

.

.

.

~8

=~

reaction (1)

§

~,

reaction (2)

.

[~l

. . . . fion ( 3 ) ~ _

+

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

.

a

"2 §

B a

+

500 550 600 Temperature, K

Chelmical reactions:! § §

al

e+/:**

650

700

400

450

500 550 Temperature,

§

600 K

650

700

Fig.2. Direct synthesis ofDME from CO2 and H~ Fig.7.Direct synthesis of DME from CO and H~

409 0.25

I

Chemical equilibrium Synthesis of dimethyl ether XOCO2=0.2, XoH2=0.8, p=9 MPa -

I-

-~ 0 . 2 0 + + + +

@

0.15-

+

4"

+

N O.lO-

+

§

§

a

*§

§ +

500 550 600 Temperature, K

oooo

a nan

~***

650

700

400

§

1

reaction (2) reaction (3)

+ v.v

+

~

anna

450

~ ~ +

or§ anna

§ ++ o§

a

noon

450

+ + + +

reaction (3) n a a I [] ~ D

§ §

~ 2 4 7 2o4 7 . . . . . . .

.-~

+

reaction (1)

a §

*vo~

o.oo- . 4oo

I

reaction (2)

§

.r. ~9 o.o5-

+ + +

~Chemicai reactions: +

1

Chemical equilibrium C h e m i c a l reactions: Synthesis of dimethyl ether X o C O = 0 . 3 , X o H 2 = 0 . 7 , p = 5 M P a -. "[" reaction (1)

~fifia

500 550 600 Temperature, K

§

4.

§ §

650

700

Fig.3. I)irect synthesis of DME from CO2and I-~ Fig.8.Direct synthesis of DME from CO and H2. 0.25

I

Chemical equilibrium

t_ tNI

Synthesis of dimethyl

.~ 0.20 :-

ether

I

/

XOCO2=0.3, XoH2=0.7, p=5 MPa I

Chemical equilibrium I Chemical reactions: Synthesis of dimethyl ether reaction (1) XoCO=0.3, XoH2--0.7, p=9 MPa 4"

C h e m i c a l reactions:

0.15

+§

:

§ +

=

~

0.10 .

+ * ~176176176

:-- 0.05

~

o" :~

§

reaction (1)

~

reaction (2)

a

reaction ( 3 ) n o +

o§

*

"

[a~ 0.00

§

400

a

450

+

o~ §

"0§

. . . . . . .

++§

@

an

=

o

v

idl

0.25

ovov

ann

~***

650

700

400

-~ 0 . 2 0

vo**

n

nna

450

09" 4 , §

XOCO2=0.4, XoH2--0.6, p?5 MPa

§ + + §

+

N 0.10 .,~9 o.o5 0.00

+ " +

"00.

+

4=

reaction (1)

~

reaction (2)

['1

reaction (3} a n [] n [] o G

+

*§ ....

4oo

noon

450

+

r * o

reaction (3)

+

+

§

+

§ r

§

+

,

~Baa

n

500 550 600 Temperature, K

§ § § §

§

650

4, !

700

[]0 0

reaction (2) §

a unnn

nnnu

reaction (1)

§ §

§ , [I §

reaction (3).

§

naaa

§

a m

I

++

500 550 600 Temperature, K

,

reaction (2)

I "Chemical reactions:"

C h e m i c a l reactions:"

0.15

"0§

o o

synthesis of DME from CO and I~.

Chemical equilibrium Synthesis of dimethyl ether

.

+§

nO a

Fig.4. Direct synthesis ofDME from CO2and FI~ Fig.9.~

II~

++++

+

o

500 550 600 Temperature, K

++§

;50

700

Fig.5. Direct synthesis ofDME from COzandH~

Chemical equilibrium I Synthesis of dimethyl ether XoCO=0.5, XoH2=0.5, p=5 MPa . . . . . . . . I . . . . . . . . . 1. . . . . . . . . I . . . . . . . . . 400 450 500 550 600 Temperature, K

I a

650

Fig.10.~syn~ofDMEfromCOandH~

§

§247 maim

700

410 2.1. S y n t h e s i s of DME f r o m CO2 a n d H2 Both temperature and pressure have a considerable effect on the equilibrium conversion degrees of (1-3). Pressure influences strongly the equilibrium conversion degrees of (1) and (2) in low temperatures as 400-550 K and has no effect in higher ones. On the contrary, pressure effects on the equilibrium conversion degree of (3) only in temperatures over 600 K. An increase in temperature decreases the equilibrium conversion degrees of (1) and (2) and increases that of (3). The effect of the initial mole fraction of CO2 on the equilibria in the system of (1-3) is also highly significant. 2.2. S y n t h e s i s of DME f r o m CO a n d H2 Pressure has rather a moderate effect on the equilibrium conversion degrees of (4-6). The equilibrium conversion degree of (4) increases distinctly with increasing pressure, while those of (5) and (6) only slightly depend on pressure. The curves of the equilibrium conversion degrees of (4) and (5) decrease with temperature. The equilibrium conversion degree of (6) increases initially with temperature, passes through a maximum, and then decreases. The position of maximum point depends on both pressure and the initial mole fraction of CO.

It is also seen that in the whole studied parameters range the equilibrium conversion ~2" of (1-3), which is proportional to the equilibrium yield of DME, is higher than respective one of (4-6). 3. R E F E R E N C E S

1 Dubois J.C., Sayama K., Arakawa H., Chemistry Letters, (1992) 1115. 2 Li J.-L, Zhang X.-G, Inui T., Appl. Catal. A: General, 147 (1996) 23. 3 Skrzypek J., Lachowska M., Serafin D., Chem. Eng. Sci., 45 (1990) 89.

Reaction Kinetics and the Developmentof Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999Elsevier Science Ltd. All rights reserved.

Thermodynamics alcohols

and

kinetics

of t h e

411

synthesis

of h i g h e r

aliphatic

M. Grzesik ab, M. Kulawska ", J. Skrzypek . and M. Witczak b aInstimte of Chemical Engineering, Polish Academy of Sciences, Gliwice, Poland bFaculty of Food Technology, Academy of Agriculture, Krak6w, Poland

Abstract The thermodynamics of the synthesis of higher aliphatic alcohols is studied in detail. Kinetic model of the synthesis of higher aliphatic alcohols is presented. Experiments were carried out in a high-pressure continuous gradientless stirred tank reactor. Unexpectedly, the reaction rate is independent of partial pressure of carbon monoxide.

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

The search for a clean combustion fuel is the most important incentive to improve the synthesis of higher aliphatic alcohols. The mixture of methanol and higher alcohols appears to be a very valuable additive to gasoline as an antiknock agent. They can be a real alternative for MTBE since they are entirely based on natural gas. It is a clean fuel without aromatics, olefins and sulphur. The last review papers on this subject were [1,2]. The stoichiometry of higher alcohol synthesis from syngas is based on the following reaction scheme: m CO + 2mH2

<==>

CmH2m+IOH+ (m-1)H20,

m=1,2...

(1)

All the reactions are reversible, exothermic and proceed with volume contraction. The water-gas shift reaction is always present and assumed as attaining to the state of chemical equilibrium at the synthesis conditions: CO+H20

<==>

CO2+H2.

(2)

In only a few studies attempts have been made at modelling the kinetics of the overall rate of the synthesis of higher aliphatic alcohols [3,4].

412 The thermodynamic data concerning the process are scarce and limited in scope. The results cited often refer to the individual reactions only, thus giving an unrealistic general description. The comprehensive work has been published by Mawson et al. [5] . The thermodynamic background of the synthesis has been shown by Xiaoding et al. [6]. Tronconi et al. [7,8] have presented a thermodynamic analysis concerning their experimental results.

2. CHEMICAL EQUILIBRIUM STUDY A thermodynamic model describing the synthesis of C1-C4 aliphatic alcohols was developed for the system of chemical reactions (1-2). The numerical solution of the nonlinear algebraic equations allowed the estimation of the equilibrium conversion degree of carbon monoxide. Selected results of numerical computations are presented in Fig. 1-4 (m=2,3,4): 1.0

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

~'~

i

0.6

Chemical reactiom :|

O

O.4-

.9~ "g~

..........

co ~P~O" l

!

I

/

i

550

n ~

,.,.

[] +

u ~

.....

[] +

! §

.

~

! ......... 600

o f higher

§

alcohoIs

X01t2=0.7,

]

p - 3 1 V I P a I.........

'" . .". . . . i . . . . " . . . . . . " . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

~][ r e a c t i o n (2) | . . .9. . . ,. . . . .,. . . . .9. . . .+. . . . *. . . . .*. . . . *. . *

0.2

0.0

Synthesis

...................................... i.......... X o C O = 0 . 3 ,

0.8

§

[

!

............... ~

i i i :.............................. *_ a m v .-.--.O----o----a----u---u----m u , . . . :: i §

+

§

I .........

+

650

Temperature,

§

i

+

+

+

I .........

700

*

§

i *

t

750

K

Fig. 1. Chemical equilibria in synthesis of higher alcohols The effect of the initial mole fraction of CO on the equilibria in the system of (1,2) is highly significant. Pressure has rather a moderate effect on the equilibrium conversion degrees of CO to ethanol and propanol (m=2 and m=3). In the whole range studied the equilibrium conversion degrees increase slightly with increasing pressure and decreasing temperature. Both temperature and pressure have a considerable effect on the equilibrium conversion degree of CO to butanol (m=4). Pressure influences strongly the equilibrium conversion degree in high temperatures as 600-750 K and has rather moderate effect in lower ones. An increase in temperature decreases the equilibrium conversion degree of CO to butanol. The equilibrium conversion degree of (2) increases initially with temperature, passes through a maximum, and then decreases. The position of maximum point depends on both pressure and the initial mole fraction of CO.

413

1.0

. "o

i

,9

A

9

reactions:

Chemical

~"

Synthesis

of higher

...... i................... Xoco=O.~,xom.--o.7,

...... ~:~;

0.8

9

~* 9

p~

I i

Meal--- I i

9

........................................... : .....:.....,~......, ......: ...................... ~.

0.6

4-

~ 0.4

~,

CO - > EtOH

i

CO - > P r O H

i

CO

->

......~ ~ ~ ~ -

,,

v

0.0

V

i

[]

9

0

4.

. . . . . . . . .

550

......~ ~ ~

n

4.

:

!

......~i ~ ~

[]

.~

ii

.

0.2

"

........................................................ i....................................... i ::

BuOH

reaction (2) ~~

alcohols

.I.

n

n

a

4.

.I.

4.

I . . . . . . . . .

600

.....r ......_~ . ~ a ~m ~" - i ~ i O ~ i 0

i 4'

T i

i

a

n

"

4.

.I.

.I.

I . . . . . . . . .

.I.

650

i §

.I.

r

§

I . . . . . . . . .

700

Temperature, K

i 4.

.I.

I

750

Fig. 2. C h e m i c a l e q u i l i b r i a in s y n t h e s i s of h i g h e r alcohols 1 . 0 ......................................... [ ...................................................................................................................

"~

i Chemical equilibrium i Synthesis of higher alcohols ,------~......~......~......~.......i................... X o C O = 0 . 3 , X o H 2 = 0 . 7 , p - 7 M P a

0.8

O

~"

Chemicalreactiom: 0.6

4-

9

~ 0.4-

..z-

1

:* "

"

9

co->EtOH/....................i.............................. :.......!.....................~......~...... co o- / co

- > B u O n ] ........................................................... !......................................

:.....r~......~......~......~......~......~......~176176176176

0.0

........

I ......... 600

550

.....~ ~ C *

I ......... 650

Temperature,

I ......... 700

I

750

K

Fig. 3. C h e m i c a l e q u i l i b r i a in s y n t h e s i s of h i g h e r alcohols 1.2 / Synthesis "IXoco=0,4,

1.0 O

9w~ -

0.8

ofhigher

alcohols I/ MPali

_ "P

Xom=0,6,p = 3

_> CO

- ............................... i.............................. .............................. ii

I

I

0.6

.

~

.~

0.4

-

9-

0.2

-~--u--.---.

9

,

~

:;--: ..............;... ..........................

............................... i ............................... i ............................. ! ............................ ~ ....................... .,......~

....... i,..o

i 0.0

~

/ EtOH [

[]

[]"

"

[]

[]

[]

'

[] u [] [] [] [ ] , [] [] [] [] i. ......... i .............................. i .............................. i

i

i

,,,,,,,,,I,H,,,,,,I,,,,,,,,,I,,,,,~,,,I,,,,,,,,~

500

550

600 Temperature,

, I

650

700

750

K

Fig. 4. C h e m i c a l e q u i l i b r i a in s y n t h e s i s of h i g h e r alcohols

414 3. K I N E T I C S T U D Y

The catalyst containing mainly CuO and ZnO with Zr, Fe, Mo, Th and Cs oxide addition has been prepared in our laboratory. It is a kind of low temperature modified methanol synthesis catalyst. Three different methods of catalyst preparation were used; the best results were obtained for the method using citric acid. The catalyst exhibits a remarkable stability during one-year experiments and high selectivity towards alcohols. Experiments were carried out in a high-pressure continuous gradientless stirred t a n k reactor. This type of reactor allows a direct determination of the reaction rate. The range of experimental parameters used was: P = 4.0-10.0 MPa, T = 553 - 653 K, H2/CO ratio = 0.85 - 3.16, GHSV = 900 - 12 000 h -1. The experimental conditions allowed the process proceeds in the intrinsic kinetic area. At low conversion degrees of carbon monoxide attained, it was far from chemical equilibrium and the reverse reactions were negligible. The liquid product consisted of methanol, ethanol, propanol, C4-C7 aliphatic alcohols and water. Hydrocarbons were practically absent but traces of methane were detected. Unexpectedly, the reaction rate is independent of partial pressure of carbon monoxide. It can be seen t h a t the values of activation energy are typical for catalytic reactions and slightly differ from each other. Using the standard fitting procedures the rate of the reaction was simply described as follows: r = k o exp(-E / RT).pH ~ , (3) by the mean error not greater than 10%. The detailed values of the parameters are included in Table 1. Table 1 Kinetic p a r a m e t e r s in eq. (3) Reaction ko [mol/g/h/MPa n m=l 218.1 m=2 252.3 m=3 84.61

E [cal/mol] 18440 19530 19520

n 2 1.5 1.5

alcohol MeOH EtOH PrOH

4. R E F E R E N C E S

1 2 3 4 5 6 7 8

J2. Hindermann, G~I.Hu/rhmgs and/k Kiennemann, Cata[ R e v . ~ Eng., 35 (1993) 1. J.C. Slaa, J.G. von Ommen and J.I~H. Ross, Catal. Today, 15 (1992) 129. EAVI.Calverley and IZ~J.Smith, In& Eng. Chem. Res. 31 (1992) 792. E.Troncom, N.Ferlazzo, P. Forzatti andI. Pasquon, In& Eng. Chem. Res., 26 (1987) 21. S. Mawson, M. S. McCutchen, P. I~ Lim and G. W. Roberts, Energy&Fuels, 7 (1993) 257. ~ Xiaoding, E. B. M. ~ u r g andJ. J. F. Scholten, Cata[ Today, 2, 125, Elsevier, 1987. E. Troncom, P. Forzatti and I. Pasquon, J. Catal., 124 (1990) 376. E. Tmncom, L. Lietti, G. Groppi, P. Forzatti andI. Pasquon, J. Catal., 135 (1992) 99.

Reaction Kinetics and the Developmentof Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science Ltd. All rights reserved.

K i n e t i c s of e s t e r i f i c a t i o n of a c r y l i c a c i d w i t h C3 a n d a l c o h o l s i n t h e p r e s e n c e of s u l f u r i c a c i d a s a c a t a l y s t

415

C4

aliphatic

M. Grzesik ab, J. Skrzypek a and M Witczak b aInstitute of Chemical Engineering, Polish Academy of Sciences, Gliwice, Poland bFaculty of Food Technology, Academy of Agriculture, Krak6w, Poland

Abstract Acrylic acid and its esters are important organic products. The results of the kinetic studies on the esterification of acrylic acid with C3 and C4 aliphatic alcohols in the presence of sulfuric acid as a catalyst are presented in the range of temperatures 45 - 100~ It was found that the reaction of acrylic acid with C3 aliphatic alcohols proceeds as the forth order - second order with respect to the acrylic acid and second order with respect to alcohol (double square kinetics).

1. I N T R O D U C T I O N Acryhc esters are very important monomers that are widely applied in production of homo- and copolymers. The kinetic data of the process of esterification of acryhc acid with lower ahphatic alcohol are scarce and bruited in scope [1,2]. The synthesis of acrylic esters can be represented by the scheme: CH2= CHCOOH + ROH

<===>

CH2= CHCOOR + H 2 0 .

The aim of our study was to develop rigorous kinetic equations for the esterification of acryhc acid with C3 and C4 ahphatic alcohols in the presence of H2SO4 as catalyst. 2. E X P E R I M E N T A L

The equipment employed in this study allowed conducting the experiments without removing of water. The main component of the apparatus was a four-necked glass flask of 1 dm 3 capacity placed in a thermostat. This flask was equipped with a thermometer, an alcohol inlet, a high-speed mixer, a head for collecting samples and an azeotropic head with a cooler. The experiments were carried out at the range of temperatures of 45 - 100~ at various acryhc acid to alcohol molar ratios and at various concentrations of H2SO4. The catalyst was used in the form of alkyl sulhtric acid.

416 3. M A I N R E S U L T S

3.1. Esterification of acrylic acid with C3 aliphatic alcohols Unexpectedly, the reactions seem to be forth order - second order with respect to acid and second order with respect to alcohol (double square kinetics). The final kinetic equations are as follows: acrylic acid + n-propanol (A-acrylic acid, P-n-propanol, E-ester, W-water, k-sulfixric acid): r=k Ck(CA2Cp2- CE2Cw2/K), where K = 2 and k = 3.16 10 -~ exp(-16400+100/RT) [m12/(mo14min)], acrylic acid + isopropanol (I-isopropanol): r=k Ck(CA2CI 2 - CE2Cw2/K), where K = 2 and k = 1.10 10 .5 exp(-17100+200/RT)

[m12/(mo14min)].

3.2. Esterification of acrylic acid with C4 aliphatic alcohols In difference to the results obtained for esterification of acrylic acid with Ca aliphatic alcohols reactions appear to be second order with respect to acid and alcohol. The final kinetic equations are as follows: acrylic acid + n-butanol (A-acrylic acid, B-n-butanol, E-ester, W-water, k-sulfimc acid): r = k Ck(CACB- CECw/K), where K = 1.5 and k = 56.6 exp(-15300+100/RT) [m6/mol2min], acrylic acid + isobutanol (I-isobutanol): r = k Ck(CACI - CEcw/K), where K = 1.5 and k = 85.4 exp(-15600+100/RT) [m6/(mol2min]. Exactness of fit for the selected examples is illustrated in Fig.l-3 (Ca aliphatic alcohols) and in Fig.4-6 (C4 aliphatic alcohols). 250 -

Esterification (acid/alkocol 1:3)'[ 9Expe'rimentai data: ~ acrylic acid + propanol

sulfur acid 2 wt %

200

o

temp.

45oc

temp. 55oc

150

I:!

temp.

O

temp. 75oc temp. 85oc kinetic model

~ 100-i

\"

so0

65oc

-

60

,

r~ u

120

180 240 300 Time, min

i

360

420

480

Fig. 1. Comparison of experimental d a t a with those obtained from kinetic model.

417 300 .-. 250 - -

~_

r Experimental data:

Esterificaion (temp" 75~ acrylic acid + propanol sulfuric acid 2.0 wt %

acid/alcohol 1:2

200

['!

acid/alcohol1:3

O

acid/alcohol 1:4

E =

kinetic model

150

o~

100 50-

o

o

u

~

0 40

0

80

120 Time, min

160

200

240

Fig. 2. Comparison of experimental data with those obtained from kinetic model. 280-=_ rsterification (acid/alcohol' 1:3) ~ Experimental data: " i acrylic acid + propanol 240 = temperature 75oc + H2SO40.5 wt % ~ O H2SO41.0 wt % ~., 200 + ~+~2,.. Q H2SO4 2"0 wt %

E =

-

160

"0 <~ 120

O

H2SO4 3.0 wt %

,

. so44o t. o

40 0

40

80

120

Time, min

160

200

240

Fig. 3. Comparison of experimental data with those obtained from kinetic model. 140 " 105 ~~

, , . Estrification (acid/alcohol 1:5) 9Experimental data: acrylic acid + isobutanol ~. sulfuric acid 1 wt % 4"~ temp.temp" 8070~176

0 =

O 70

"<

~'<---5. o

0

§

§ ~

'

0

temp. 90 oc kinetic model

''

50

100

150 Time, min

200

250

300

Fig. 4. Comparison of experimental data with those obtained from kinetic model.

418 200

l ~ , Estrification (temp. 90oc)] acrylic acid + isobutanol I sulfuric acid 1 wt % ~

"

160 . ~ O I"

"~

120

Experimental data:

4"

acid/alcohol 1:3

.~

~

acid/alcohol 1:4

~ ~k

~i

acid/alcohol 1:5

O

acid/alcohol 1:10

= nkx ~

80

~

4t1

-

--..'

kinetic model

~-- -------~

0 0

30

60

90 Time, min

120

150

180

Fig. 5. Comparison of experimental data with those obtained from kinetic model. 2 0 0 - Estrifiction (acidialcohoi 115) acrylic acid + isobutanol temperature 90oc

160 "~

E =

120 80

~

~ '~

Experimental data:

9

4"

H2SO 4 0.1 wt %

~

H2SO 4 0.2 wt %

Q

H2SO 4 0.5 wt %

e

H2SO 4 1 wt %

r------.~

kinetic model

,

<

0

50

100

150 Time, min

200

250

300

Fig. 6. Comparison of experimental data with those obtained from kinetic model. The kinetic results will be applied to the modelling and optimization of industrial reactors. 4. R E F E R E N C E S 1 G.A. Czubarnow, S.M. Danow, W.I. Logutow, O.A. Starikowa, Z. prik. chem., 55 (1982) 1204. 2 G.A. Czubarnow, S.M. Danow, W.I. Logutow, T.N. Obmieluchina, Z. prik. chem., 57 (1984) 203. The w o r k was s u p p o r t e d by K B N ( G r a n t No 3T09C 010 13).

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

419

Hydrogenation o f carbonaceous adsorbed species f o n n e d d t ~ n g the CO/H2 reaction on a Ru/AI203 catalyst: experimental and kinetic modeling.

H. AHLAFIh, M. NAWDALI"and D. BIANCHI a. a-Laboratoire d'Application de la Chimie ~i l'Environnement, LACE, Unit6 mixte 5634. Universit6 Claude Bernard Lyon I, CPE-Bat. 308, 69622 Villeurbanne Cedex, France. b- Laboratoire de Chimie Physique, Unit6 de Catalyse, Facult6 des Sciences , E1-Jadida, Maroc. Abstract:

The carbonaceous adsorbed species detected during the CO/H2 reaction on a RH/AI203 catalyst are mainly the linear CO species and some hydrocarbon species CxHy.It is shown that the CO species can be hydrogenated in pure hydrogen according to a kinetic model assuming a limiting step: the dissociation of CO assisted by hydrogen. 1. INTRODUCTION

In the past two decades, the adsorbed carbonaceous species formed during the CO/H2 reaction on metal supported catalysts were widely studied by transient experiments. The isothermal hydrogenation of the adsorbed species into CH4 is of particular interest because the species can be characterized according to their reactivities versus H2. Previous works on a 10 wt% Fe/A1203 catalyst [ 1-2] have shown that this method gives various information such as: a) the number of hydrogenable adsorbed species, b) the number of elementary steps controlling the kinetic of the hydrogenation, c) the activation energies of hydrogenation and d) the evolution of the number of hydrogenation sites during the reaction. On this catalyst it has been shown [3] that the coverage of the surface by molecular adsorbed CO is very small during the 10%CO/H2 reaction (at T> 200~ and 1 atm total pressure). Mainly surface carbon and CxHy species are present. On ruthenium catalysts, during the CO/H2 reaction, a high coverage of the surface by molecularly adsorbed CO is observed, as on rhodium supported catalysts [4], associated with the presence of CxHv species. However, the adsorbed CO species can be rapidly hydrogenated at the reaction temperature in pure hydrogen. The objectives of the present work is to obtain more insight on the mechanism of this hydrogenation using both FTIR and mass spectrometer as detectors. 2. EXPERIMENTAL 2.1. Catalyst The 3.5%Ru/AI203(in weight percent) catalyst is prepared via impregnation of alumina (Alon-C, Degussa) by the incipient wetness method with an aqueous solution of RuNO(NO3)3, 2 H20 (Johnson Matthey). After drying 24h at room temperature and then 12h at 383 K, the solid is treated in air 12h at 623 K (heating rate 5K/min), leading to a surface area of 90 m2/g.

420 Before either the adsorption of CO or the CO/H2 reaction, the solid is treated in-situ according to the following procedure: helium (T = 713 K, t = 10 min)-> hydrogen (T =713 K, t = 2 h) -> helium (T= 713 K, t = 10 min) -> helium (at adsorption or reaction temperatures). The metallic dispersion of the reduced catalyst is determined by hydrogen chemisorption at 383 K using a volumetric method (ASAP 2000, Micromeritic) and according to a procedure previously described [5]. The quantity of hydrogen irreversibly adsorbed is 74 gmolH2/g of catalyst which leads to a dispersion of D=43%, assuming a chemisorption ratio H/Rus= 1. 2.2. Analytical procedure Two analytical systems (TE 1 and TE2) allowing transient experiments are used. TEl :The transient experiments are performed using a mass spectrometer as detector, according to a procedure previously described [3].Briefly, a quadrupole mass spectrometer is used to determine the composition of the gas mixture (1 atm. total pressure) at the outlet of a quartz micro-reactor. This allows to study the evolution of the rates (lamol/(g of cat. min.)) of either the appearance or the disappearance of the various compounds during a switch between two controlled flows of gas, and finally to determine the amount of adsorbed species on the surface. TE2: The transient experiments are performed using a FTIR spectrometer as detector. A stainless steel IR cell [6] having a small dead volume (transmission mode, pellet of catalyst), with CaF2 windows, allows to execute transient experiments in the temperature range 300-900 K according to experimental conditions similar to those of a differential reactor. 3. RESULTS AND DISCUSSION The 3.5%Ru/A1203 catalysts is active for the 10•CO/U 2 reaction at temperature higher than 453 K. In the present work, all the experiments are performed at 478 K. In a preceding study [7], it has been shown that during the CO/H2 reaction various adsorbed species are formed on the ruthenium particles. The main species are the linear CO (IR band at 2044 crn-~) and some CxHy species (IR bands at 2930, 2857 and 1461 cm1, shoulders at 2960, 2895 and 2830 cm-a). This two species are hydrogenated into CH4 at the reaction temperature during a switch 10%CO/Hz-> H2. The present study concerns the mechanism of hydrogenation of the linear CO species. 3.1 Adsorption of CO at 478 K and isothermal hydrogenation Figure 1 gives the IR spectra recorded after a switch He-> 10%CO/He. According to the litterature, the main IR band at 2044 cml is attributed to linear CO species adsorbed on Ru particles. The other weak IR bands and shoulders are attributed to gem-dicarbonyl species (2140, 2080 crn~) and bridged species (1775 cm-~) on Ru and bicarbonate species adsorbed on the alumina support (1651, 1443 and 1230 cm-1). The disappearance of the linear CO species during a switch 10%CO/He-> He, at 478 K, is very slow, the intensity decreases by 22 % atier 20 min in helium. In another hand the hydrogenation of the linear CO species is very fast as observed in figure 2. Note that during the hydrogenation: a) the IR band of the linear CO species shills to lower wavenumbers due to the decrease of the coverage (2010 cm-1 after 80 s) and b) no IR bands are detected in the range 3100-2800 c m "1 characteristic of C.~I-Iyadsorbed species (there are no detectable adsorbed intermediate species CxHy). The quantity of CO adsorbed (TEl) at 478 K during a switch He--> 5%CO/3%Ar/He is 69 gmol/g of catalyst. CO2 is not detected indicating that the rate of dissociation of CO is very low. After various times on stream in CO/He followed by 2 min in He, a switch He-> Hz

421 is realized. A sharp peak of CH4 with a decreasing exponential profile is detected as observed in figure 3, without the detection of CO (water is not quantified). The quantities of CH4 formed increase slightly with time on stream in CO/He, from 56 lamol/g after 1 min to 62 after 25 min. This increase can be attributed to the slight accumulation of carbon on the Ru surface. It} 9

9

,~r

~o.~ a

o O

9

~000

aooo l~o'o Wavenumber (cm-1)

i-~O0

Wavenumber (cm-1)

Figure 1. Adsorption of CO at 478K. a) lmin, b) 20 min

Figure 2. Hydrogenation of the adsorbed CO a) CO/He, b-d) U2: 20s, 80s, 260s

The difference between the amount of CH4 formed and the quantity of CO adsorbed which is around 13 gmol/g, is attributed to the formation of non hydrogenable species, detected in particular by FTIR aiter more than 4 min of hydrogenation (figure 2, spectrum d). If the hydrogenation of CO is performed at lower temperatures (463 K and 443 K) the same decreasing exponential profile for the production of CH4 is recorded, but with lower values for the rate maximum. T ~ (s~ 0 300 [

IHe[H--d2

I

40

80

120

160

I

-0.~

" ~ 2()0 -

-1

-~.~

C) _= "~ 100 -

-2

-2.5

0

l()0

Time (s)

200

Figure 3. Rate of CH4 production at 478K. COMe: a)l min, b)5min, c)15min, d)25min

Figure 4. Evolution of the intensity of the linear CO species: a) in He, b) in H2

3.2 Kinetic model for the hydrogenation of the linear CO species. The decreasing exponential profile of the CH4 production (figure 3) indicates that a limiting step controls the overall process [1-2]. The various elementary steps which can be considered during the hydrogenation of the linear CO are: a) Desorption of CO: COao~-k'0-> COg,~ b) Dissociation of CO: COad~-k0 -> C,ds + Oars

422 c) H2 chemisorption: H2--> 2 Ha~ d) Hydrogenation of Caa~: Ca~ + Ha~ -> CHaa~ e-g)Successive hydrogenation of the CHx intermediates: CHaa~+ Haa~-->CH2~a~->..->CH4aa~ h) Desorption of CH4 : CI-I4,~ --> CI-I4g,~ Steps a) and b) can be studied separately using the IR data recorded during the treatment in He. The rate of disappearance of the CO species is v0=-d[CO~as] /dt =(k'0+k0)[COaa~] leading to Ln([COaa~]/ [COaas]0)= -(k'0+k0) t, with: [] the superficial concentration and t the time of treatment. This relation can be verified using the change of the intensity, A, of the linear CO species with the time t. Assuming that the integrated absorption intensity of the linear CO species is constant with the coverage: Ln([COaa~]/ [COaa~])= Ln(A/Ao). Figure 4, curve a) shows that the experimental data fit the linear relation relationship very well. This leads to a value (k'o+ko)= 2.1 10-3 s-1. This value is in good agreement with the one found by Cant and Bell [8] on a 4.3%Ru/SiO2 catalyst, i.e 0.7 10-3 s-1 at 473 K. The same relationship can be verified using the experimental data recorded during the hydrogenation of CO (figure 2). A linear relationship is again observed, figure 4 curve b), but the slope leads to a rate constant, 1.6 10-2 s-1, greater than the one determined in curve a). This cannot be attributed to a higher rate of desorption due for instance to a competition with hydrogen chemisorption, because CO is not detected during the production of CH4 (TEl). It seems that during the hydrogenation, a new step is involved which increases the rate of disappearance of the linear CO species. We consider the hydrogen assisted dissociation of CO [9]: COaa~ + Haas -k"0-> C,a~ + OHaa~. The rate of disappearance of CO in presence of hydrogen is now: v~=-d[COaa.~]/dt= (k"0 [Had~]+ k'o+ko ) [CO~dd and assuming [Had~] constant during the hydrogenation then k"0 [Haas]= K"0 = (1.6-0.21)10 .2 s-~ =1.4 10~ s-1. The exponential decreasing profile of the CI-I4 production during the isothermal hydrogenation (figure 3) indicates that a limiting step controls the process. If it is assumed that this step is the H2 assisted dissociation of CO, the rate of formation of CI-h is : v(CH4) = K"o [COaa~]0 exp(-K"0 t) and the initial rate (at t = 0) is v(CH4)o= K"0 [COaa~]0. Figure 2 shows that the value of the initial rate of Ct-I4 production at 478 K is around 190 [amol/(g.min), while the quantity of linear adsorbed CO is around 60 lamol/g. This leads to K"0-5 10-2 s-~, a value not very different than the one determined using the FTIR data (1.4 10-2 s~). This seems to indicate that the production of CH4 is controlled by the hydrogen assisted dissociation of CO. 4. CONCLUSION The linear adsorbed CO species present with a high concentration on the ruthenium particles during the 10%CO/H2 reaction at 478 K is rapidly hydrogenated into CH4. A limiting step controls the CH4 production: the hydrogen assisted dissociation of CO. 5. 1 2 3 4 5 6 7 8 9

REFERENCES D. Bianchi and J.L Gass, J. Catal., 123 (1990) 298. H. Ahlafi, C.O. Bennett and D. Bianchi, J. Catal., 133 (1992) 83. D. M. Stockwell, D. Bianchi and C.O. Bennett, J. Catal., 113 (1988) 13. E.M. Efstathiou, T. Chafik, D. Bianchi and C.O. Bennett, J. Catal., 148 (1994) 224 H. Ahlafi, M. Nawdali, A. K. Bencheikh and D. Bianchi, Bull. Soc.Chim. Fr., 133 (1996) 461 T. Chafik, O. Dulaurent, J.L. Gass and D. Bianchi, J. Catal (in press) H. Ahlafi, M. Nawdali, A. K. Bencheikh and D. Bianchi, Bull. Soc.Chim. Bel., 106 (1997) 245 N.W. Cant and A.T. Bell, J. Catal., 73 (1982) 257. S. Y. Wang, S. H. Moon and M. A. Vannice, J. Catal., 71, (1981) 167

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

423

Application of the continuous two impinging streams reactors in chemical absorption Morteza Sohrabi and Amir Masoud Jamshidi Department of Chemical Engineering, Amirkabir University of Technology, Tehran 15914, Iran

Abstract A stochastic model for the residence time distribution in continuous two impinging streams reactors with spray nozzles (CISR) has been developed based on Markov chain processes. The performance capability of the CISR in chemical absorption operations has been demonstrated by a typical example of reaction between carbon dioxide and monoethanolamine.

1. INTRODUCTION The coaxial two impinging streams devices which utilize a unique flow behaviour to intensify transfer processes in heterogeneous systems was first described by Elprin [1]. In such an apparatus, two streams flowing counter currently on the same axis collide with each other at the zone in which the two streams impinge. Continuous two impinging streams reactors with spray nozzles (CISR) have shown significant efficiencies in various chemical processes such as mixing of gases and solids [2], adsorption [3,4], drying I5], dissolution [6] and two phase chemical reactions [7]. In the present study, an attempt has been made to develop a model for the residence time distribution (RTD) in CISR. In addition, the performance capability of such devices in two phase gas-liquid reactions has been investigated.

2. EXPERIMENTAL

2.1. Reactor system The reactor consisted of a cylindrical vessel made of "Pyrex" glass, length 60 cm and internal diameter 9.2 cm. The vessel was equipped with two spray nozzles made of stainless steel type 316, placed on two movable coaxial circular plates made of Teflon positioned against each other at the two ends of the reactor. Thus the length of the reaction compartment could be varied by moving the plates away from or toward each other.

424 The design of nozzles has a significant effect on the intensity of liquid dispersion within the gas phase, the droplet size distribution and the velocity of phases. Each nozzle consisted of two basic sections: the main body and the middle part. The latter was screwed inside the main body. The flow of liquid phase to each nozzle was via a central port (3 mm diameter), while that of the gas phase was through three openings, each 3 mm in diameter, spaced around the central port.

2.2. Start-up procedure The experimental system is shown in figure 1. In each run the system was first adjusted using water and air flows. When RTD was to be determined, a pulse of a colour tracer was injected instantly into the water inlet and the exit stream from the vessel was collected in a series of sampling bottles until no further tracer was observed in the effluent. The concentration of tracer within each sampling bottle was determined by UV spectroscopy. The pattern of change in effluent colour was also observed by a high speed photographic technique and video recording.

Fig. 1. The experimental sot-up. 1, Cylinderical vessel; 2, thermometer; 3, funnel; 4, Teflon ring; 5, spaff nozzels; 6, gas flow meter: 7, liquid flow meier, 8, liquid tank: 9, compressor; 10, gas tank.

In the case of chemical absorption experiments, the flows of water and air were replaced by those of monoethanolamine (MEA) solution and carbon dioxide gas respectively. When steady state conditions were established, samples were drawn to the sampling bottles and the extent of reaction was determined by chemical quantitive analyses. The results were fully reproducible with a mean absolute deviation of 5-7 %. Further details concerning the reaction system and analytical procedures may be found elsewhere [8].

3. RESULTS AND DISCUSSION 3.1. Modelling the residence time distribution in the reactor A model for the RTD in CISR was developed based on models first proposed by Van de Vusse [9]. As the collision of the droplets in the impingement zone is random, a suitable mathematical technique to handle such a process could be Markov chain models [ 10,11].

425 According to discrete-time Markov chains, the probability of an event at time t + l (t = 0, 1, 2, ...) given only the outcome at t i m e t is equal to the probability of the event at time t-t- 1 given the entire history of the system. In other words, the probability of the event at t + 1 is not dependent upon the state history prior to time t. Thus, the values of the process at the given time t determines the conditional properties for future values of the process. These values are called the state of the process and the conditional properties are thought of as transition propbabilities between the states i and j, PO" These values may be displayed in a matrix (P = [P0]) called the one step transition matrix. The matrix P has has N rows and N columns, where N is the number of possible states for transition of the system. The rows of matrix P consist of the probabilities of all possible transitions from a given state and so sum to 1. )v

~",P ij = 1

(1)

j=l

This matrix completely describes the Markov process. Further details of Markov models may be found elsewhere [8, 10, 11]. By considering the patterns of liquid flow within the vessel, the reaction compartment was divided into eight regions with equal volumes (Fig. 2). Each region represents a state in the Markov process. A recycle stream R was also assumed due to counter current flows. A typical RTD curve calculated for the reactor is shown in Fig. 3. Rt2

Q/2

7

(Q+ R)/2

Fig. 2. The flow regions proposed for the reactor: 1-7,perfect mixing regions; 8,plug flow region.

Fig. 3. R e s i d e n c e

~

Time

Distribution

k

--

model

0.8 ~ (3

b

0.6

0.4 0.2

0

i

i

i

i

i

i

l

2

4

6

8

10

12

14

i~-"=--,..I---=~

16

18

=

i

a

20

22

24

,,. _

26

Time (s) Mean residence time = 5 91 s; Variance = 12,75s^2; Recycle ratio = 0.53; Inter nozzle diameter = 14 cm; Reactor diameter =9.2 cm

426 3.2. Chemical absorption of carbon dioxide in monoethanolamine (MEA)

Using the experimental set up shown in Fig. 1, absorption of C O 2 gas in MEA was investigated. The chemical reaction between the two reactants is given by the following equation: [12,13] C O 2 q-

2 RNH2 --+ RNCO-2 + RNH+3

(2)

where RNH2 represents MEA. The chemical reaction is very rapid so that the diffusion of C O 2 in liquid droplets may be assumed to be involved in the overall rate controlling step of the process. By considering a material balance around a differential thickness (dx) within the reaction zone adjacent to the interface and integration of such an equation, the following relationship may be obtained:[14,15] CO*

roveran= a ( 2Dc I r(C~)dC~

(3)

0

where a is the interfacial area between gas and liquid phases; D~ is the effective diffusivity of CO2 in MEA; C~ and Co* are the bulk and interfacial concentrations of carbon dioxide respectively and /'overall is the overall rate of reaction. It may be possible, therefore, to estimate the interfacial area between the two phases in the impingement zone by applying eqn (3). The overall rate may be calculated from the RTD model [7]. Under a wide range of experimental conditions, the interfacial area was round to vary between 0.15 and 1.5 cm 2 cm -3. 4. REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15

I. Elprin, Inzhenerno Fizicheskii Zhurnal, 31 (1961) 62. A. Tamir, AIChE.J., 31 (1985) 1744. A. Tamir, Chem. Eng. Sci., 40 (1985) 214. A. Tamir, Chem. Eng. Sci., 41 (1986) 3023. A. Tamir, Chem. Eng. Prog., 85(9) (1989) 35. A. Tamir and M. Grinholts, Ind. Eng. Chem. Res., 26 (1987) 726. M. Sohrabi, T. Kaghazchi and F. Yazdani, J. Chem. Tech. Biotechnol., 58 (1993) 363. A. M. Jamshidi, Graduate Thesis, Amirkabir University, Tehran, 1995. J.G. Van de Vusse, Chem. Eng. Sci., 17 (1962) 507. H. Stark and J. W. Woods, Probability, Random Processes and Estimation Theory for Engineers, Prentice-Hall, Englewood Cliffs, N.J., 1986. A. Papoulis, Probability, Random Variables and Stochastic Processes, McGraw Hill Inc., New York, 1991. C. Alvarez Fuster, N. Midoux, A. Laurent and J. C. Charpentier, Chem. Eng. Sci., 35 (1980) 1717. S. S. Laddha and P. V. Danckwerts, Chem. Eng. Sci., 16 (1981) 479. M. Sohrabi, Chimia, 37 (1983)465. M. Sohrabi, Afinidad, 43 (1986) 34.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

427

A kinetic s t u d y of H e c k r e a c t i o n of iodobenzene a n d m e t h y l a c r y l a t e using homogeneous Pd/TPP catalyst F.-G. Zhao, B. M. Bhanage, M. Shirai and M. Arai* Institute for Chemical Reaction Science, Tohoku University, Katahira, Aoba-ku, Sendal 980-8577, Japan Abstract A kinetics of Heck vinylation of iodobenzene with methyl acrylate has been studied using homogeneous Pd(OAc)2/TPP catalyst in the temperature range of 50-70~ The rate of reaction shows linear dependence with iodobenzene concentration. The effect of catalyst and methyl acrylate concentrations shows linear dependence initially and marginally increases afterwards. The effect of triethylamine concentration passes through a maximum. An empirical rate equation has been derived to fit the experimental data, showing accuracy in the range of -+8%, and the activation energy was found to be 23 kcal/mol. 1. INTRODUCTION Heck vinylation of aryl halide with olefins is one of the important methods for the formation of new C-C bonds and finds several industrial applications in the synthesis of various chemicals [1,2]. In the literature, there are numerous publications on Heck reactions using homogeneous and heterogeneous catalysts. In most of these papers, the influence of various operating variables including catalysts, promoters, solvents, substrates, and others on the reaction rate and selectivity has been reported. Although reaction mechanisms were proposed [1,2], little attention has been paid to detailed kinetic analysis of this reaction. Then we have embarked in the investigation of kinetic analysis of this important reaction. Heterogeneous Heck reactions by supported metal catalysts and in biphasic operation modes are also gaining interest, since they avail advantage of catalyst and product separation. Kinetic analysis using homogeneous catalyst is fundamental and useful for the kinetic analysis of heterogeneous reactions. In the present work, we have studied kinetics of Heck reaction of iodobenzene and methyl acrylate using homogeneous Pd/TPP (TPP: triphenyl phosphine) catalyst considering applications of methyl cinnamate ester [3]. The effect of various reaction parameters has been studied in the temperature range of 50-70~ and a suitable rate model has been derived.

428 2. EXPERIMENTAL All the experiments were carried out using a glass reactor well stirred with a Teflon stirrer. In a typical experiment, Pd(OAc)2 (0.1 retool), TPP (0.2 retool), Et~N (10 mmol), methyl acrylate (10 retool), and N-methyl pyrrolidone (solvent, 15 ml) were mixed together under atmospheric pressure of argon. The reaction mixture was heated to the desired temperature and then iodobenzene (10 retool) was added. The reaction mixture was sampled repeatedly and analyzed by gas chromatograph. Under the present conditions used, trans-methyl cinnamate was observed with 100% selectivity. The rate of reaction was calculated from the amount of iodobenzene consumed for initial 20% conversion. 3. RESULTS A N D DISCUSSION The Heck vinylation of idobenzene with methyl acrylate was studied using palladium acetate and TPP catalyst system. The reaction scheme is as follows:

[~l+ H2C"~'~COOCH3 Pd(OAc)2~COOCH3

The effect of various parameters s u c h a s catalyst, iodobenzene, methyl acrylate, triethylamine and ligand concentrations was studied at temperatures of 50-70~ Effect of agitation speed has been studied and it has been ensured that the reaction is occurring in kinetic regime. Figure I shows the initial rate of reaction is linearly proportional to iodobenzene concentration. According to well accepted mechanism [2] for Heck reaction given in Scheme 1, iodobenzene is oxidatively added to bis(triphenylphosphane)palladium(0) generating r palladiumaI) complex. This is the first step in the mechanism and as concentration of iodobenzene increases it leads to generate more and more active species and hence linear dependence can be explained. Effect of methyl acrylate concentration is shown in Figure 2. It has been observed that initially the rate increases linearly and increases marginally afterwards. Addition of olefm to catalyst complex leads to formation of u complex in an equilibrium reaction and this is followed by c complex formation. Since this is equilibrium reaction, at higher methyl acrylate concentration the rate of reverse reaction is likely to be predominant leading to marginal rate dependence. The effect of triethyl_Amine concentration passes through a maximum as shown in Figure 3. The maximum appears at I kmol/m ~, which is equivalent to 0.7 on stoichiometry with other reactants. BeUer and Riermeier [4] have reported similar observation in vinylation of aryl bromides with butyl methacrylate. Base takes part in reductive elimination of HX to regenerate active catalytic species (step E, Scheme 1). With excess of base, it may be inhibiting active catalytic sites and hence maximum is observed. The

429

7~0

Et=N.HI Phi ~'~ ,,[(PPh~Pd] ( oxida~veaddit~ Et~N ~ A ~

0.008

tHpd(PPh3)2r J

? 0.006

~/",~'c __#5~x__. 'P

OOCHs]O

"~ 0.004

syn e~iminatk)nH~,r

0.608

/

0

0.5 Iodobenzene,

internalrotation

1 1 kmol/m 3

Effect o f i o d o b e n z e n e concentration

S c h e m e 1.

"

i

~ : ,yn insertion

e'n~#Td(PPh~21

H

Figure 1. 7

a ~'"~"COOCH~

pd(PPh3),i --

Ph~C_OOCl%XX

d 0.002

0

[Ph-Pd(~

COOCH~

M e c h a n i s m of H e c k r e a c t i o n

0.01

? 0.006

0.008

"~ 0.004

~ 0.006 0.004

,90.002

~o.oo2

t~

0

0

0.5 1 1 Methylacrylate, kmol/m 3

F i g u r e 2.

7

Effect of m e t h y l a c r y l a t e concentration

0.01

0

0.5 Triethylamine,

F i g u r e 3.

1 1 kmol/m 3

Effect of t r i e t h y l a m i n e concentration

0.02

? 0.008 -

~

~

0.006

. _ j ~ ~ 6 o oc

O

0.004

"~ 0.01

oC

d 0.002 0

0

0

F i g u r e 4.

.

,

.

I

.

,

.

I

.

,

.

I

.

0.002 0.004 0.006 Pd, k m o l / m 3

Effect of c a t a l y s t concentration

,

.

0.008

0

0

F i g u r e 5.

0.002

0.004

0.006

TPP, k m o l / m 3

Effect of T P P c o n c e n t r a t i o n

430 effect of catalyst concentration (Figure 4) shows increase in rate with increase in catalyst concentration, but not linearly as observed in many cases. This may be compounded effect of ligand and catalyst precursor concentrations, since, in this case, we varied both ligand and catalyst precursor to keep P d ~ P P ratio constant at 2. Effect of TPP concentration shows (see Figure 5) negative order dependence. In the absence of TPP, maximum rate is observed; however, it is followed by Pd precipitation. So TPP gives stability to the catalyst at the expense of reaction rate. Herrmann et al. [5] observed similar trends for the effect of TPP concentration. Afterwards we have analyzed the data obtained for kinetic analysis. In order to fit the rate data, several rate equations were examined using nonlinear regression analysis. The optimization program based on Marquadt method was used. The following equation was found to fit the rate data with - 8 % accuracy. kAB2C D Rate = (1 + KBB~) (1 + KcC 4) (1 + KDD)3 (1 + KEE) where, k = rate constant, ml~/(kmoP s); A = concentration of iodobenzene, kmol/m3; B = concentration of methyl acrylate, kmol/m~; C= concentration of triethylamine, kmol/m3; D = concentration of catalyst precursor, kmol/m~; E = concentration of TPP, kmol/m3; KB ,Kc ,KD and KE are respective constants. The optimized values of these constants are given in Table 1. Tablel Optimized values for various constants Temperature, ~ 50 60 70

k, m12/(kmoPs) 7.6363 23.2900 67.4565

KB, m6/kmol2 4.3251 5.4616 9.1013

Kc, m'2/kmoP 0.3307 0.4348 0.3948

KD, m~/kmol 47.0966 6.9880 1.3547

KE, m3/kmol 10.628 163.21 251.94

The activation energy was found to be 23 kcal/mol. 5. REFERENCES [1] R. F. Heck, Organic Reactions, Vol. 27, John Wiley & Sons, New York, 1982, 345-393. [2] A. De Meijere, F. E. Meyer, Angew. Chem. Int. Ed. Eng., 33 (1994) 2379. [3] D. Garbe in Ullmann's Encyclopedia of Industrial Chemistry, A7 (1992) 99. [4] M. Beller, T.H. Riermeier, Tetrahedron Lett., 37 (1996) 6535. [5] W. A. Herrmann, C. Brobmer, K. Ofele, M. BeUer, H. Fischer, J. Mol. Catal. A, 103 (1995) 133.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

431

Kinetic and catalytic aspects in the synthesis of polyethylene terephtalate (PET), also through the use of model molecules B. Apicella., E. Santacesaria ,M. Di Serio Dipartimento di Chimica, University of Naples "Federico [I", Via Mezzocannone 4 (80134) Napoli, (Italy). Abstract Kinetic and catalytic aspects of the two steps of formation of poly(ethylenterephtalate) (PET) have been studied in this work using model molecules. Many kinetic runs have been performed using different catalysts. We have made kinetic runs on some catalysts under different operative conditions to evaluate the effect of catalyst concentration and temperature on the reaction rate. All kinetic runs have been interpreted and kinetic parameters determined. Suggestions on the reaction mechanisms conclude the work. 1. I N T R O D U C T I O N The production of polyethylene terephtalate (PET) from dymethylterephtalate (DMT) is industrially performed in two steps. In the first, the melt transesterification of DMT with ethylene glycol (EG) occurs with the formation of 1-10 monomeric units oligomers, which are of three different types having respectively hydroxyl-hydroxyl, hydroxyl-methyl or methyl-methyl terminal groups. The second step is a polycondensation occurring at high temperature and under vacuum with the elimination of EG. The authors have recently developed a kinetic model describing the evolution with time of all the oligomers formed during the reaction as regards the first step of PET synthesis [1]. This model has also been applied to different catalytic systems particularly to zinc acetate catalyst [2]. Despite the complexity of the kinetic model characterized by a great number of reactions and products, the comparison between the different catalysts is relatively simple depending on the activity in two steps of these reactions: the reaction of a methyl group with an hydroxyl of ethylene glycol (k~) and that of a methyl group with an hydroxyl terminating chain (k2), being both equilibrium reactions of second order. Thus, if the reaction scheme is truncated at the fourth stage, the model consists of 58 reactions and 24 oligomeric species but only the two mentioned kinetic parameters are necessary to describe the behaviour of the system. In order to deep kinetic studies and semplify the comparison between the different catalysts, considering the analytical difficulty in determining the concentrations of all possible oligomers, we have studied the kinetic behaviour of transesterification test reactions [3]. Then, the use of test reactions have been extended also to the study of policondensation reaction [4], that is, the second step in the synthesis of PET. _

2. R E S U L T S A N D D I S C U S S I O N The test reactions used for studying the first step of the synthesis of PET are: COOCH 3

CCK)CH2CH2OH

X

X

432 COOCI{2CH2OH C O O C H 3

COOCH2CH200C

X

X

where X = H, CH3,NO2,OCH3. In other words, two reactions occur with different rates and equilibrium constants, the former involving one hydroxyl of the free ethylene glycol molecules, the latter the residual hydroxyl of the bonded EG, that is, the same reaction types occurring in the DMT transesterification. Such reactions can more easily be followed by analyzing reagents and products of the test reaction by gas-chromatographic analysis. We have observed, however, that the nature of the substitutent X strongly affects the catalytic activity and, only for X=NO2, that is a group with a strong electron withdrawing character, we have an activity comparable with that of DMT. Another important remark is that catalytic activity depends on the catalyst concentration in a complex way, that is different for each catalyst. Thus, it is not correct to give a simple correlation, as made in some papers published in the literature [5-8], between activity and acid character of the metal ion (volcano shaped curves), normally made at one fixed catalyst concentration. A comparison of the activities obtained in the transesterification of different model molecules in the presence of different metal ions is reported as a function of metal acidity in Fig. 1. F i g . l - O v e r a l l activities of proven catalysts for different substrates (MA=methyi p-anisate, MT=methyl p-toluate, MNB=methyi p-nitrobenzoate, MB=methyibenzoate, DMT=dimethylterephtalate) as a function of the acidity factor [3 (stability constant of dibenzoyimethane complex) --II-- MA

K M + K D ( L m o l l m i n -~)

oo18

~

'

[~

'

l-~ ,

0.016

'

i

0 014

[ --O-- MT

[~r~' j,,

l --A-- MNB I --V1 El

:

MB DMT

0 012 0 010 0 008

,

0006

, ,

, ,

0.004

".\

-~ , -\

---- -1~:~'-

o 002 0 000

.,.,

-...

: I '

s5

' ~

! ,I

90

, '

I

,

.\.

: 9s

ff~---

~

', , i

\... e-..# I

~0.0

, "iF

,

l

~os

Logl3

In the same figure the activities obtained for dimethylterephtalate (DMT) are also reported for comparison. As it can be seen, the maximum of activity changes for any kind of used model molecule. With regard to the second step of PET synthesis the kinetic study is much more difficult, since starting with a mixture of oligomers (a circumstance generally neglected in the literature by assuming that reaction starts with pure 2, bishydroxyethylterephtalate BHET), we obtain, removing EG, polymers of high molecular weight with a scarcely known molecular weight distribution. It is difficult for such a system to perform catalytic screening in a fast and simple way and to study the catalytic mechanism. For this purpose the following test reaction could be suggested [4]:

433

2

--.

-

,

X

X

+

HOCI-L-,CH2OH

X

The evolution with time of this reaction can be easily followed by analyzing the composition of the reacting mixtures by a gaschromatograph. The forward and reverse reactions turned out to be of second order. Therefore, the kinetic and equilibrium parameters can be easily determined. Moreover, we have observed that the influence of the substituent X, in this reaction, is very small for tri- or tetravalent metals, normally used as catalysts of this reaction, on the contrary, the influence is large for bivalent metals, suggesting a different operating reaction mechanism. The equilibrium constant turned out to be about 0.17-0.20 at 200~ Kinetic constants obtained for some of the experimental metals are reported for comparison in Fig. 2 as a function of metal acidity. From data collected it results that titanium has a high activity at a very low metal concentration, but activity is poorly affected by the increase of metal concentration. Fig.2-Activities of proven catalysts in polyeondensation reaction test as a function of another acidity factor (IP)Z/R, with IP=ionization potential, Z--valency of cation, R=radius of cation 1

log k

Ti (IV)

0.5

Z~_~ _ ~ Sb(Ill)

-0.5

=Mn(ll)

1.5

2

2.5

3

log ((IP) Z/R)

In fact, the reaction order with respect to the catalyst turned out to be 0.09 for titanium, 0.2 for antimonium and 0.5 for molibdenum. As other authors have already proposed [5, 8,1 l], the mechanisms of condensation are probably very similar to those of transesterification. Two mechanisms seem to be operative promoted respectively by bivalent and tri- or tetravalent metal catalysts. Bivalent catalysts attack the carbonylic oxygen favouring the successive nucleophilic attack of a glycoxide oxygen to the carboxylic carbon atom, as in the following scheme: M2+

CH2CH2OR

Nf2+

It," /

II " ~ ~

X

_

+

HOCH~CH2OH

X

L R= ~ o r H

with x and X - H or CH3 In this case the metal has a strong inductive effect on the carbonyl group and influences the aromatic ring resonance. For this reason, para substituent strongly affects the performances,

434 as observed in particular in transesterification reaction. On the contrary, tri- and tetravalent metals, are preferibly coordinated to the acylic oxygen and this favours the nucleophilic attack to the adjacent carbon from the alkoxide coordinated to the metal [9,11, 12], in two or three steps, as in the following example: o ~---(_~Z.:H~-CHe-OH +

\H

- , -

CH2

CI~

x

0 ~--O---CH2-CH--OH /

HOCH2--X2t{2/(;k~ b ' ~ j ' H2 CH~

x

_

\0

HOCI-b-Ct~,b/~l

NO

%.x

H2

O o ~--O--CH2--CHI~-O-~

~- ~ , , \ H/

aH~ +

O/"

X

0 x

where X = H or CH3.

4. R E F E R E N C E S 1 E. Santacesaria, F. Trulli, L. Minervini, M. Di Serio, R. Tesser, S. Contessa- J. Appl. Polym. Sci., vol.54, 1371-1384 (1994) 2 M. Di Serio, R. Tesser, F. Trulli and E. Santacesaria- J. Appl. Polym. Sci., vol.62, 409415(1996) 3 M. Di Serio, B. Apicella, G. Grieco, P. Iengo, L. Fiocca, R. Po, E. Santacesaria- J. Molecul. Catalysis A Chemical, 1691 (1997) 4 B. Apicella, M. Di Serio, L. Fiocca, R. Po, E. Santacesaria- J. of App. Pol. Science in press 5 J.S.Chung- J. Macromol. Sci.: Chem. A 27, vol.4, pp.479-490, (1990) 6 K. Tomita, H. Ida - Polymer, vol. 16, 185-190 (1975) 7 K. Tomita, H. Ida - Polymer, vol. 14, 55-60 (1973) 8 K. Tomita - Polymer, vol. 17, 221-224 (1976) 9 S.B:Maerov- J. Polym. Sci. Polymer Chem. Ed., vol. 17, pp.4033-4040 (1979) 10 G.Rafler, G. Reinish, E.Bonatz-Acta Chim. (Budapest), vol.81, pp.253-260 (1974) 11 Parshall- Homogeneous catalysis, pp.269-275 (1990) 12 J. Otton, S. Ratton, V. Vasnev, G. Markova, K.Nametov, V.Bakhmutov, L. Komarova, S. Vinogradova and V. Korshak- J. Polym. Sci., Part A, 26 (8), 2199-2224 (1988)

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

435

H / D isotopic exchange between oxide surface and spiitover hydrogen on nickel supported c a t a l y s t s .

V. Almasan, Mihaela Lazar, P. Marginean.

Institute of Isotopic and Molecular Technology, P. O. Box 700, 3400 Cluj Napoca 5, Romania.

Abstract Exchange of D 2 with OH groups of six supported nickel catalysts was performed in a typical reaction system for transient experiments. The rate determinig step is the hydrogen spillover. The HD yield is independent on the support nature, but strongly dependent on the number of OH groups on the oxide surface.

1. I N T R O D U C T I O N . The hydrogen - deuterium (H/D) isotopic exchange was performed between D2 gas and the hydrogen binding to oxide surface of six supported nickel catalysts (Ni/A1203 , Ni/Cr203 , Ni/MgO, Ni/SiO2, Ni/ZnO, Ni/ZrO2). The catalyst samples were prepared by coprecipitation method. BET surface area and metal dispersion were measured by krypton adsorption and hydrogen chemisorption, respectively.

2. E X P E R I M E N T A L The experimental investigations were made in a typical reaction system for transient experiments, which consist in a catalytic reactor on line with a quadrupole mass spectrometer (QMS). In every case the isotopic exchange reaction follows the same way: the input gas (D2) passes through the catalytic reactor, and the output gases are continuos monitored by QMS. Isothermal isotopic exchange (I.I.E.) was performed at seven different temperatures ranging between 21 ~ and 300~ In all the I.I.E. experiments the QMS monitories 2, 3, 4 masses (H2, HD, D2 respectively), and recorded an MS diagram like in Fig 1.

3. R E S U L T S A N D D I S C U S S I O N The MS diagrams were transformed in D 2 adsorption and H 2 desorption kinetic curves with an adequate mathematic program [1]. Fig.2 presents such of kinetic curves. In the same conditions the shape of D 2 adsorption plot is different from the H 2 desorption plot. If are

436 compared the slopes of these two kinds of curves it is found that the processes associated to D2 adsorption at the beginning are more rapid than the spillover process which caused the H2 desoption curve. In this case it can be remarked that the rate-determining step (rds) is the hydrogen spillover [2]. The adsorption plot passes through a maximum before it reaches the plateau value, and the desorption plot have no extreme point.

Figure 1. The MS diagram for I.I.E. reaction on Ni/SiO2 100

90 80 H2

7O o~

t,,.-

=

60

.E

50

a.

40

v

(9

9 HD ._ D2

30 20 10

0

5

10

15

20

25

30

Time(min)

The maximum of the adsorption plot arises in the same time with the maximum yield of HD fig.1. The adsorption plot was obtained by the summarisation of D 2 yield and 1/2 of HD yield, and similar for H 2 desorption curve. The source of H2 is only the H atoms transferred from oxide surface. The HD molecule is produced only on the metal surface. The HD molecules produced are desorbed directly from this surface, and are responsible for the difference between adsorption and desorption. The yield of HD is a measure of hydrogen amount binding from the oxide surface. For this reason the HD production have the same variation like OH number upon the oxide surface, which is illustrated in fig. 3 and table 2. The hydrogen spiltover that is the rds of IIE takes place very easy [3] and is not influenced by the surface coverage. The values from table 1 of the apparent activation energy determined for the six catalysts are independent of the oxide nature and are very low (up to 1 kcal / mol).

437 Table 1. The apparent activation energies for hydrogen spillover process (kcal/mol). Catalysts Ni/Cr203 Ni/MgO Ni/SiO2 Ni/ZrO2 Ni/A1203 El 0.5 0.8 0.7 0.4 0.7 E2 0.9 0.4 0.5 0.4 0.5

Ni/ZnO 0.9 0.7

El, E2 - the apparent activation energy for low and high hydrogen coverage of the catalyst surface, respectively.

Figure 2. The D 2 adsorption and H2 desorption curves on Ni/SiO 2

25

_L

20 03

E

D2

. ~ , = = = ~ 50 ~

H2

15

300 ~

,._..

E --= O >

10

H2 0

I~0

- ;

10

15

20

2'5

3'0

35

40

Time(min)

Table 2. The influence of temperature over the OH population from oxide support (OH number/rim2).

Ni/Cr203 Ni/MgO Ni/SiO2 Ni/ZrO2 Ni/A1203 Ni/ZnO

21 9.6 20.0 9.8 16.2 6.8 5.5

50 9.6 19.8 9.6 16.0 7.5 8.4

100 8.3 16.4 8.2 16.4 7.2 6.5

Temperature 150 8.3 14.7 6.3 12.5 7.0 8.3

200 7.4 11.2 5.1 9.6 6.0 7.3

250 6.6 10.9 4.9 8.0 6.6 6.0

300 6.0 10.9 5.1 7.2 4.6 6.0

438 Figure 3. The HD yield as function of temperature on Ni/SiO2 16

21 ~

14

100~

12 --eq

10

"6" E

8

o >

6

E r

150~ 200~ 250~ 300~

0

5

10

15

20

25

time(min)

4. C O N C L U S I O N S The following conclusions can be drawn from this study: - The associated processes to D2 adsorption (dissociative chemisorbtion of hydrogen isotopes and liE of hydrogen) can not be rds. The rds of the reaction H2 + D2 - 2HD over Ni supported catalysts is the spillover of hydrogen. - The HD yield is proportional to the occupied degree of oxide surface with OH groups.

5.

REFERENCES

1 2 3

V. Almasan, I. Hodor, P. Marginean, Appl. Surf. Sci., 120 (1997) 335. D. Martin, D. Duprez, J. Phys. Chem.B, 101 (1997) 4428. U. Roland, T. Braunschweig, F. Roessner, J. Mol. Catal. A: Chem., 127 (1997) 61.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

439

Transient Studies o f Adsorption Kinetics J. Kanervo b , L.B. Backman b, A.O.I. Krause b, S-L J~ims~i-Jounela a aLaboratory of Process Control and Automation, Helsinki University of Technology, P.O. Box 6100, FIN-02015 HUT, Finland bLaboratory of Industrial Chemistry, Helsinki University of Technology, P.O. Box 6100, FIN-02015 HUT, Finland 1. ABSTRACT This paper deals with the adsorption of hydrogen on nickel. The kinetic p a r a m e t e r s of a Langmuir-type adsorption model were estimated based on hydrogen pulse experiments conducted under isothermal conditions at several temperatures. 2. INTRODUCTION In general the term 'transient' refers to changing one or more of the system parameters. Transient state methods have applications in reactor modelling, optimisation and control. In transient kinetic studies a dynamic change is introduced into a reactor system, and its response in some reaction quantity is observed. Transient state is utilised in m a n y different kinds of kinetic studies. Roughly, transient kinetic methods can be classified into two main groups. In the first group the thermodynamic s t a t e of the system is changed. A dynamic change is introduced in flow rate, pressure, t e m p e r a t u r e or inlet concentration. The most common methods of this type are concentration step or pulse response methods and TPD. The second group where the system remains in the state determined by thermodynamics includes various isotopic labelling techniques. Transient responses can be analysed to obtain more kinetic information t h a n is possible using traditional steady-state experiments. The goal of this work is to study the mechanism of toluene hydrogenation by using transient methods [1]. Due to the complexity of the considered model hypotheses and the large number of required model parameters, the strategy was to separately examine the adsorption-desorption dynamics of the reactants and the products. This paper contains adsorption studies of hydrogen on a supported nickel catalyst. 3. EXPERIMENTAL The hydrogen adsorption experiments were carried out at 30-180 ~ under atmospheric pressure using an Altamira AMI-100 pulse micro reactor. The outlet s t r e a m was analysed by a mass spectrometer. Extra attention was given to the calibration of mass spectrometer, because very low concentrations were to be measured reliably. The concentration as a function of ion current was calibrated at multiple concentrations. The calibration result was validated to be satisfactory by comparing the actual molar amount of the input hydrogen pulse with the molar amount obtained by the integration of concentration signal. The pulse experiments were assumed to be dynamically more beneficial t h a n step

440 changes, because the catalytic system is in transient state having both increasing and decreasing surface coverage during the hydrogen input. Hydrogen was pulsed into the inert carrier gas every 180 second. Hydrogen concentration at pulse m a x i m u m was kept below 1 mol-% in order to m a i n t a i n isothermal conditions. The mass spectrometer data was recorded about 4 times/second. A commercial Ni/AlzOa catalyst, particle size < 0.1 mm, was used. The catalyst was reduced at 335~ for 120 min. The mass of the catalyst used in these experiments was 20 mg. 4. M O D E L L I N G If dissociative Langmuir-type chemisorption of hydrogen (1) is assumed, H~.+ 2*

++ 2H*

(1)

the t r a n s i e n t pseudohomogeneous one-dimensional mass balance in a packed bed reactor becomes (2) : G~CH2

Ot

= -v

C~CH2

8z

9 -+ (k,,C_H. _ kACHC,2 ) 1 ,7.,, ~,

8CH* = 2(kACH C,~ &

-

k,,C2H .

(2)

)

where CHz is the concentration of hydrogen in the gas phase, CH, is the concentration of adsorbed hydrogen, Cv is the concentration of the vacant sites, t is time, z is the axial distance along the catalyst bed, ~b is the void fraction of the catalyst bed, kD is the desorption rate constant and kA is the adsorption rate constant. In the actual calculations the concentrations were transformed into molar fractions in the gas phase and fractional coverages on the catalyst surface (3). Applying the ideal gas law for the gas phase gives CH~ = xHz*ptot/RT. The surface concentration is written as fractional coverage of CH*= Ntot OH*, where N t o t - C v + CH* r 0 v - 1- 0 H*. The modified system becomes: (~XH2

8t Ot

= -v =2

OXH2

Oz

1 -+ (k,,N ,,,,0 2 H. 2 R T --kAXH~N,,,,2 ( 1 - 0 H.)2 ) ~

p,,,,

kAN,,,, XH2 (1 -- OH. )2 -P,,,, --~--

0%

s

(3)

kt~N'"'O""2

For a unique solution the required boundary and initial conditions are: XH~ = XH,- (O,t) = f ( t )

x.2 (z,0) = 0 0.. (z,0) = 0

(4)

The conditions above correspond to the case of imposing a concentration pulse of hydrogen on the carrier gas when no hydrogen species exist initially in the reactor. The function fit) is the time dependency of the hydrogen fraction at the reactor inlet. This function was obtained by a blank run of the input pulse through an empty reactor and

441 was given to the pde-solver in numerical format. The model assumes that the adsorption enthalpy is independent of the coverage on the catalyst surface. The model also assumes that the linear gas velocity is constant in time and place. The catalyst bed should be isothermal, axially equally efficient, and no radial concentration gradients are allowed. The model assumes only one type of active sites. However, temperature programmed desorption studies have demonstrated the presence of surface heterogeneities on the nickel catalysts: hydrogen can be modeled to exist in different adsorption states (I-III), state I being responsible for the main part of the adsorption [2]. Consequently, the adsorption model (3) is acknowledged to be a simplification of the existing physical-chemical situation. 5. SOLUTION SCHEME AND COMPUTATIONS The parameter estimation was performed by nonlinear regression. The problem set-up leads to a minimisation task including solving of coupled partial differential equations [3]. All the computations were implemented in the MATLAB environment. The solution scheme and the numerical tools used are shown in figure 1. The pde-solver was chosen from the commercial NAG fortran library. This NAG routine d03pef was made runnable in Matlab by utilising the mex-file mechanism. The object function m-file performed the calculation of weighed residual sum between the model output and the experimental output. It also performed minor data handling to complete the solution scheme.

J

Boundary conditions

J

Initial parameters

Solving the pde-system of Ithe model equations

Generation of new Parameter estimates

~

Matlab minimization routine: fmins based on Nelder-Mead algorithm

~AG-fortran library routine: [d03pef based on the method of lines

J

Object function calculation

Experimental data

I

I Optimal parameters

I

Figure 1. The solution scheme and the numerical tools used

6. RESULTS AND CONCLUSIONS Two examples of the model output and the experimental output are given in Figure 2. The estimated parameter values with their sum of squares of residuals for four different temperatures are shown in Table 1.

442 Due to the correlation between parameters the identifiability of the parameters was improved by lumping together the products of kA and Ntot 2 and kD and Ntot 2 and the identifiability was further improved by using simultaneously two dynamically different pulse responses in parameter estimation. The parameter sensitivity was studied with the aid of contour plots. Figure 2 shows that the model is in good agreement with the experimental data. The model is able to describe the adsorption dynamics in each specific temperature for at least two different kinds of pulse responses. However, the single-site Langmuir adsorption model cannot be supported by comparing the obtained parameter values for different temperatures. The number of active sites should remain nearly constant and the values of the rate constants should regularly increase with temperature. Transient methods can give valuable kinetic information, but the performance of the analyser is of key importance. Extra precaution must also be taken to deal with the parameter identifiability. Figure 2. The 8x~o-' model (continuous line) and the experimental data .~+ ij:!i! \~;'..+..~++.+~,~. . . . _..... (circles) at T=90 ~

10

30

20

time/=

40

50

- - -60 . . . . .

~

8 x 10"3

Table 1. Estimated Parameters. oo+ . . . . .

o

T/~

kA

30 90 120 180

2.2e-3 2.9e-3 5.3e-4 5.9e-4

~o

mGmol-2s

3'o

-1

,~+0-+~--- -;'o. . . . . . .

kD

m3mol-ls

"1

;,o . . . . .

Ntot

mol]m3cat

~(Ccal-Cobs)

2

bed

1.2e-3 2.9e-4 1.2e-4 8.6e-5

132 175 388 621

2.5e-6 3.6e-6 3.5e-6 3.9e-6

7. REFERENCES 1. Kanervo, J., Master's Thesis, Helsinki University of Technology, Department of Chemical Technology, Espoo 1998. 2. Smeds, S., Salmi, T., Lindfors, L. P., Krause, O., Appl. Catal. A., 144 (1996) 177. 3. van der Linde, S. C., Nijhuis, T. A., Dekker, F.H.M., Kapteijn, F., Mouljin, J. A., Appl. Catal. A.,151 (1997) 27.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science Ltd. All rights reserved.

443

Modelling the Voltammetric Behaviour of Cobalt Cations Inside Zeolites M.A.N.D.A. Lemos (1), P. Sousa 0), F. Lemos_(2), A.J.L. Pombeiro 0) and F. Ram6a Ribeiro (2) (1)Centro de Quimica Estrutural, Complexo Interdisciplinar, Instituto Superior T6cnico, Av. Rovisco Pais, 1096 Lisboa Codex, Portugal (2)Centro de Engenharia Biol6gica e Quimica, Instituto Superior T6cnico, Av. Rovisco Pais, 1096 Lisboa Codex, Portugal

ABSTRACT In the present study we have applied digital simulation techniques to the investigation of the electrochemical behaviour of a metal cation, cobalt, inside the porous structure of an Y zeolite. The CoNaY zeolite was prepared by ion-exchange and the electrochemical behaviour was studied by cyclic voltammetry using a CoNaY/graphite composite electrode in a specially designed apparatus. 1. INTRODUCTION The use of zeolites in catalysis, adsorption and ion exchange is closely linked to their ability to act as ion-exchangers, due to the presence, in their intracrystalline pore structure, of cations that are electrostactically bound to the inner surface of the framework. Suitable exchange of these cations allows us to change the catalytic properties of a particular zeolite and also to introduce additional catalytic functions. The use of protons is extremely important for their use as acid catalysts, but transition metal cations are also used to introduce specific traits to these catalysts or to modify the acid-base properties of the structures. Electrochemical studies have been used for a long time to study the redox behaviour of species in solution; they can also be used to provide a good insight into the dynamics of these cations inside the pore structure of zeolites and this line of research has attracted some attention lately [ 1]. Digital simulation has been a very useful technique in interpreting the complex patterns that are obtained in dynamic electrochemical experiments, namely in the enlightenment of mechanistic aspects [2], and also to study species entrapped in solid electrodes [3]. Its extension to zeolite modified electrodes could help in this kind of studies. Experimental studies with zeolites containing cobalt complexes relevant in catalysis have already been investigated by others [4, 5].

444 2. EXPERIMENTAL The starting zeolite material, an Y zeolite with a Si/A1 ratio of 2.5 (LZY-52 from Union Carbide), was used in its sodium form. Cobalt was introduced by an ion-exchange procedure: 1 g of zeolite was stirred in 5 ml of a 0.86 M cobalt nitrate solution for about 20 hours. After the exchange the zeolite was filtered off, dried under vacuum and then calcined for 8 hours at 500 ~ in a furnace, using a self-steaming procedure. The final zeolite that was obtained can be represented by the following formula: Co0.26Na0.asA1Si2.507. The electrodes were prepared by mixing 2.1 mg of zeolitic material with an equal amount of graphite. The mixture was homogenised and pressed in a 4 mm press with a total applied pressure of 0.5 ton. The pellet thus produced was placed in a special support where it contacted with the electric circuit by means of a platinum disc that was pressed against the pellet, as shown in figure 1. This assembly was then placed in a conventional electrochemical micro-cell containing 3 ml of a 0.2 M solution of KC1 in water (which was used as the electrolyte solution). Potentials were measured using a silver wire immersed in the electrolyte solution and separated from the main cell compartment by a Luggin tube.

Figure 1 - Schematic representation of the support used to study the electrochemical behaviour of transition metal cations entrapped inside zeolites.

The cyclic voltammetric experiments were carried out using a Radiometer (model DEA 101) digital electrochemical analyser controlled by a computer, which was also used to acquire the data.

445 Simulations were carried out with a commercial spreadsheet (Excel 97 - 9 Microsoft Corp.) in Personal Computer. The Euler method was used to integrate the relevant differential equations. 3. THEORY

Although cations within the zeolite framework have a certain degree of mobility, since each cyclic voltammetry experiment occurs during a relatively narrow time-window, we considered that no significant exchange with cations in solution occurred during the potential sweep. We also assumed that a classical current-potential equation [6] could be applied to obtain the kinetics of the electron-transfer process that occurs within the zeolite framework. As a first approach a uniform field was considered within the whole pellet. Since the results obtained using this approach were quite acceptable by comparison with the experimental results, this item was not further elaborated. The equations that describe the Co2+/~ redox process, which can be assigned [5] to the process observed in our pellet, can, thus, be written as: d

nco2+

j~nF= dt

d nc~176=" k~ ! e-an(E-E~176

. e(1-a)n(E-E~

)

where nco2+and ncoOcorrespond, respectively, to the number of moles of C o 2+ and Co o inside the pellet at any given time. 4. RESULTS AND DISCUSSION Voltammograms of species in constrained spaces present a peculiar shape, quite different from the one presented when in solution; this is due to the fact that the diffusion-reaction problem does not have a semiinfinite boundary condition but, on the contrary, it is restricted to a limited amount of space. Figure 2 shows a voltammogram obtained for a complete cycle on a CoNaY/graphite pellet. The results indicate that the cobalt cations are in the divalent state [5]. When the

y j/

f,_,

/

J

A~j/ -2000 -~500 -~000 -500

0

500

1000 1500

E (mV) Figure 2 - Voltammogram obtained at 20 mV/s scan rate from a CoNaY/graphite pellet (see text for other conditions).

446 potential is decreased the cations 3.5 are reduced to the metallic state, 3 with a wave which is extensively 2.5 superimposed on the reduction of 2 the electrolyte solution itself. ~1.5 Upon inversion of the potential -9 1 sweep, the reduced atoms are 0.5again re-oxidised, showing a shape that is typical of redox -0.5 reactions occurring in a -1000 -500 0 500 1000 constrained space. E (mV) Since the reduction wave is extensively superimposed on the wave observed on a electrolyte solution reduction, no Figure 3 - Oxidation CoNaY/graphite pellet after the reduction cycle attempt was made to simulate this section. Thus, only the oxidation (experimental data: ~ 20mV/s, A 10 mV/s; was simulated, using the simulations: m). assumptions presented above and taken into account the balance between the two species (Co 2+ and Co ~ at the beginning of the section of the voltammogram being simulated. Figure 3 compares simulations with experimental results at two different scan rates. The electrochemical parameters obtained are: E ~ - -720 mV (vs. s.c.e.); k ~ = 1.3 x 103 sl; an = 0.1.

J

5. CONCLUSIONS The use of dynamic electrochemical techniques in general, and cyclic voltammetry in particular, may prove to be very helpful techniques in the characterisation of electrochemically active species inside zeolite cavities. By resorting to digital simulation, the electrochemical parameters can be estimated and used to quantify the redox properties of these species. This may also open the possibility of relating the catalytic properties of metal cation-loaded zeolites with the redox properties of these cations.

6. REFERENCES 1. J.-W. Li, K. Pfanner, G. Calzaferri, J. Phys. Chem., 99 (1995) 2119; D.R. Rolison, Chem. Rev., 90 (1990) 867. 2. M.A.N.D.A. Lemos, A.J.L. Pombeiro, J. Organomet. Chem., 438 (1992) 159. 3. P.J. Peerce, A.J. Bard, J. Electroanal. Chem., 114 (1980) 89. 4. K. Balkus Jr., A.G. Gabrielov, S.L. Bell, F. Bedoui, L. Rou6, J. Devynck, Inorg. Chem., 33 (1994) 67 5. F. Bedioui, E. De Boysson, J. Devynck, K.J. Balkus Jr., J. Chem. Soc. Faraday Trans., 87(24) (1991 ) 3831. 6. A.J. Bard, L.R. Faulkner, Electrochemical Methods, Wiley, New York, 1980.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

447

Modelling Transient Tracer Studies o f C o m p l e x Activation M e c h a n i s m s o f T w o - A t o m Labelled Molecules. A.A. Shestova'b, R. Burch b, J.A. Sullivanb and V.S. Muzykantov a a Boreskov Institute of Catalysis, RAS, Pr Akademika Lavrentieva 5, Novosibirsk, 630090, Russia. b Catalysis Research Centre, Chemistry Department, University of Reading, Whiteknights, Reading, RG6 6AD, UK. Abstract

The work reported here concerns the investigation of several different complex mechanisms of activation of molecules containing two atoms of a labelled element in open systems (Plug Flow Reactors, Continuously Stirred Tank Reactors) and the relationship between kinetic exchange parameters and the parameters of elementary steps. General sets of kinetic equations are used to simulate the responses of different types of reactor and to illustrate the special characteristics of these systems.

1. INTRODUCTION The experimental investigation of isotopic exchange of doubly atomically labelled molecules RA2 (e.g. 02, N2, H2, CO2, NO2 etc.) has been examined extensively. In molecules represented as RA2 where A represents a labelled chemical element which has two isotopes (A and *A) and R represents a part of the molecule which does not contain the A-element. The basics of the kinetics of isotopic exchange for these systems have been derived [1-4]. However in real catalytic systems exchange is most likely to occur through complex mechanisms consisting of several steps. In this case a new peculiarity arises and it is necessary to consider the overall variations of isotopic variables due to all reactions involving the components. This results in kinetic exchange parameters which will depend on the nature of the intermediates and on the relationships between the rates of the various single steps and hence on the specific conditions of the reaction. However, the investigation of complex multi-step mechanisms systematically (as distinct from the use of compartmental modelling) both in general cases and in open systems, has not yet been conducted. Such investigations are important for the further development of methods of Transient Kinetics e.g. in the Steady State Isotopic Transient Kinetic Analysis (SSITKA), Isotopic Exchange Transient Kinetic Analysis (IETKA) and Temporal Analysis of Products (TAP) techniques.

2. RESULTS AND DISCUSSION To obtain a general set of kinetic equations for isotopic redistribution in open systems, in a system containing gaseous RA2 molecules and surface monatomic (A) species (where usually these are lattice atoms of a catalyst e.g. lattice oxygen atoms in an oxide catalyst),

448 three types of exchange mechanism were considered [2,4]. These are differentiated by the number of exchanging atoms per elementary exchange act through which gaseous RA2 molecules might exchange their atoms in the presence of surface (A) species. These are as follows; 1. without participation of surface (A) species, 2. with participation of one surface (A) species and 3. with participation of two surface (A) species. Kinetic equations describing isotopic redistribution in systems containing gaseous RA2 and surface (A) species for a plug flow reactor are:

N Oa 1 Oa ~ +- N = e.R(as - a ) Ot r Og

(1)

N Oz 1 N O Z --+- -eXz + eJf3(a - a) 2 Ot r Og

(2)

das

U, dt

(3)

- 2R(a s - a)

with initial conditions: t=0, a(0,s ) = ot~ z(0,g )=z ~ cts(0, g ) = a f, V g ~ [0,1]. and boundary conditions: g = 0, a(t,0) = FI~P(t), z(t,o)= F2mp(t), V t > 0. where --~0and c3 represent partial differentiation with respect to time (t) and the Ot

Og

dimensionless length of catalyst bed (g). N and his are the concentrations of gaseous RA2 and surface (A) species; R, K and 1s are the rates of hetero-, homoexchange and exchange of the third type. a and ot~ represent the atomic fraction of the heavy isotope (*A) in the gaseous RA2 and the surface (A) species respectively, z is the difference between the fraction of isotopically doubly labelled molecules R'A2 and their isotopic equilibrium value (x2eq=a2) V m z2=(x2-a2); x = - - , gas phase residence time within the reactor, e = where m = mass of v ~' catalyst, V = volume of reactor and v = gas flow rate. The general systems of kinetic equations for a CSTR are not presented due to limitations on space in this article but will be published later [7]. Homoexchange processes are those where the isotopic redistribution takes place within isotopic molecules of the gas phase, RA2. Heteroexchange processes are those in which there is a change in the fraction of heavy isotopes in the gas phase (RA2) and on the surface (A). The rates of homo- and heteroexchanges (K and R) are related to the rates of the three types of exchange by the equations: R = O.5K2 + K3 and

(4)

K = KI + K2 + K3

where Ki is the rate of i-type exchange. ,,.,

List of mechanisms considered .............................. N Reaction Step I 1 A2+2z2zA 1 2 A2+z+()<x>zA+(A) 0 3 A2+2()c:>2(A) 0 4 A2+zc~>zA2 0 5 zA2+zc:>2zA 0 6 zA+()c~>z+(A) 0

H 0 1 0 0 0 0 i

,,

Mechanism IV V 1 0 0 0 0 0 0 1 0 1 1 0

HI 0 0 1 0 0 0 i

,

i

i

VI 0 1 0 0 0 1

VII 0 0 0 1 1 1

VIII 1 1 1 0 0 1 ,,,,

449 Kinetic analysis of isotopic redistribution has been performed for a series of one step and complex mechanisms with different intermediates and different numbers of steps. The various mechanisms examined are represented in the table presented above. In the above table values of 1 and 0 indicate the participation, or lack of participation, respectively, of the relevant step in the overall exchange process for a given mechanism. The molecule RA2 is represented as A2 for the sake of simplicity. The symbols z and 0 identify adsorption sites and surface or lattice vacancies, respectively. In other words zA(zA2) and (A) can be taken to indicate weakly and strongly bound surface atoms, respectively. Low and high concentrations of zA (zA2) and (A) species will be assumed. Mechanisms I and V and those that involve associative adsorption of RA2 with adsorbed three or four atomic intermediates, e . g . A2+zAzA3 can provide examples of the first type of exchange [3,5]. Mechanism II and Eley-Rideal mechanisms involving exchange via triangular intermediates ( A A * A ) a d s , with the participation of one surface atom (A) give the second type of exchange [3,5]. Finally mechanisms III, IV (Bonhoeffer and Farkas), VI, VII and VIII (all with fast reaction 6) and associative mechanisms via four-atomic complex (AAA*A*)ad.~ with the participation of two (A) species [6] may supply the third type of exchange mechanism. Under steady state conditions the rates of the forward reactions and of the reverse reactions are the same for every step. When the switch RA2 ::> R'A2 is carried out, the redistribution response of isotopically labelled molecules R'A2 will be observed. It has been established that kinetic equations in isotopic variables for both types of reactors have a general form and that this general form is independent of the type and complexity of the reaction mechanism. Kinetic exchange parameters, R, K and K3, are generally complex functions of the rate of steps and are determined by the mechanisms of the reaction leading to exchange; to be more precise - by the character of the steps and their rate. For example in one case for the Bonhoeffer-Farkas Mechanism IV R = PlOP

K3 = p 1 (1)2

K = p]

(5)

where ~p =

P6 - a dimensionless parameter and/~ is the rate of step i (see (2,ol + P6) table). If/96 > >,ol, then ~ = 1 and K = R = K3 =/91 - the third type of exchange. The ratio of the rates of heteroexchange to homoexchange (B=R/K) is an important parameter. It is shown that in the ease where only one surface (A) atom is involved in the elementary heteroexchange act, then B = 0.5. If, along with this mechanism, homoexchange also proceeds (in any mechanism) then B < 0.5. If both "A" atoms of a doubly labelled molecule are involved in the elementary act of the heteroexchange with the catalyst then B = 1. It is shown that the mean number of surface (A) atoms which participate in one act of exchange, i . e . the Coefficient of Multiplicity, M = 2oB [7]. When considering mechanism IV M = 2o~. [

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

1

.

,., 0.8 ~z. 0.6

D.

~j_ 0.6 ><- 0.4

x:o.4

0.2

0.2 0

0.8

|

|

,I-

,~-

-v.

9

0

1

2

3

4

5

Ome/a.u.

0 0

1

2

3

4

5

time / a.u.

Fig 1. Simulated fraction of isotopic molecules RA2.i*Ai (X~ in the gas phase as a function of time following the ideal step switch RAe -~R*Aefor (a) the 2~ type of exchange and (b) the 3rdtype of exchange.

450 1.1 ~

1.1

0.9

0.9

0.7

0.7

N T /~ ~0t'step-exp ~-0.5 ~in-Pe~ 0.3 ~] /~/ / ~ e~::mPe~

p

=p

2

3

N .~j~,._~ep. ~.~r~f'-st "'~Peotp -e~P ~" 0.5 ~ Z e ~ -

~

p

-

0.3

0.1

0.1

-0.1 0

1

2

3

time / a.u.

4

5

-0.1

- .....

,.~----, [h [ I ~ I

1

4

5

time / a.u.

Fig 2. Simulated response curves of a and z variables obtained during isotopic exchange transient kinetic analysis for the 2~ (a) and 3rd (b) type of exchange with different boundary conditions. Open data points represent an ideal step change while filled data points represent a simulated experimental step change

We have modelled all types of exchange mechanisms using different initial and boundary conditions and different sets of parameters. Some of the results are shown in Figs. 1-2. Fig. 1 illustrates typical dependence of the fraction of isotopic molecules RA2 at the outlet of a reactor as a function of time after an ideal step switch RA2 --) R'A2 for the 2nd and 3rd type of exchange. The simulation was conducted using a small residence time (x). Initial redistribution of the isotopic product at time t=x shows that for the second type the main isotopic product is RA*A and for the third type the main isotopic product is R'A2. The same data (using an ideal step switch) are shown in the unfilled data points in Fig. 2 (for the ot and z coordinate). As seen, z(t)=0 at all times of response for the 2nd type of exchange. This indicates that the mixture of isotopic products is at isotopic equilibrium. This differs from the data seen representing the third type of switch (Fig 2b) where the isotopic molecules (RA2.i*Ai) are not statistically mixed. Figure 2 (filled data points) also illustrates the same data for a simulated "experimental" stepchange of RA2 to R'A2. It can be seen that statistical equilibrium is not present at all times, even for the 2nd type of exchange. In these cases, for the purpose of discrimination of exchange mechanisms we need to conduct the switch in a more precise manner. This requirement is generally true for transient kinetic experiments as the beginning of the transient trace contains the most valuable information concerning the reaction mechanism on the catalyst surface. REFERENCES

.

4. 5.

K. Klier, J. Novakova and P. Jiru, J. Catal., 2 (1963) 479-484. V.S. Muzykantov, V.V. Popovskii and G.K. Boreskov, Kinet. Catal. 5 N4, (1964) 624-629. G.K. Boreskov, V.S. Muzykantov, Ann. N.Y. Acad. Sci. 213 (1973), 137-160. V.S. Muzykantov, React. Kin. Catal. Lett. 33 (1987) 937-947. G.K. Boreskowin "Catalysis, Science and Technology", Springer-Verlag, New York, v3, (1982), p. 39. E.R.S. Winter, J. Chem. Sot., (A), (1968) 2889-2902. Shestov, A.A., ~ h , R., Sullivan, J.A. and Muzykantov, V.S., in prep.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 1999 Elsevier Science B.V.

451

Preparation of ZSM-5 zeolite film on metal suport A. Brehm, U. Antons, A. Bekurdts University of Oldenburg, Germany Introduction

In the last years we studied a number of liquid phase catalytic reactions with the aim to influence the kinetic and mass transfer parameters. Reactions of interests were: - the isomerisation of long-chain n-alkanes

- the hydrogenation of 1-alkenes (C 8 - C20 ) - the hydrogenation of the methylesters of vegetable oils - the catalytic removal of dissolved oxygen from water Small catalyst particles (normally small zeolite or active components on zeolitic carriers) were fluidized in the liquid phase by agitation. In liquid systems the point of incipient fluidization is near to the point of excessive entrainment. This results in a small opportunity for the variation of the actual velocity between Particles and the liquid phase. To obtain a fluidized bed, the hydrodynamic conditions can be varied only in a small range. The parameters for an external mass transfer limitation can be extensively influenced in a fixed bed. The relatively dense packing and the resulting higher pressure drop is sometimes desired to provide a more even distribution of the liquid flow. To lower the permanent pressure drop across a fixed bed, the catalytic particles may be shaped in a form to have increased void volume. To improve the hydrodynamic aspects (heat and mass transfer, pressure drop and uniformity of distribution of concentration, temperature and velocity) the Sulzer Chemtech Ltd. (Winterthur, Switzerland) developed open crossflow channel catalysts and catalyst supports based on the well known structured packing/1/. The preparation of thin films of molecular sieves (MFI-types) on various supports (especially metallic substrates) is described by Jansen, Nugroho and van Bekkum. Calis et al. /3/ prepared such films on wire (stainless steel) gauze. In our lab ZSM-5 films have been produced by the in situ crystallization on structured packings as supports. E x p e r i m e n t a l and C h a r a c t e r i z a t i o n of Z S M - 5 zeolite films

The hydrothermal synthesis of ZSM-5 zeolite films (in this work) is based on the preliminary studies on the alkaline-free synthesis of large crystals of ZSM-5 by Mtiller and Unger/4/. A mixture of the following composition was applied: 9 TPABr + 182 (NH4)20 + AI203 + 30 SiO 2 + 2460 H20

The wire supports (for example Sulzer-structured packings) were added to the hydrogel.

452 To check the correct zeolite structure we used the X-ray diffraction. Micrographs and scanning electron micrographs (SEM) of the samples covered with zeolite film give information about the crystallization, especially the habit, the size, the orientation and the degree of coverage. Figure 1 and 2 clearly show a support of more than 90 % covered with a continuous crystal film. The degree of coverage (increase in sample weight is caused the zeolite film on the supports) was determined after intensive cleaning with a high-pressure water jet (flow" 300 dm3/h, nozzle diameter: 4 mm, distance between nozzle and samples: 2 cm) and defined drying conditions. The stability of the crystal on the support material was tested by a mechanical scratch with a sharp spatula. The chemical stability against formic acid (pH - 1.5) was proved during a period of 5 as well as 48 hours. To check the distribution of the catalytic activ compound ( palladium as well as platinium ) we used EDX-spectroskopie. Results

The preparation and the test of ZSM-5 zeolite film has been studied on various supports. In this presentation we present the results of the use of wire (stainless steel) gauze only. Other supports we used were copper and different kinds of glasses. Using wire gauze we varied (apart from the typical parameters like temperature, reaction time, ...) a) the type of stainless steel (5 typical materials of Sulzer-structured packings), b) the form of the support (for example Sulzer-structured packings), c) the surface structure, influence of corrosive action by HC1. ad a) All analysis, especially the micrographs (SEM) and the analysis by weighting, show an excellent regular zeolite film with as many crystal faces as possible exposed. The crystals show a minimum intergrowth. There is no difference between the 5 types of stainless steel. ad b) Using Sulzer-stmcmred packings we analyzed the regularity of the zeolite film as a function of the packing position. Figure 1 shows the micro graphs (SEM) for 4 typical positions (inside - center, bottom, periphery - inner side and outside). ad c) clearly shows pronounced intergrowths of the crystals in the zeolite film. A corrosive action by HC1 can be very promising with regard to systems with an insufficient wettability and special adsorption processes. ad b) The experiment clearly shows pronounced intergrowths of the crystals in the zeolite film ( Figure 2 ). A corrosive action by HC1 can be very promising with regard to systems with an insufficient wettability and special adsorption processes.

453 In differnt reactors ( external as well as internal loop ) the activity of the catalytic packings we studied with the aim to influence the kinetic and mass transfer parameters. Reactions of interests were: - the catalytic removal of dissolved oxygen from water - the hydrogenation of 1-tetradecene References

/1/ ,,Katalysatoren ftir heterogene Systeme", J.-P. Stringaro, J. Luder, Chemie-anlagen + verfahren, 4, 1992 /2/,,Controlled growth of thin films of molecular sieves on various supports", J.C. Jansen, W. Nugroho, H. van Bekkum, Proc. 9th Int. Zeolite Conf., Montreal, S. 247, 1992 /3/ ,,Anwendung von Zeolithen bei der selektiven katalytischen Reduktion von NOx in Industrieabgasen", H.P. Calis, O.L. Oudshoorn, A.W. Gerritson, K.J.C. Jansen, C.M. van den Bleek, H. van Bekkum, Chem.-lng.-Tech. 67, S.777, 1995 /4/,,Preliminary studies on the synthesis of alkaline-free large crystals of ZSM-5", U. Mtiller, K.K. Unger, ZEOLITES 8, S. 154, 1988 For further informations please contact Priv.-Doz. Dr. Axel Brehm, University of Oldenburg, FB 9 - Department of Industrial Chemistry, P.O.Box 2503, D 26111 Oldenburg; Phone: +49 441 798 3841; Fax: +49 441 798 3330 Email: [email protected]

This Page Intentionally Left Blank

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

455

A Steady State Isotopic Transient Kinetic Analysis o f N O Reduction over Pt / SiO2 under Lean Burn conditions. A.A. Shestova'b, R. Burch b and J.A. Sullivan b Boreskov Institute of Catalysis, RAS, Pr. Akademika Lavrentieva, 5, Novosibirsk, 630090, Russia. b Catalysis Research Centre, Chemistry Department, University of Reading, Whiteknights, Reading, RG6 6AD, UK. a

1.

INTRODUCTION Steady State Isotopic ;Fransient Kinetic Analysis techniques have been used, in conjunction with Mathematical Modelling, in an attempt to clarify the mechanism of the NO/C3HdO2 reaction over supported Pt catalysts. The isotopic switches carried out under steady state reaction involved replacing NO in an NO/O2/C3H6 stream with an equivalent volume and pressure of labelled 15NO. Profiles of reactants and products after the switch, along with the profile of an inert tracer gas (At), are monitored as a function of time. The profiles of the unlabelled products of the reaction, N2, N20 and NO2 yield information about the surface coverages of reactive intermediates on the catalyst surface during the steady state reaction and can be used to analyse the fraction of the surface involved in reaction. The profiles can also be used to estimate the time constant for the production of the product molecules and thus also gives information about the reactivity of the adsorbed species. From the corresponding experimental data the fraction of heavy (15N) isotope in the products (N2 and N20) was calculated; (at (N20) can be calculated in a similar way).

a(N~) =

0.5['~N'N] +['N2 ] [ ' % ] +['W"N] +["N2 ]

During and after a switch the masses 28 (14N2), 29(14N15N), 40(At), 45(~4N15NO) and 46(~5N20) are monitored as a function of time. The system is left in ~5NO for 2 minutes and then the 15NO is replaced by 14NO. The upward trend for the signals at 28 (14N2) and the downward trend for m/e = 46 (~5N20) are then taken to equal the profiles expected for 1~N2 and 14N20 expected during the original NO to ~sNO switch. These species cannot be directly monitored due to overlaps with 14NO and CO2 respectively. The balanced nature of the switch is checked against the expected 15N2 and 14N20 switches (since the sum of the overall normalised N2 and N20 concentrations should equal 1). Switches were carried out at three different temperatures; 225, 263 and 328 ~ over 100 mg of l%Pt SiO2 under the following conditions, 1.63% NO, 8900 ppm C3I-I6, 8.8% 02 in a total flow of 111 cm3 min1. The catalysts were held at a temperature of 480 ~ for 1 h prior to reaction and held at the corresponding reaction temperature for 30 min prior to the switches taking place.

456 2.

RESULTS AND DISCUSSION

Mathematical modelling of the data followed the methodology outlined below. At first, for a qualitative analysis of the experimental data, a number of mechanisms following the general scheme below

have been studied on the basis of continuity equations for isotopically labelled components. Here A represents an initial substance (reactant), B represents a final substance (product) and $1 and $2 represent two pools of surface species. This scheme includes all possible variations (consecutive, parallel, buffer, impact and direct as well as their combinations) in the production of B from A from two intermediate pools. The basis for modelling such systems has been previously derived [1-3]. More complex cases, with more than two intermediates will generally be resultants of these simple cases. The qualitative dependence of the fraction of labelled atoms in the gas phase with time ct = c~(t) gives some important information about the mechanisms of redistribution of isotopic molecules of product B in terms of compartmental modelling. Every mechanism is characterised by the number and nature of intermediates present as well as a definite interdependence between its kinetic parameters. This will be reflected in the types of the isotopic responses of both product and reactant molecules. For example simulated results for the buffer mechanism (a) and the one intermediate mechanism (b) with 50% reversibility of step 1 are shown in figure 1. A<

1 ~) S l

$2

(a)

J~2

~/ - B

A

~

1

~

~

S1

(b)

We can see that convex shapes in the semi-logarithmic co-ordinate (figure lb) result for the expected profiles of reactant and product in the case of the buffer mechanism and that straight line profiles result in the case of reactant and product in the case of the one intermediate mechanism. This is true even with a reversible first step in the reaction. The profile of the unreacted reactant A, or more precisely the deviation of t~A from the inert response can be used to give information regarding the nature of adsorption of A and to determine the extent of reversibility of the adsorption/desorption process (see Fig.2). Another feature of the response of the reactant A which has heretofore been ignored is that information can be determined regarding the reaction network in the case of reversible adsorption of reactant A. For

457 ,

I

,

,|

,

9

,,~,

4.5

0.8

_. 4 . 0 3.5

90.6 =9

-~

2.o

~ 2.5

0.4

=oo,~ ........ 0

2

4

,, I 6

8

" ;o'.5

10

0 0"-0

time / a.u.

2

4

6

8

10

time / a.u.

Fig 1. Simulated response curves of a variables for both reactant (.4) and product (11) (fig 1 (a)) and a logarithmic presentation for the buffer mechanism and one pool mechanism (fig l(b)) following an n

ideal step isotopic switch from A + A ,,tabmed .... tam --) *A + A ,,tabetted.... ~.t~ 1-7 Reactant A (Buffer mechanism); O, Product B (Buffer mechanism); A, Reactant A (one intermediate mechanism); 0 Product B (one intermediate mechanism); x represents the ideal response o f an inert tracer and the Reactant A response when the desorption rate is zero. The simulation was conducted using a small residence time.

example for the buffer mechanism when the rate of step 3 is zero (or much faster than the rate of steps 1 and 2) then the response for (Za in logarithmic co-ordinate must be a straight line. Thus the Ota profile may reflect the presence of different intermediates and show their "position" in the reaction network. If there is more than one labelled product from the reaction then the order of responses of each of the products, orBs(t), can yield a useful criterion for the discrimination of different mechanisms. It is useful in these situations to establish which product is isotopically first. By the "isotopically first" product we mean the product which has the lowest residence time on the catalyst relative to all the other reaction products. If we consider a reactant A(g) and two products Bl(g) and B2(g), then generally after an isotopic switch the fraction of heavy atoms in the gas phase of the isotopically first molecule (BI) will be higher than that for the isotopically second (or last) molecule at all times, i.e. or(B1) > or(B2). This is represented graphically in Figure 3 where it is seen that cc (N20) > ot (N2), i.e. that the area between the ct (N20) profile and the Ar profile is less than the area between the ct (N2) profile and the Ar profile. This can be displayed in either of the proposed linear networked reaction schemes presented below where A is the reactant and B] and B2 represent the isotopically first and isotopically second product molecules, respectively.

A ~> SI<:=)... Si ~ ~ > ? B1

...<~> . . . S k ~

.? B2 (i <j _
In these eases the production of the last product, B2, normally involves the terminal (S,) intermediate. "Isotopically first" molecules might also be produced by readsorption of the B1 molecule followed by its transformation on the surface to form the B2 molecule through additional pool(s) of intermediates. All the experiments we have carried out have consistently shown that, regarding the a(t) functions, N20 is a l w a y s the isotopically first product (Fig 3). In

458

i~o00: I ~

o

.

~

I ~

3

o

Time / s

Fig 2. Comparison o f ~sNO traces seen following the 14N0/C~o/02 -) 15N0/C~tIc,/O2 switch over Si02 (A) at room temperature and over a 1%Pt/Si02 catalysts at 263 ~ (0) and at 328 ~ (L~. A profile for the inverted normalised Ar trace is also shown as a full line.

=,ot

t I

o . ~ - " ' 14o' lio " 1~o ' 1,1o ' 1so Time / s

Fig 3. a (N9 and a(NeO) as a function o f time following the ~4NO/C~-Io/02 "-) 25N0/C~c,/02 switch over 1% Pt/SiO: at 225 ~ E7 a (Nz); O, a(N:~9). An invertedAr profile is also shown as a full line

terms of the reaction mechanism this means that N20 must be produced from surface intermediates formed earlier, whereas N2 is formed from later intermediates. In our present case the observed difference between the ~t (N20)(t) and a (N2)(t) profiles might also indicate the presence of intermediate pool(s) of relatively large concentration. Another reason for this might be that the production of N20 follows the interaction between terminal and earlier precursors whereas the production of N2 takes place exclusively through terminal precursors. Another important criterion for qualitative analysis is the degree of statistical equilibrium, as a function of time, of the isotopic molecules in the products. This gives information both about the rate of independent homoexchange within the products and the nature of the final surface intermediates. In addition we suggest some criterion functions both for testing the validity of, and for qualitative analysis of, experimental SSITK data [4]. Our experimental SSITK results indicate that the parallel and impact schemes can be excluded and that we must account for some consecutive-buffer scheme. Further details of proposed reaction schemes and experimental methods are available elsewhere [5]. The full set of partial differential equations coupled with ordinary differential equations are integrated using the numerical Method of Lines and a Backward Differentiation Formula Method. Estimation of the Kinetic Parameters was obtained by minimisation of the objective function of the sum of squares of deviations for both reactant (*NO) and products (*N2 and *N20) simultaneously. The minimisation was achieved with Simplex and Marquardt Methods. References.

.

3. 4. 5.

Happel, J., "Isotopic Assessment of Heterogeneous Catalysts" Academic Press, Orlando, FI (1986). Happel, J., Walter, E. and LecourtierY., J. Catal., 123, 12, (1990). Shannon, S.L. and Goodwin, J.G., Chem Rev., 95, 677, (1995). Shestov, A.A., Burch, R. and Sullivan, J.A., in prep. Burch R., Shestov, A.A. and Sullivan, J.A., submitted to J. Catal.

Reaction Kinetics and the Developmentof Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

459

Modelling the Dynamics of the Surface of a Carbon C. Palma ~ I. Santos Silva(t), F. Lemos(2), L. Sousa Lobo~ (1)Departamento de Quimica, Centro de Quimica Fina e Biotecnologia, Faculdade de Ci~ncias e Tecnologia, Universidade Nova de Lisboa, Quinta da Torre, 2825 Monte da Caparica, Portugal. (2)Centro de Engenharia Biol6gica e Quimica, Instituto Superior Trcnico, Av. Rovisco Pais, 1096 Lisboa Codex, Portugal

Abstract In this work we aim to present a preliminary description of the use of a numerical deconvolution technique to analyse the temperature programmed decomposition of the surface of a carbon material. The evolution of CO and CO2 from the carbon surface has been monitored by a mass spectrometer and the resulting signal was numerically processed so as to give kinetic information on the individual processes that occur on the surface, assuming that these can be described as irreversible first-order kinetic processes obeying an Arrhenius law. 1. INTRODUCTION Carbon materials have a significant importance in many fields of catalysis [ 1]. Not only they are important reactants, but also they have attracted a lot of attention as potential supports for other catalysts [ 1-2]. In particular, the analysis of their potential use as supports for environmental catalysis, such as NO and CO conversions, has been increasing lately [ 1-2]. Carbon surfaces contain many oxygenated groups, like carbonyl, carboxyl, phenol, etc, which, when heated, decompose and release CO or CO2; their highly reactive surface produces very intimate interactions with reactants and products, and is also likely to change during the course of reaction. Surface characterisation of carbon materials, and in particular its chemical dynamics is extremely important and can be performed by several methods, such as TPD. TPD is extremely useful in the characterisation of the surface of carbon materials and has been extensively used to characterise surface functionality of modified carbons [3-4]. Analysis of the oxygen functionality has been performed by a wide range of methods [1-4]. In the present work, we have applied digital deconvolution techniques to inspect the release of CO and CO2 from the surface of a carbon material. This procedure, which has been successfully applied to the study of the acidity in zeolite samples [5], will allow a more quantitative description of the sites that are present in the surface of the carbon.

460 2. METHODOLOGY TPD can be used to characterise the composition of the carbon, by following the amount of CO and CO2 released as a function of temperature. The observed thermogram can be rather complex, due to the superposition of species released from surface sites of different chemical nature and/or with different chemical environment. This complexity makes it difficult to interpret the results that are obtained. Numerical deconvolution can be used to decompose the thermogram into the various peaks that it includes. In the following, all decomposition processes will be considered as being of first order and as obeying an Arrhenius law in relation to the change of rate with temperature. In addition, we considered that all processes are irreversible and that no re-adsorption occurs. Due to the nature of these reactions, which correspond to decomposition of surface species with the evolution of a gaseous compound it is likely that these two assumptions will hold. With the set of assumptions listed above, the amount of CO that is released (per unit weight) by groups of a particular chemical nature can be computed as r~O

d Ci i - E ~/RT = dt =k~ e a Ci

(1)

where Ci represents the concentration of surface sites of chemical nature i, while k6 and E~ are, respectively, the frequency factor and activation energy for that reaction. The frequency factor and the activation energy values will be characteristic of the particular group that is being decomposed and can be used as a more reliable parameter to identify these chemical groups than the maximum decomposition temperature. If one considers that all the processes occur independently, the total amount of CO that is released from a unit mass of carbon at any given time, will then be the summation, for all the types of groups that liberate CO upon decomposition, of their respective rates of decomposition

~/Rr

rco evolution = E k6 e -E'a

i=1

Ci

(2)

Although this last assumption may prove to be somewhat restrictive, since it is quite possible that the decomposition of one type of groups may influence the remaining groups on the surface, the results obtained show that it is able to describe, in a general way, the results obtained. Equation 2 can be used to estimate the kinetic parameters of the various groups that decompose within the temperature range of any particular experiment, by fitting it to experimental data with i varying from 1 to n, where n corresponds to the number of sites that are taken to be different; the value of n itself must be considered as a parameter in the fitting, since it is generally not possible to know beforehand how many sites of different chemical nature or of different chemical environment will be present in any given surface. The fitting procedure will involve the solution of n differential equations, one for each of the different surface species that is considered, with an initial value condition that Ci is equal

461 to Cio, the total amount of that species on the surface in the initial sample. These Cio values also correspond to adjustable parameters to be estimated in the fitting procedure. If one resorts to a simple method to solve the differential equations, such as the Euler method with a sufficiently small step, this can be carried out using general-purpose software, such as a spreadsheet. The procedure described here in relation to CO applies equally well to the analysis of CO2 evolution. 3. EXPERIMENTAL A commercially available activated carbon (BDH 33033) was used, in the form of a powder, for this study. The apparatus used for TPD experiments consists of a U-tube quartz microreactor. The reactor exit was connected via a capillary tube to a mass spectrometer (Fisons MD800) set up for continuous analysis of the gases evolved in MID (multiple ion detection) mode. For standard TPD experiments, typically 50 mg of sample were dried in a flow of helium at 110 ~ prior to heating. After treatment the sample was heated in a flow (360 ml/min (STP)) of helium. Subsequently the temperature was linearly increased up to 1000 ~ at a rate of 10 ~ 4. RESULTS AND DISCUSSION The numerical decompositions were carried out using a commercial spreadsheet (Excel 97 for Windows- 9 Microsoft Corp.), in a personal computer. Both the temperature programmed desorption thermograms corresponding to CO and to CO2 were processed. The quality of the fittings that were obtained can be seen in figure 1, while in tables 1 and 2 the values estimated for the various parameters are presented, both for CO and for CO2. Alongside with the assignment of these peaks to individual processes. (a)

(b)

A.

.

~ ~ -

0

200

400

T (~

600

800

1000

i

0

_~

200

....~ i

~

400

600

Fitting

800

1000

T (~

Figure 1 - Profile obtained experimentally for the evolution of CO (a) and CO2 (b) from a sample of BDH 33033 under temperature programmed desorption conditions (heating rate = 10 ~ Symbols correspond to experimental data (not all available data shown) and lines to the fitting obtained by the method described in the text.

462 Table 1 - Peak assignments for the individual processes obtained by the deconvolution of the CO and CO2 evolution temperature programmed desorption thermogram. Peak

Factor (s-1)

CO#1

Initial Conc. (mmol g-l) 0.008

Frequency 1.1 • 103

Activation Energy (kcal mo1-1) 12.2

CO #2

0.023

3.9 • 105

25.5

Lactone or lactone-like moieties [6] Carboxylic acid anhydride

CO #3

0.370

3.2 • 103

22.8

Carbonyl and/or quinone

co2 #~

0.0t0

1.8

t1.0

Carboxylic acid (strong)

CO 2 #2 CO 2 #3

0.015 0.026

3.2 • 102 1.7 x 101

9.6 9.1

Carboxylic acid (weak) [7] Carboxylic acid anhydride

C02 #4

0.010

2.2x 107

40.5

Reaction CO + C(O) ~ C~ + CO2 [3]

• t0 4

Assignment

[4]

[7] [7] [4]

5. CONCLUSIONS From the results that have been presented, it is clear that the decomposition of a single temperature programmed desorption thermogram can provide useful information on the individual processes that occur on the carbon material during the experiment. This information can be useful not only for the chemical characterisation of the functional groups that are present on the surface of the carbon material, but also to ascertain their reactivity and, thus, to improve the understanding of catalytic reactions carried-out over these materials. The power of modem computers allows us to perform this kind of analysis on a routine basis on a common desktop computer without need for special knowledge on programming. However, to have a more accurate description of the surface, additional work will have to identify processes, which may not abide by the assumptions that were used and which may interfere with each other. 6. REFERENCES

1. L. Radovic and F. Rodriguez-Reinoso, Chemistry and Physics of Carbon, Ed. Peter Thower Marcel Dekker, Vol. 25 (1997) 243, NewYork 2. J. Illn-Gomez M., A. Linares-Solano and C. Salinas-Martinez de Leccea, Energy and Fuels, 9 (1995) 976 3. P. Hall and J. Calo, Energy & Fuels, 3 (1989) 370. 4. Q. Zhuang, T. Kyotani and A. Tomita, Energy & Fuel, 8 (1994) 714. 5. C. Costa, J.M. Lopes, F. Lemos and F. Ramba Ribeiro, Catal. Lett., (1997) 44, 255. 6. D. Fagan and T. Kuwana, Anal. Chem., 61 (1989) 1017. 7. V. Zielke, K.J. Huttinger and W.P. Hoffman, Carbon, 34 (1996) 983.

Reaction Kinetics and the Development of Catalytic Processes G.F. Froment and K.C. Waugh (Editors) 9 1999 Elsevier Science B.V. All rights reserved.

463

I n v e s t i g a t i o n o f I s o m e r i z a t i o n K i n e t i c s of m-Xylene Over Zeolite Based Catalysts O. Akpolat*and G. Gi~ndi~z* *Ege University, Faculty of E n g i n e e r i n g , Chemical E n g i n e e r i n g D e p a r t m e n t , 35100 Bornova Izmir, Ti~rkiye. Introduction; In literature, much attention has been focused on the shape selective catalysts, especially on ZSM-5 type zeolites, for the isomerization of m-xylene (1,2,3). The aim of this study is to find a novel catalyst which is selective for pxylene in the isomerization of m-xylene, to investigate the reaction mechanism and kinetics over the most selective catalyst using by a fixed bed catalytic reactor. In this work, gas-phase isomerization reaction of m-xylene was studied over various catalysts including Ni, Sn, Pt, Ga, Re or Zr in different amounts . Alumina and several type of zeolites have been employed as carriers. The catalysts prepared were characterized and the most selective catalyst was determined after a selectivity-screening test. The most selective catalysts were found to be the ones prepared over ZSM5 carriers and consisting of Ga and/or Pt as active components. Methods: Catalyst preperation: Eighteen different catalysts with different types of carriers such as Alumina, Natural Zeolite (Clinoptilolite) and several synthetic zeolites ( MEM5766, MEM1510, ZSM5(Na) and ZSM5(H)-Pentasyl), were prepared by impregnation method to add the active components, e.g. the zeolite type carriers were contacted with aqueous solutions of salts of several metals such as Ni, Sn, Re, Pt, Ga, Zr, Sn. Chemical compositions of the catalysts prepared were determined by IR, Omnisorp 1000CX and DSC. Experimental Procedure: In all the experiments, a known amount of catalyst and hydrocarbon are charged to the reactor and the saturator, respectively, and they are equilibrated to the reaction conditions. The nitrogen flow carrying the hydrocarbon from the saturator into the reactor is opened and the experiment is started. Throughout the run the gas stream is measured continuously by the rotameter and was kept constant. The ovens heating the reactor and the saturator are connected with the automatic temperature control device and the temperatures of reactor and saturator were kept constant and recorded continuously during the experiment. The products and unreacted hydrcarbon were condensed in the traps cooled by salt-ice mixture and liquid nitrogen, and the products and unreacted m-xylene collected in the traps were analyzed by a gas chromatography Hewlett .Packard 5890 series 2; on a capillary column of HHP-FFAP using flame ionization dedector (FID). 9The experimental reaction conditions are varied as follows:

464 Reaction Time (h): 4 (for selectivity exp.) and 8 (for kinetic exp.) Reaction Temperature (~ 270 - 380. Reaction Pressure (atm): 1 Nitrogen flow (mYsec): 1 - 2. Catalyst weight (g): 0.5 - 1 Evaluation of the Experiments: A kinetic model given below in Figure 1 including isomerization and disproportination reactions was developed and tested for the goodness of fit of the model with the experimental data. Following assumptions were done during the evaluation of the experimental data; i. The reactions between o-xylene and p-xylene and ii. The reactions between methyle groups were neglected The kinetic parameters of the model were determined and Arrhenius law dependency of the rate costants were obtained. Experimental rate costants were compared with those given in literature. Results:Selectivity experiments were carried out at several conditions given before for all the catalysts prepared. At the end of selectivity experiments, two ZSM5 catalysts with the active component of 1.8*103% Pt and with the active components of 1.8*103% Pt and 3.04*102% Ga were found to be as the most selective ones in the isomerization reaction of m-xylene. Kinetic runs EQUILIBRIUM REACTIONS OF XYLENES

CH3 o-Xvlene

CH

~.-"CH3 r

m-Xylene ~ C ~ . WH3 ~%

~

C ~ . pX .yelne

SIDE REACTIONS I. DISPROPORTIONATION IL DEMETHYLATION

&

IL

2 Xylcnes

CH3 ~ CH,

+H+ CH3 ~

~ 1 , 2 , 3 or 1,2,4-TMB -q- L~.._~ CH3 Toluene CH3 ~

CH~

Me.vie -%

CH3-

CH3 2

~ C H 3

_q_

CH3- ~

CH3- -I- CH3- ~CH3-CH3

~

CH3 -~CH3 CH3- ~

4"

H+ -'i- CH4

Figure 1: Reactions in Isomerization of Xylenes

465 were m a d e at a t e m p e r a t u r e range of 270-380~ a n d at several initial concentrations of hydrocarbon over 1 g of catalyst.First, conversions to products a n d total conversion of m-xylene in all the experiments were calculated. The curves of conversion vs. space time were plotted for the most selective catalysts a n d the comparision of total conversion vs. W/FAo curves at different t e m p e r a t u r e s were p r e s e n t e d in Figure 2. It is obvious t h a t there is an appreciable increase in conversion to each product as the t e m p e r a t u r e increases from 270~ to 340~ Above 340~ such a significant change can not be observed in the conversions. The performances of the catalysts are almost the same at these reaction t e m p e r a t u r e . Because of the fact that total conversions achieved in all the experiments are greater than 10 %, the reactor is treated as an integral. Selectivities of products have been defined as the ratios of related rate constants (4). First order reaction mechanism (for isomerization reaction reversible, for disproportionation and demethylation irreversible) was found to fit the experimental data fairly well. The experimental results over the most selective catalysts, the kinetic parameters, activation energies and frequency factors calculated by Arrhenius equation were given in Table 1.1 and 1.2.

Figure 2: Comparision of total conversion vs. W/FAo curves at different temperatures

Table 1.1" E x p e r i m e n t a l Results T C

p-x

p-/e-

p4(ToI+TMB)

270

0.15-0.18

1.82-1.86

0.19-0.20 0.16-0.18 0.18-0.19 0.18-0.19

1.46-1.51 1.83-1.85 1.58-1.60 1.47-1.50

9.44-9.80 14.68-15.07 5.88-6.00 4.71-4.83 4.01-4.04 1.13-2.24

13 PtGa 340 | 5 Pt

S%

x%

380 270 340 380

0.18-0.20 1.80-1.86

466 T a b l e 1.2 : Kinetic p a r a m e t e r s calculated

kpx

, I

k-p, I ko, I k~, I kTo,~ [ kTMS k=A*e -v~vr A:Frequency Fac. E:Cal/mole

A 13 PtGa E

Corr.

15 Pt

A E Corr.

4.41982

2.07052

9.73376

2.88839

1.12524

1.14821

1887.65 0.98 3.6704

1887.55 0.98 1.29693

3040.11 0.98 1.80399

3040.11 0.98 3.67811

79.48 0.92 3.67811

99.35 0.99 635.84

1331.29 0.92

834.54 0.64

1510.12 0.90

437.14 0.90

1271.68 0.82

5358.4 0.99 .

As seen from Table 1.1, addition of Ga promoter to ZSM5Pt catalyst increases the selectivity to p-xylene at 340~ although the conversions to p-xylene do not change significantly at the same or higher temperatures. For comparison with literature as seen Table 2, the studies done by Dimitriu et al and Li et al. (5) have been selected. Dimitriu et al. tested the zeolites SK-500, HndY and HZSM5 in the reaction of m-xylene isomerization. It was found t h a t HZSM5 h a d the best selectivity for p-xylene formation. Three kinetic models including isomerization and disproportionation reactions were employed to explain the kinetic data. It was observed t h a t as the t e m p e r a t u r e was increased, the rate constants were also increased, as expected. Li and co-workers used a pulse micro reactor chromatography technique to study the xylene isomerization reaction on HZSM5 zeolite catalyst. A mathematical model including diffusion, adsorption and reaction steps was developed. They also declared that the reaction rate constants increased with increasing temperature.Rate constants calculated in this study also increase with increasing temperature, which is the right trend for the constants. Activation energies obtained in this study are very low with respect to the activation energies given by Li et al.. But there is a good fitness between the values of the reverse reaction rate constants of the isomerization reaction m e a s u r e d in this study over ZSM5PtGa catalyst and by Li et al. over HZSM5 catalyst. It is obvious t h a t the numerical values of the kinetic p a r a m e t e r s given by Dumitriu et al. are very different from those given by this study and by Li et am..

References: 1. Sachtler, A., and et al., US Patent 4,957,891, 1990. 2. Onodera, T., and et al., 1987, US Patent 4,700,012. 3. Karge, H. G., 1996, "Zeolites- Their Acidic and Basic Properties Related to Catalysis", Proc. of the Sec. Turk. Chem. Eng. Congress, i.T.U. Istanbul, Proc. Inv. Lect., 32-55. 4. Dumitriu, E., Oprea, S., Hulea, V., Activity and Selectivity of Zeolite Catalysts in m-Xylene Izomerization, Revue Roumanie de Chimie, 32, 5, 525-532, 1987. 5, IA, Y., Chang, X., Zeng, Z., Kinetic Study of the Izomerization of Xylene on HZSM-5 Zeolite. 1. Kinetics Model and Reaction Mechanism, Ind. Eng. Chem. Res., 31, 1, 187-192, 1992. 6. Akpolat,/~, Investigation of Isomerization Kinetics of m-Xylene, Ph.D Thesis, Eng. Fac of Ege Univ., 1998.

467

Abatzoglou Ch. Ahlafi H. Akpolat O. AI-Khowaiter S.H. AI-Kinany M.C. Almasan V. Amon B. Angelidis T.-N. Ansell G.P. Antons U. Apicella B. Aracil J. Arai M. Arena G.E. Ashour I. Backman L.B. Baddeley C.J. Barelko V. Barth D. Bekurdts A. Bellobono I.R. Bennett P.S. Bhanage B.M. Bhargava A. Bianchi D. Bianchini A. Borio D.O. Braun S. Brehm A. Brenner G. Brundage M.A. Burch R. Casanave D. Cavers M. Centi G. Chuang S.S.C. Colaris A.H.J. C6t~ A.S. Cox J.P. Cullen B. Cunningham R.H. Dalmon J.A. Datta R. Davidson J.M. de Lange W.

175 419 463 375 375 435 247 341 157 451 431 317,325 427 109 229 439 11 191 403 451 385 157 427 219 419 109 117 307 451 307 83 447, 455 367 65 109 83 93 199 157 125 393 367 275 65 255

468 D~camp T. Delgass W.N. Dennison P.R. Di Serio M. Diskin A.M. Doepper R. Dubuis S. Durst F. Elnashaie S.S. Emig G. Errazu A.F. Evans J.M. Fiaty K. Fishtik I. Forissier M. Freygang M. Friedrich G. Froment G.F. Frost J.C. Garcia D. Garcfa R. Garcfa T. Gatica J.E. Gawlik B.M. Gladden L.F. Gray P.G. Grenfell J. Grzesik M. G~ndiJz G. Haq S. Harkness I.R. Harmsen J.M.A. Hedrick S.A. Hoebink J.H.B. HOnicke D. lengo P. Jackson S.D. Jalibert J.C. J~ims~-Jounela S.-L. Jamshidi A.M. Jones A.-M. Kalantzopoulos A. K~immerer M. Kanervo J. Kantzas A. Katsanos N.A. Kennedy D.R. King F. Kiwi-Minsker L. Klemm E. Kourtakis K. Krause A.O.I.

133 199 125 267, 431 393 359 359 307 229 247, 307 141 157 367 275 367 237 237 333 157 325 317 31 7, 325 117 385 167, 183 157 149 407, 411,415 463 11 65 101 83 93, 101 47 267 125, 149, 291 133 439 423 157 175 255 439 219 175 125 291 191 247 73 439

469 Kulawska M. Kung H.H. Langford C.H. Lazar M. Lemos F. Lemos M.A.N. Lennon D. Litorell M. Louafi S. Lukyanov D.B. Marginean P. Martens G. Martfnez M. Matheson I.M. McCracken J.S. McDougall G.S. McLeod A.S. Metcalfe I.S. Mikkola J.-P. Mills P.L. Mirodatos C. Moroni A. Moustafa T.M. Murray P. Muryn C. Muzykantov V.S. Nawdali M. Norskov J.K. Ormerod R.M. Ortega Lorenzo M. Palma C. Paul S. Pombeiro A.J.L. Porras J.A. Rajaram R.R. Ramkrishna D. Ram6a Ribeiro F. Randall H.T. Raval R. Rees L.V.C. Renken A. Rigby S.P. Roubani-Kalantzopoulou F. Ruf S. Sahibzada M. Salmi T. Santacesaria E. Santos Silva I. Schbib N.S. Sch6nfelder H. Schouten J.C. Schunk S.

411 23 219 435 443, 459 443 125 157 11 299 435 333 31 7, 325 149 209 65 167 403 351 73, 209 133 385 229 11 11 447 419 3 35, 393 11 459 283 443 141 157 199 443 73, 209 11 65 191,359 183 175 307 403 351 267, 431 459 11 7, 141 255 93, 101 307

470 Schuurman Y. Shestov A.A. Shirai M. Sj6holm R. Skrzypek J. Smedler G. Sohrabi M. Sousa Lobo L. Sousa P. Starosud A. Stitt E.H. Sullivan J.A. Sultan M. Tzitzios V. Vanhove D. Vazzana F. Veser G. Vitali P. Walker A.P. Webb G. Weber Th. Williams J. Witczak M. Yuranov I. Zengerle R. Zhao F.-G.

133 447, 455 427 351 407, 411,415 157 423 459 443 219 291 447, 455 283 341 283 109 237 109 157 125, 149 307 11 411,415 191 237 427

471 STUDIES IN SURFACE SCIENCE AND CATALYSIS

Advisory Editors: B. Delmon, Universit~ Catholique de Louvain, Louvain-la-Neuve, Belgium J.T. Yates, University of Pittsburgh, Pittsburgh, PA, U.S.A. Volume 1

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Preparation of Catalysts I.Scientific Basesfor the Preparation of Heterogeneous Catalysts. Proceedings of the First International Symposium, Brussels, October 14-17,1975 edited by B. Delmon, P.A.Jacobs and G. Poncelet The C~ntrol of the Reactivity of Solids. A Critical Survey of the Factors that Influence the Reactivity of Solids, with Special Emphasis on the Control of the Chemical Processes in Relation to Practical Applications by V.V. Boldyrev, M. Bulens and B. Delmon Preparation of Catalysts II. Scientific Basesfor the Preparation of Heterogeneous Catalysts. Proceedingsofthe Second International Symposium, Louvain-la-Neuve, September 4-7, 1978 edited by B. Delmon, P. Grange, P.Jacobs and G. Poncelet Growth and Properties of Metal Clusters. Applications to Catalysis and the Photographic Process. Proceedings of the 32nd International Meeting of the Soci~t6 de Chimie Physique, Villeurbanne, September 24-28, 1979 edited by J. Bourdon Catalysis by Zeolites. Proceedings of an International Symposium, Ecully (Lyon), September 9-11, 1980 edited by B. Imelik, C. Naccache, Y. Ben Taarit, J.C. Vedrine, G. Coudurier and H. Praliaud Catalyst Deactivation. Proceedings of an International Symposium, Antwerp, October 13-15,1980 edited by B. Delmon and G.F. Froment New Horizons in Catalysis. Proceedings of the 7th International Congress on Catalysis, Tokyo, June 30-J uly4, 1980. Parts A and B edited by T. Seiyama and K. Tanabe Catalysis by Supported Complexes by Yu.l. Yermakov, B.N. Kuznetsov and V.A. Zakharov Physics of Solid Surfaces. Proceedings of a Symposium, Bechyhe, September 29-October 3,1980 edited by M. LazniEka Adsorption at the Gas-Solid and Liquid-Solid Interface. Proceedings of an International Symposium, Aix-en-Provence, September 21-23, 1981 edited by J. Rouquerol and K.S.W. Sing Metal-Support and Metal-Additive Effects in Catalysis. Proceedings of an International Symposium, Ecully (Lyon), September 14-16, 1982 edited by B. Imelik, C. Naccache, G. Coudurier, H. Praliaud, P. Meriaudeau, P. Gallezot, G.A. Martin and J.C. Vedrine Metal Microstructures in Zeolites. Preparation - Properties- Applications. Proceedings of a Workshop, Bremen, September 22-24, 1982 edited by P.A. Jacobs, N.I. Jaeger, P.Jir~ and G. Schulz-Ekloff Adsorption on Metal Surfaces. An Integrated Approach edited by J. Benard Vibrations at Surfaces. Proceedings of the Third International Conference, Asilomar, CA, September 1-4, 1982 edited by C.R. Brundle and H. Morawitz Heterogeneous Catalytic Reactions Involving Molecular Oxygen by G.I. Golodets

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Preparation of Catalysts III. Scientific Basesfor the Preparation of Heterogeneous Catalysts. Proceedings ofthe Third International Symposium, Louvain-la-Neuve, September 6-9, 1982 edited by G. Poncelet, P. Grange and P.A. Jacobs Spillover of Adsorbed Species. Proceedings of an International Symposium, Lyon-Villeurbanne, September 12-16, 1983 edited by G.M. Pajonk, S.J. Teichner and J.E. Germain Structure and Reactivity of Modified Zeolites. Proceedings of an International Conference, Prague, July 9-13, 1984 edited by P.A. Jacobs, N.I. Jaeger, P.JiE=,V.B. Kazansky and G. Schulz-Ekloff Catalysis on the Energy Scene. Proceedings ofthe 9th Canadian Symposium on Catalysis, Quebec, P.Q., September 30-October 3, 1984 edited by S. Kaliaguine and A. Mahay Catalysis by Acids and Bases. Proceedings of an International Symposium, Villeurbanne (Lyon), September 25-27, 1984 edited by B. Imelik, C. Naccache, G. Coudurier, Y. Ben Taarit and J.C. Vedrine Adsorption and Catalysis on Oxide Surfaces. Proceedings of a Symposium, Uxbridge, June 28-29, 1984 edited by M. Che and G.C. Bond Unsteady Processes in Catalytic Reactors by Yu.Sh. Matros Physics of Solid Surfaces 1984 edited by J. Koukal Zeolites: Synthesis, Structure, Technology and Application. Proceedings of an International Symposium, Portoro~-Portorose, September 3-8, 1984 edited by B. Dr:;aj, S. HoEevar and S. Pejovnik Catalytic Polymerization of Olefins. Proceedings of the International Symposium on Future Aspects of Olefin Polymerization, Tokyo, July 4-6, 1985 edited by T. Keii and K. Soga Vibrations at Surfaces 1985. Proceedings ofthe Fourth International Conference, Bowness-on-Windermere, September 15-19, 1985 edited by D.A. King, N.V. Richardson and S. Holloway Catalytic Hydrogenation edited by L. Cerven~ New Developments in Zeolite Science and Technology. Proceedings of the 7th International Zeolite Conference, Tokyo, August 17-22, 1986 edited by Y. Murakami, A. lijima and J.W. Ward Metal Clusters in Catalysis edited by B.C. Gates, L. Guczi and H. Kn6zinger Catalysis and Automotive Pollution Control. Proceedings of the First International Symposium, Brussels, September 8-11, 1986 edited by A. Crucq and A. Frennet Preparation of Catalysts IV. Scientific Basesfor the Preparation of Heterogeneous Catalysts. Proceedings ofthe Fourth International Symposium, Louvain-laNeuve, September 1-4, 1986 edited by B. Delmon, P. Grange, P.A. Jacobs and G. Poncelet Thin Metal Films and Gas Chemisorption edited by P.Wissmann Synthesis of High-silica Aluminosilicate Zeolites edited by P.A. Jacobs and J.A. Martens Catalyst Deactivation 1987. Proceedings of the 4th International Symposium, Antwerp, September 29-October 1, 1987 edited by B. Delmon and G.F. Froment Keynotes in Energy-Related Catalysis edited by S. Kaliaguine

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Methane Conversion. Proceedingsof a Symposium on the Production of Fuels and Chemicals from Natural Gas, Auckland, April 27-30, 1987 edited by D.M. Bibby, C.D. Chang, R.E Howe and S. Yurchak Innovation in Zeolite Materials Science. Proceedings of an International Symposium, Nieuwpoort, September 13-17, 1987 edited by P.J. Grobet, W.J. Mortier, E.F.Vansant and G. Schulz-Ekloff Catalysis 1987. Proceedings of the 10th North American Meeting of the Catalysis Society, San Diego, CA, May 17-22, 1987 edited by J.W. Ward Characterization of Porous Solids. Proceedings of the IUPAC Symposium (COPS I), Bad Soden a. Ts., April 26-29,1987 edited by K.K. Unger, J. Rouquerol, K.S.W. Sing and H. Kral Physics of Solid Surfaces 1987. Proceedings of the Fourth Symposium on Surface Physics, Bechyne Castle, September 7-11, 1987 edited by J. Koukal Heterogeneous Catalysis and Fine Chemicals. Proceedings of an International Symposium, Poitiers, March 15-17, 1988 edited by M. Guisnet, J. Barrault, C. Bouchoule, D. Duprez, C. Montassier and G. P6rot Laboratory Studies of Heterogeneous Catalytic Processes by E.G. Christoffel, revised and edited by Z. Padl Catalytic Processes under Unsteady-State Conditions by Yu. Sh. Matros Successful Design of Catalysts. Future Requirements and Development. Proceedings ofthe Worldwide Catalysis Seminars, July, 1988, on the Occasion of the 30th Anniversary of the Catalysis Society of Japan edited by T. Inui Transition Metal Oxides. Surface Chemistry and Catalysis by H.H. Kung Zeolites as Catalysts, Sorbents and Detergent Builders. Applications and Innovations. Proceedings of an International Symposium, W~rzburg, September 4-8,1988 edited by H.G. Karge and J. Weitkamp Photochemistry on Solid Surfaces edited by M. Anpo and T. Matsuura Structure and Reactivity of Surfaces. Proceedings of a European Conference, Trieste, September 13-16, 1988 edited by C. Morterra, A. Zecchina and G. Costa Zeolites: Facts, Figures, Future. Proceedings of the 8th International Zeolite Conference, Amsterdam, July 10-14, 1989. Parts A and B edited by RA. Jacobs and R.A. van Santen Hydrotreating Catalysts. Preparation, Characterization and Performance. Proceedings ofthe Annual International AIChE Meeting, Washington, DC, November 27-December 2, 1988 edited by M.L. Occelli and R.G. Anthony New Solid Acids and Bases. Their Catalytic Properties by K. Tanabe, M. Misono, Y. Ono and H. Hattori Recent Advances in Zeolite Science. Proceedings of the 1989 Meeting of the British Zeolite Association, Cambridge, April 17-19, 1989 edited by J. Klinowsky and RJ. Barrie Catalyst in Petroleum Refining 1989. Proceedings of the First International Conference on Catalysts in Petroleum Refining, Kuwait, March 5-8, 1989 edited by D.L. Trimm, S. Akashah, M. Absi-Halabi and A. Bishara Future Opportunities in Catalytic and Separation Technology edited by M. Misono, Y. Moro-oka and S. Kimura

474 New Developments in Selective Oxidation. Proceedings of an International Symposium, Rimini, Italy, September 18-22, 1989 edited by G. Centi and F.Trifiro Olefin Polymerization Catalysts. Proceedings of the International Symposium Volume 56 on Recent Developments in Olefin Polymerization Catalysts, Tokyo, October 23-25, 1989 edited by T. Keii and K. Soga Volume 57A Spectroscopic Analysis of Heterogeneous Catalysts. Part A: Methods of Surface Analysis edited by J.L.G. Fierro Volume 57B Spectroscopic Analysis of Heterogeneous Catalysts. Part B: Chemisorption of Probe Molecules edited by J.L.G. Fierro Introduction to Zeolite Science and Practice Volume 58 edited by H. van Bekkum, E.M. Flanigen and J.C. Jansen Heterogeneous Catalysis and Fine Chemicals II. Proceedings of the 2nd Volume 59 International Symposium, Poitiers, October 2-6, 1990 edited by M. Guisnet, J. Barrault, C. Bouchoule, D. Duprez, G. Perot, R. Maurel and C. Montassier Chemistry of Microporous Crystals. Proceedings of the International Symposium Volume 60 on Chemistry of Microporous Crystals, Tokyo, June 26-29, 1990 edited by T. Inui, S. Namba and T. Tatsumi Natural Gas Conversion. Proceedings of the Symposium on Natural Gas Volume 61 Conversion, Oslo, August 12-17, 1990 edited by A. Holmen, K.-J. Jens and S. Kolboe Characterization of Porous Solids II. Proceedings of the IUPAC Symposium Volume 62 (COPS II), Alicante, May 6-9, 1990 edited by F. Rodriguez-Reinoso, J. Rouquerol, K.S.W. Sing and K.K. Unger Preparation of Catalysts V. Scientific Bases for the Preparation of Heterogeneous Volume 63 Catalysts. Proceedings of the Fifth International Symposium, Louvain-la-Neuve, September 3-6, 1990 edited by G. Poncelet, RA. Jacobs, P. Grange and B. Delmon Volume 64 New Trends in CO Activation edited by L. Guczi Catalysis and Adsorption by Zeolites. Proceedings of ZFOCAT 90, Leipzig, Volume 65 August 20-23, 1990 edited by G. (~hlmann, H. Pfeifer and R. Fricke Dioxygen Activation and Homogeneous Catalytic Oxidation. Proceedings of the Volume 66 Fourth International Symposium on Dioxygen Activation and Homogeneous Catalytic Oxidation, Balatonffired, September 10-14, 1990 edited by L.I. Simandi Structure-Activity and Selectivity Relationships in Heterogeneous Catalysis. Volume 67 Proceedings of the ACS Symposium on Structure-Activity Relationships in Heterogeneous Catalysis, Boston, MA, April 22-27, 1990 edited by R.K. Grasselli and A.W. Sleight Catalyst Deactivation 1991. Proceedings of the Fifth International Symposium, Volume 68 Evanston, IL, June 24-26, 1991 edited by C.H. Bartholomew and J.B. Butt Zeolite Chemistry and Catalysis. Proceedings of an International Symposium, Volume 69 Prague, Czechoslovakia, September 8-13, 1991 edited by P.A.Jacobs, N.I. Jaeger, L. Kubelkova and B. Wichterlova Poisoning and Promotion in Catalysis based on Surface Science Concepts and Volume 70 Experiments by M. Kiskinova Volume 55

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Catalysis and Automotive Pollution Control II. Proceedings of the 2nd International Symposium (CAPoC 2), Brussels, Belgium, September 10-13, 1990 edited by A. Crucq New Developments in Selective Oxidation by Heterogeneous Catalysis. Proceedings of the 3rd European Workshop Meeting on New Developments in Selective Oxidation by Heterogeneous Catalysis, Louvain-la-Neuve, Belgium, April 8-10, 1991 edited by P. Ruiz and B. Delmon Progress in Catalysis. Proceedings of the 12th Canadian Symposium on Catalysis, Banff, Alberta, Canada, May 25-28, 1992 edited by K.J. Smith and E.C. Sanford Angle-Resolved Photoemission. Theory and Current Applications edited by S.D. Kevan New Frontiers in Catalysis, Parts A-Co Proceedings of the 10th International Congress on Catalysis, Budapest, Hungary, 19-24 July, 1992 edited by L. Guczi, F. Solymosi and R Tdtenyi Fluid Catalytic Cracking: Science and Technology edited by J.S. Magee and M.M. Mitchell, Jr. New Aspects of Spillover Effect in Catalysis. For Development of Highly Active Catalysts. Proceedings of the Third International Conference on Spillover, Kyoto, Japan, August 17-20, 1993 edited by T. Inui, K. Fujimoto, T. Uchijima and M. Masai Heterogeneous Catalysis and Fine Chemicals III. Proceedings ofthe 3rd International Symposium, Poitiers, April 5-8, 1993 edited by M. Guisnet, J. Barbier, J. Barrault, C. Bouchoule, D. Duprez, G. Perot and C. Montassier Catalysis: An Integrated Approach to Homogeneous, Heterogeneous and Industrial Catalysis edited by J.A. Moulijn, P.W.N.M. van Leeuwen and R.A. van Santen Fundamentals of Adsorption. Proceedings of the Fourth International Conference on Fundamentals of Adsorption, Kyoto, Japan, May 17-22, 1992 edited by M. Suzuki Natural Gas Conversion II. Proceedings of the Third Natural Gas Conversion Symposium, Sydney, July 4-9, 1993 edited by H.E. Curry-Hyde and R.F. Howe New Developments in Selective Oxidation II. Proceedings of the Second World Congress and Fourth European Workshop Meeting, Benalm&dena, Spain, September 20-24, 1993 edited by V. Cortes Corberan and S. Vic Bellon Zeolites and Microporous Crystals. Proceedings of the International Symposium on Zeolites and Microporous Crystals, Nagoya, Japan, August 22-25, 1993 edited by T. Hattori and T. Yashima Zeolites and Related Microporous Materials: State of the Art 1994. Proceedings of the 10th International Zeolite Conference, Garmisch-Partenkirchen, Germany, July 17-22, 1994 edited by J. Weitkamp, H.G. Karge, H. Pfeifer and W. H61derich Advanced Zeolite Science and Applications edited by J.C. Jansen, M. St6cker, H.G. Karge and J.Weitkamp Oscillating Heterogeneous Catalytic Systems by M.M. Slin'ko and N.I. Jaeger Characterization of Porous Solids III. Proceedings of the IUPAC Symposium (COPS III), Marseille, France, May 9-12, 1993 edited by J.Rouquerol, F. Rodriguez-Reinoso, K.S.W. Sing and K.K. Unger

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Catalyst Deactivation 1994. Proceedings of the 6th International Symposium, Ostend, Belgium, October 3-5, 1994 edited by B. Delmon and G.F. Froment Catalyst Design for Tailor-made Polyolefins. Proceedings of the International Symposium on Catalyst Design for Tailor-made Polyolefins, Kanazawa, Japan, March 10-12, 1994 edited by K. Soga and M. Terano Acid-Base Catalysis II. Proceedings of the International Symposium on Acid-Base Catalysis II, Sapporo, Japan, December 2-4, 1993 edited by H. Hattori, M. Misono and Y. Ono Preparation of Catalysts VI. Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings of the Sixth International Symposium, Louvain-La-Neuve, September 5-8, 1994 edited by G. Poncelet, J. Martens, B. Delmon, RA. Jacobs and R Grange Science and Technology in Catalysis 1994. Proceedings of the Second Tokyo Conference on Advanced Catalytic Science and Technology, Tokyo, August 21-26, 1994 edited by Y. Izumi, H. Arai and M. Iwamoto Characterization and Chemical Modification of the Silica Surface by E.F.Vansant, R Van Der Voort and KoC.Vrancken Catalysis by Microporous Materials. Proceedings of ZEOCAT'95, Szombathely, Hungary, July 9-13, 1995 edited by H.K. Beyer, H.G.Karge, I. Kiricsi and J.B. Nagy Catalysis by Metals and Alloys by V. Ponec and G.C. Bond Catalysis and Automotive Pollution Control III. Proceedings of the Third International Symposium (CAPoC3), Brussels, Belgium, April 20-22, 1994 edited by A. Frennet and J.-M. Bastin Zeolites: A Refined Tool for Designing Catalytic Sites. Proceedings of the International Symposium, Quebec, Canada, October 15-20, 1995 edited by L. Bonneviot and S. Kaliaguine Zeolite Science 1994: Recent Progress and Discussions. Supplementary Materials to the 10th International Zeolite Conference, Garmisch-Partenkirchen, Germany, July 17-22, 1994 edited by H.G. Karge and J. Weitkamp Adsorption on New and Modified Inorganic Sorbents edited by A. Dqbrowski and V.A. Tertykh Catalysts in Petroleum Refining and Petrochemical Industries 1995. Proceedings of the 2nd International Conference on Catalysts in Petroleum Refining and Petrochemical Industries, Kuwait, April 22-26, 1995 edited by M. Absi-Halabi, J. Beshara, H. Qabazard and A. Stanislaus 1lth International Congress on Catalysis - 40th Anniversary. Proceedings ofthe 1lth ICC, Baltimore, MD, USA, June 30-July 5, 1996 edited by J. W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell Recent Advances and New Horizons in Zeolite Science and Technology edited by H. Chon, S.I. Woo and S. -E. Park Semiconductor Nanoclusters - Physical, Chemical, and Catalytic Aspects edited by RV. Kamat and D. Meisel Equilibria and Dynamics of Gas Adsorption on Heterogeneous Solid Surfaces edited by W. Rudzifiski, W.A. Steele and G. Zgrablich Progress in Zeolite and Microporous Materials Proceedings of the 1lth International Zeolite Conference, Seoul, Korea, August 12-17, 1996 edited by H. Chon, S.-K. Ihm and Y.S. Uh

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Hydrotreatment and Hydrocracking of Oil Fractions Proceedings of the 1st International Symposium / 6th European Workshop, Oostende, Belgium, February 17-19, 1997 edited by G.F. Froment, B. Delmon and R Grange Volume 107 Natural Gas Conversion IV Proceedings of the 4th International Natural Gas Conversion Symposium, Kruger Park, South Africa, November 19-23, 1995 edited by M. de Pontes, R.L. Espinoza, C.R Nicolaides, J.H. Scholtz and M.S. Scurrell Volume 108 Heterogeneous Catalysis and Fine Chemicals IV Proceedings of the 4th International Symposium on Heterogeneous Catalysis and Fine Chemicals, Basel, Switzerland, September 8-12, 1996 edited by H.U. Blaser, A. Baiker and R. Prins Volume 109 Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis. Proceedings ofthe International Symposium, Antwerp, Belgium, September 15-17,1997 edited by G.F. Froment and K.C. Waugh Volume 110 Third World Congress on Oxidation Catalysis. Proceedings of the Third World Congress on Oxidation Catalysis, San Diego, CA, U.S.A., 21-26 September 1997 edited by R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons 'olume 111 Catalyst Deactivation 1997. Proceedings ofthe 7th International Symposium, Cancun, Mexico, October 5-8, 1997 edited by C.H. Bartholomew and G.A. Fuentes Volume 112 Spillover and Migration of Surface Species on Catalysts. Proceedings of the 4th International Conference on Spillover, Dalian, China, September 15-18, 1997 edited by Can Li and Qin Xin Volume 113 Recent Advances in Basic and Applied Aspects of Industrial Catalysis. Proceedings ofthe 13th National Symposium and Silver Jubilee Symposium of Catalysis of India, Dehradun, India, April 2-4, 1997 edited by T.S.R. Prasada Rao and G. Murali Dhar Volume 114 Advances in Chemical Conversions for Mitigating Carbon D|oxide. Proceedings of the 4th International Conference on Carbon Dioxide Utilization, Kyoto, Japan, September 7-11, 1997 edited by T. Inui, M. Anpo, K. Izui, S. Yanagida and T. Yamaguchi Volume 115 Methods for Monitoring and Diagnosing the Efficiency of Catalytic Converters. A patent-oriented survey by M. Sideris Volume 116 Catalysis and Automotive Pollution Control IV. Proceedings of the 4th International Symposium (CAPoC4), Brussels, Belgium, April 9-11, 1997 edited by N. Kruse, A. Frennet and J.-M. Bastin Volume 117 Mesoporous Molecular Sieves 1998 Proceedings of the 1st International Symposium, Baltimore, MD, U.S.A., July 10-12, 1998 edited by L.Bonneviot, F. B61and, C. Danumah, S. Giasson and S. Kaliaguine Volume 118 Preparation of Catalysts VII Proceedings of the 7th International Symposium on Scientific Bases for the Preparation of Heterogeneous Catalysts, Louvain-la-Neuve, Belgium, September 1-4, 1998 edited by B. Delmon, RA. Jacobs, R. Maggi, J.A. Martens, R Grange and G. Poncelet Volume 119 Natural Gas Conversion V Proceedings of the 5th International Gas Conversion Symposium, Giardini-Naxos, Taormina, Italy, September 20-25, 1998 edited by A. Parmaliana, D. Sanfilippo, F. Frusteri, A. Vaccari and F.Arena Volume 120A Adsorption and its Applications in Industry and Environmental Protection. Vol I: Applications in Industry edited by A. Dabrowski

Volume 120B Adsorption and its Applications in Industry and Environmental Protection. Vol I1:Applications in Environmental Protection edited by A. Dabrowski Volume 121 Science and Technology in Catalysis 1998 Proceedings of the Third Tokyo Conference in Advanced Catalytic Science and Technology, July 19-24, 1998 Edited by H. Hattori and K. Otsuka Volume 122 Reaction Kinetics and the Development of Catalytic Processes Proceedings ofthe International Symposium, Brugge, Belgium, April 19-21, 1999 Edited by G.F. Froment and K.C. Waugh

Reaction Kinetics and the Development of Catalytic Processes (Studies in Surface Science and Catalysis)

Fluid Catalytic Cracking VII:, Volume 166: Materials, Methods and Process Innovations (Studies in Surface Science and Catalysis) (Studies in Surface Science and Catalysis)

Catalytic Polymerization of Cycloolefins (Studies in Surface Science and Catalysis, 131)

Catalysts for Upgrading Heavy Petroleum Feeds, Volume 169 (Studies in Surface Science and Catalysis) (Studies in Surface Science and Catalysis)

Catalytic Kinetics

Catalysis: An Integrated Approach to Homogenous, Heterogenous and Industrial Catalysis (Studies in Surface Science and Catalysis)

Past and Present in DeNOx Catalysis: From Molecular Modelling to Chemical Engineering, Volume 171 (Studies in Surface Science and Catalysis) (Studies in Surface Science and Catalysis)

Reaction Kinetics and the Development of Catalytic Processes (Studies in Surface Science and Catalysis)